Solubility Product Constant & The Common Ion Effect Paper

 

 

CHEMISTRY 213, 214, 272

Laboratory Manual

Fall

 

201

8

 

 

Suzanne W. Slayden

 

word image 3206

and

Faculty in the Department of Chemistry & Biochemistry

George Mason University

2018

 

CHEM 213/214/272 Laboratory Manual

Fall 2018

George Mason University

 

Table of Contents

Preface iv

General Safety Guidelines and Operations in the General Chemistry Labs v

KEEPING A LABORATORY NOTEBOOK vii

CHEM 213 / CHEM 272 EXPERIMENTS 1

Experiment 1 Pennies and good Sense 2

Experiment 2 Measurements 7

Experiment 3 Qualitative Analysis of Cations 19

Experiment 4 Density of Solutions 24

Experiment 5 Empirical Formula of Zinc Iodide 35

Experiment 6 Synthesis of Aspirin 42

Experiment 7 Volumetric Analysis: Titration of Vinegar 51

Experiment 8 The Ideal Gas Law 66

Experiment 9 Heat of Reaction: Hess’s Law 79

Experiment 10 Solvent Polarities from a Solvatochromic Dye 93

Experiment 11 VSEPR and Molecular Modeling 101

Experiment 12 Estimation of Avogadro’s Number 108

Experiment 13 Graphing 111

Experiment 14 Thin-Layer Chromatography 115

Experiment 15 Chemical Interactions 121

Experiment 16 Absorption Spectroscopy 129

Experiment 17 Periodic Properties: Cation Analysis 145

Experiment 18 Periodic Properties: Anion Analysis 153

CHEM 214, CHEM 272 EXPERIMENTS 157

Experiment 19 Synthesis of a Coordination Compound 158

Experiment 20 Determination of Water of Hydration in KTOF3 166

Experiment 21 Redox Titration of Oxalate in KTOF3 172

Experiment 22 Determination of Molecular Mass by Freezing Point Depression 179

Experiment 23 Chemical Kinetics 186

Experiment 24 Measurement of an Equilibrium Constant 195

Experiment 25 Acidity Constant from a Titration Curve 207

Experiment 26 Dissociation Constants of Acids and Bases 213

Experiment 27 Solubility Product Constant (Ksp) and Common-Ion Effect 222

Experiment 28 Thermodynamics of the Dissolution of Borax 231

Experiment 29 .Electrochemical Measurements 238

Experiment 30 Identification of an Organic Acid 245

APPENDIX 250

A 1. Instructions for the pH probe 252

A 2. Temperature probe 256

A 3. Spectro Vis instructions 258

 

PREFACE

The experiments is in this lab manual are designed for students that are currently enrolled in a science / engineering program at George Mason University.

Overall, there are two categories of experiments: those that are meant to complement a lecture discussion on the same topic and those for which there are no relevant lecture text commentaries but which exemplify experimental principles and techniques. Visible-region spectroscopy, which is emphasized throughout the lab courses, is the precursor to the study of ultraviolet, infrared, and NMR spectroscopies covered in the organic chemistry laboratory courses. Organic compounds also figure prominently here because rarely is there time in lecture to survey organic chemistry and because biology majors (our main constituency) are exposed early to organic and biochemical structures in their biology studies.

As might be expected for a new edition, many people have contributed toward its improvement. GMU chemistry professors Gerald Weatherspoon, James Schornick and John Schreifels suggested procedures which were gratefully received. Most consequential have been those students (graduate teaching assistants Arion Ross, Yukiko Yarnall, Allison Leninger, Megan Devine, and Christopher Kenedy) who offered constructive comments from their practical experience with the previous edition.

 

 

GENERAL SAFETY GUIDELINES AND OPERATIONS IN THE GENERAL CHEMISTRY LABS

 

  1. Safety goggles must be worn in the laboratory at all times, whether you are actually performing an experiment or not. (The wearing of contact lenses is strongly discouraged.)
  2. In case of fire, accidents or injury (no matter how minor) notify the lab instructor at once.
  3. Note the location of fire extinguishers, safety showers, eyewash and blankets so that you can use them quickly in an emergency.
  4. If either a corrosive acid, base, or any other chemical comes in contact with your skin or clothing, wash the affected area immediately with copious amounts of water. Notify the instructor.
  5. Always read a label carefully before removing a chemical from its container.
  6. Exercise great care in noting the odor of fumes, and avoid breathing in fumes of any kind.
  7. Never point a test tube or other vessel containing a reacting mixture (especially when you are heating it) toward your neighbor or yourself.
  8. NO SMOKING, EATING, DRINKING, OR CHEWING OF GUM in the laboratory. Do not taste anything in the lab. (This applies to cigarettes, food, drinks and gum as well as chemicals.)
  9. Under no circumstances should you perform any unauthorized experiments.
  10. Keep aisles and benches clear of books, clothing, etc. Store unused belongings in the front area of the lab or in the cabinets provided. Do not use lab benches as seats.
  11. Work areas are to be cleaned and put back in order prior to exiting the lab (includes electronic balance area, fume hoods, etc.)
  12. Full and proper attire is required for all students entering the teaching labs—including proper shoes.
  13. Students are required to wear FULL SHOES—-no abbreviated shoes means no open-toed or open-heeled shoes. Ballet flats, flats,pumps and heels, wedges (including wedge boots), and moccasins with socks are not acceptable for lab entry nor any other shoes that do not completely cover the top of the students’ feet since they do not offer enough protection to the body. Flip flops, sandals, top-sider/boat shoes, espadrilles, and loafersare not acceptable forms of protective shoes for entering the chemistry lab.
  14. Cotton fabric shoes are absorbent and do not provide the necessary protection from chemical spills.
  15. Low-cut socks are unacceptable for lab entry. Full socks (crew length) must be worn at all times.
  16. Tights/leggings are not acceptable as “pants”. Pants worn to lab should allow enough space so that chemicals coming in contact with the body will not be held tightly to the body resulting in additional harm.
  17. Skinny jeans/pants are not allowed in the chemistry labs.
  18. Yoga pants and running tights are not allowed in the chemistry labs.
  19. Shorts, capri pants and any other form of abbreviated lower body wear are not allowed. Long pants worn must completely reach your shoes and not leave any area of the legs or feet exposed.
  20. Skirt lengths must adhere to the same requirements as pants; skirts should extend all the way to the shoes and not leave any areas of the legs or feet exposed. Skirts that do not meet this criterion are forbidden in the chemistry lab.
  21. Always keep long hair confined.
  22. Tank tops and abbreviated upper body wear are not allowed in the chemistry labs. As an additional layer of protective wear, the upper body should be covered by short-sleeve clothing.
  23. Lab coats are provided by the chemical preps room OR may be purchased directly from the GMU Bookstore. Lab coats purchased individually must be knee length; waist length laboratory coats are not allowed in the laboratory.
  24. Children, acquaintances and pets are not allowed in the lab.
  25. Cell phones, laptops and other portable electronic devices, may not be used in the laboratory. Cell phones should be stored in backpacks prior to entering lab and should not be removed until the student exits the lab. Violation will result in a loss of 40 points for the lab exercise being performed that day.
  26. Students that refuse to wear the proper protective equipment (PPE) will be dismissed from the lab for that day and a grade of zero recorded for the laboratory exercise. Additional information regarding PPE violations can be found on Blackboard.
  27. Students will receive one general warning about PPE violations, which is given during the first class meeting. The second violation will result in a loss of 40 points for the lab exercise being performed that day. Repeat violators will be referred to the General Chemistry Laboratory Coordinator, with the recommendation that the student be permanently removed from the lab course.
  28. Refer to the lab syllabus (posted on Blackboard) regarding laboratory operations.
  29. You will sign an acknowledgement sheet on the first day of the lab, stating that we went over these rules and regulations and you read them on Blackboard too.

 

 

KEEPING A LABORATORY NOTEBOOK

 

General Instructions

  1. Students enrolling in General Chemistry lab courses (CHEM213/ CHEM214 / CHEM251) are required to purchase the carbonless spiral bound notebook from the George Mason University Bookstore.
  2. Put your name and your instructor’s name on both the outside and inside of the notebook. This will prove useful in case the lab notebook is misplaced.
  3. Leave the first page blank, except for writing the word “Index” at the top of the page. The Index will be a chronological listing of the experiment numbers, the experiment titles and their page numbers.
  4. Make all notebook entries in PERMANENT INK using a blue or black ballpoint pen. Use the right-hand pages (white pages with light blue grid lines) for main entries and the left-hand pages (yellow backside of the carbon pages) for scratch work only.
  5. Do not try to erase or white-out; draw a single line through the incorrect work or entry.
  6. Maintain chronological entries; do not leave blank pages. Each main page must be dated at the top of the page.
  7. All data and observations are to be recorded directly in your lab notebook. Never record data or observations anywhere except in your notebook.
  8. You will turn in the carbon copy of the raw data at the end of each formal lab exercise, retaining the notebook for yourself. Additional details will be provided on the first day of lab and posted on Blackboard.

Outline of Notebook Format

Generally, you will find that each experiment can be organized in your notebook under the headings: Title; References; Purpose; Apparatus, Equipment and/or Chemicals; Procedure and Observations; Data and Results (includes sample calculations); Discussion; and Conclusion. Some of these terms are explained as follows:

References:

This states the source, including page numbers, of the information used to perform the experiment, usually the laboratory manual. Your textbook and other sources also should be read and referred to.

 

Purpose:

This should be a brief simple statement of the intent of the experiment. It should be written so that later, when you reread it, you have no difficulty remembering the experiment. Usually, the manual contains such a statement and you may choose to adapt that wording.

Apparatus, Equipment, Chemicals:

All important materials used in the experiment should be listed. If any new equipment is used, it should be described with a drawing which includes essential features such as approximate dimensions and relative locations. If apparatus described previously is used, the original description should be cited, i.e., “p. 37”.

Procedure and Observations:

This section is much like a diary and should describe what you did and what you saw. The procedural steps should be written, as a part of the pre-lab, prior to entering the lab. The observations should be entered in the notebook at the time the operation is performed, not before. Do not simply write the same instructions which are in the manual.

Number the steps so that it is easy for you to keep track of where you are in the experiment. Make sure that the steps are clear and easy for you to understand (no praragraphs). Do not write in first person.

Never record on a pad of paper or on scrap paper and then transcribe into the notebook. Always record directly in the notebook. The notebook is not the spot for polished writing. Instead, be as detailed as possible so that someone else can duplicate what you have done by reading your account. This means including contradictions and failures. If there is a conflict, enter a description of it rather than omitting it. The negative results may be important later for figuring out what happened.

Data should be entered directly into the lab notebook as it is gathered before performing calculations.

Leave adequate room for observations and, where necessary, calculations. Do not write in a jumbled/congested and disorganized manner. Refer to the grading rubric (see Blackboard).

Data and Results:

This section must begin on a new page, titled “Data and Results”. You may wish to model your Data Sheet on the sample report sheets found in your lab manual. The quantitative data, as well as qualitative data, is to be organized, neat, and recorded with the appropriate significant figures and units. Data must be written legibly.

Discussion:

In this section, you will address the discussion questions from the lab manual or posted on Blackboard. In addition, your individual experimental outcomes, data analysis and error analysis will be addressed and discussed as it relates to outcomes for the entire class. Proper sentence structure and grammar are to be used at all times.

Conclusion:

In this section you should record any thoughts or conclusions about your experiment. You may offer explanations for low yields for example, or for improvements in procedure or you may comment on the experiment itself.

 

***A sample lab report from a previous semester is posted on Blackboard. The report is to be used as a guide—-not for plagiarism purposes.

 

word image 857

“A contemporary instance of the importance of laboratory notebooks is provided by the case of Gordon Gould, who as a young physicist filed an application for a basic laser patent in 1959. Gould failed to get the patent, which was awarded instead in 1960 to Charles Townes and Arthur Schawlow. Gould went to court, claiming he was the true inventor. His challenge was based in part on his research notebook which showed, among other items, a sketch, a statement of the main idea, and a derivation of the acronym LASER–Light Amplification by Stimulated Emission of Radiation.

“In October 1977, after a series of litigated oppositions, Gould was granted a patent for optically pumped laser amplifiers. The world market has been estimated at between 100 million and 200 million dollars.

“As this example suggests, well-kept notebooks are valuable documents. They provide complete, accurate records of ongoing work. In the event of litigation or contests for patent rights, they are submitted as evidence. They serve the important role of corroboration should the researcher or inventor have to prove origin or substantiate statements and conclusions. They are valuable documents to validate a company’s claims to funds spent for research, particularly in support of tax deductions.

“The uses of laboratory notebooks are not limited to legal issues. They are vehicles for organizing and focusing the thinking of the writer, as well as being receptacles for detailed procedural information that might not be available in highly compressed journal articles. Finally, they may serve not only the researcher or inventor but also the public. If properly maintained, they are a record of success and failure, a safeguard against error and carelessness in such important areas as the testing of drugs and chemicals.”[1]

word image 858

 

CHEM 213 / CHEM 272 EXPERIMENTS

 

 

word image 859

.

 

 

Experiment 1 PENNIES AND GOOD SENSE

 

 

 

word image 860 word image 861 word image 862 word image 863

 

Adapted from Reference: “An Experiment in Thinking Scientifically”, D. J. Sardella, J. Chem. Ed. 69, 933 (1992).

Purpose: The purpose of this exercise is to determine experimentally the variation in mass of pennies both within a single mint year and over a range of years.

Equipment: Electronic balance, pennies

word image 864Procedure and Observations: Obtain a vial of pennies from your instructor and record the mint year. Using a single balance, determine the precise mass (in grams) of each of your pennies (up to a maximum of 10) by placing them on the balance and weighing them using a weigh boat. Make sure to zero the balance with the weigh boat before measuring the mass of each penny. Record all your data in your laboratory notebook as you weigh the pennies. Be sure to record the mass to the nearest 0.001 g, which is the limit of precision of the balance. Make sure to observe the penny’s condition – corroded, dirty, etc.

Data: After you have recorded all the masses, enter your own data into the appropriate Web-based data entry form, using one of the computers in the lab. Setup a table in your lab notebook that has columns for masses, deviations, and deviations squared. Calculate the average mass and standard deviation for the mint year of pennies assigned by your lab instructor. Verify the calculated standard deviation value with your lab instructor prior to entering data into the database.

Results: Outside of class, retrieve the class data in spreadsheet form from: http://chem.gmu.edu/results/. Use Excel to calculate the average mass and the standard deviation for each mint year, using the AVERAGE() and STDEV() spreadsheet functions. Make sure the precision of the calculations matches the precision of the data. Finally, calculate the standard deviation for ALL the pennies, considered as a single large sample.

Use Excel to prepare a scatter plot (select CHART, then options) showing the average mass vs. year on the primary y-axis and standard deviation vs. mint year on the secondary y-axis. A sample is provided on the next page. Row and column labels should always be printed for each Excel sheet.

Examples of a data spreadsheet and charts are shown following this page.

 

Discussion: In your discussion section address the following questions:

 

  • Describe the trend observed in the variation of average mass with mint year.
  • Describe the trend observed in the variation of the standard deviation with mint year.
  • How does the standard deviation among the pennies in a given year compare with the standard deviation among all the pennies (all years)? What can you conclude from this?
  • Give a possible explanation for these trends and discuss how your observations support this explanation. Include a discussion of the contribution of the density of each (look it up for each metal) and the composition of each penny to the total mass of the pennies.

 

 

 

 

 

Chemistry 21

3

A1 Measurements I: Sample Sheet of Data and Results

2

 

 

 

 

 

 

 

Avg.

Mint mass1 mass2 mass3 mass4 mass5 mass6 mass7 mass8 mass9 mass10 Mass Std. year number (grams) (grams) (grams) (grams) (grams) (grams) (grams) (grams) (grams) (grams) (grams) Dev.

1971

 

9

 

3.041

 

3.048

 

3.064

 

3.09

0

 

3.112

 

2.966

 

3.083

 

3.089

 

3.108

 

3.067

 

0.045

 

1972

 

6

 

3.088

 

3.029

 

3.14

 

3.046

 

3.075

 

3.09

0

 

3.078

 

0.039

 

  1. 10 3.132 3.072 3.105 3.090 3.039 3.110 3.096 3.113 3.128 3.084 3.097 0.028
  2. 10 3.089 3.098 3.043 3.067 3.032 3.053 3.098 3.092 3.123 3.078 3.077 0.028
  3. 7 3.111 3.155 3.084 3.098 3.127 3.053 3.080 3.101 0.034

3.049

 

  1. 8 3.080 3.072 3.011 3.085 3.094 3.128 3.090 3.076 0.034
  2. 10 3.143 3.068 3.056 3.131 3.107 3.074 3.117 3.054 3.093 3.102 3.095 0.031
  3. 10 3.095 3.130 3.135 3.073 3.056 3.053 3.117 3.091 3.063 3.126 3.094 0.032
  4. 10 3.094 3.096 3.078 3.094 3.058 3.131 3.076 3.065 3.124 3.094 3.091 0.023
  5. 10 3.072 3.083 3.100 2.999 3.083 3.074 3.108 3.102 3.115 3.064 3.080 0.033
  6. 10 3.020 3.092 3.040 3.086 3.088 3.062 3.028 3.105 3.084 3.055 3.066 0.029
  7. 10 2.509 2.517 2.586 2.535 2.543 2.519 2.517 2.496 2.514 2.514 2.525 0.025
  8. 8 2.545 2.550 2.492 2.486 2.527 2.539 2.580 2.518 2.530 0.031
  9. 10 2.529 2.529 2.553 2.530 2.464 2.532 2.522 2.582 2.519 2.525 2.529 0.029 1985 10 2.534 2.579 2.507 2.575 2.525 2.501 2.502 2.543 2.492 2.511 2.527 0.031
  10. 10 2.485 2.561 2.490 2.520 2.569 2.494 2.614 2.482 2.557 2.560 2.533 0.045
  11. 10 2.511 2.545 2.468 2.508 2.503 2.496 2.521 2.465 2.485 2.500 2.500 0.024
  12. 10 2.544 2.504 2.479 2.470 2.516 2.473 2.505 2.492 2.518 2.486 2.499 0.023

1990 10 2.485 2.502 2.539 2.505 2.458 2.515 2.487 2.519 2.499 2.498 2.501 0.022

Shaded areas indicate where measurements were not performed.

 

word image 3207

Experiment 2 MEASUREMENTS

 

 

 

word image 865 word image 866 word image 867 word image 868

In all sciences, careful measurements of materials and conditions are essential. There must be well-understood, consistent methods of describing the weight, volume, temperature, and other characteristics of a sample or method that is investigated. The metric system is the accepted system for scientific measurements worldwide and will be employed throughout your chemistry lecture and laboratory. Tables listing the relationships between the English and metric systems of measurement can be found in your textbook and other reference works. You should be familiar with the metric prefixes and with the units used in measuring mass, volume, temperature, and distance.

word image 869 It is never possible to make an exact measurement because no matter how sophisticated the measuring instrument or how carefully the measurement is made, errors cannot be avoided. We distinguish between two categories of experimental errors: systematic and random. It is possible to eliminate or correct systematic errors; we cannot eliminate random errors. Random errors are the errors which are inherent in every physical measurement and are beyond the observer’s control. Systematic errors often are avoidable and may arise from faulty equipment or poor technique. (However, a systematic error is not the same as a mistake or a blunder.) In order to balance random errors and to more closely approximate the accepted or true value, at least two, and preferably three or more, measurements should be made. The best value is usually the average of the individual results (although often the question arises whether to include a result that seems out of line with the others). The arithmetic mean, or average, of this set of results is calculated by the equation:

Mean  Average  xave  [x1 x2 x3 … xi ]n

where X1,2,3,…,i indicates the measured values in the sample and n is the number of measurements in the sample.

The error of a measurement is the numerical difference between the observed value and the “true” or accepted value. The smaller the error, the more accurate the observed value. Thus, accuracy is the nearness of a result to the true value. If the true or accepted value is known, the average and/or individual values can be compared to it and the accuracy judged. A useful comparison is the percent error:

 

accepted value observed value

% Error 100

accepted value

Because the “true” value is seldom known (and hence the accuracy), we often rely on the precision of the data as an indication of its accuracy. Precision is the closeness or agreement of a set of results among themselves. It is thus an indication of the reproducibility or reliability of the measurements. Precise results may be accurate or inaccurate; that is, measurements of high precision, made with a faulty measuring instrument, can be wrong.

We often speak of the uncertainty of the measurement using a particular piece of equipment. Since no measurement, or derived quantity, can be exactly determined, the indication of the uncertainty of a measurement is expressed by the number of significant figures used in recording its magnitude. All certain digits are recorded, while the final digit is significant but estimated, or uncertain. (Consult your textbook for a discussion on the correct use of significant figures.) For example, some electronic balances weigh an object with an uncertainty of + 0.001 g (plus or minus 1 milligram), while other, less precise balances weigh only to + 0.01 g (plus or minus 1 centigram). If the digital display on a milligram electronic balance reads 10.382 g, then the sample’s maximum mass is 10.383 g and its minimum mass is 10.381 g (10.382 + 0.001 g). Therefore, the measurement and its uncertainty indicate the range a data value may have because of the need to estimate the last, uncertain digit. If the final digit of a display is a zero, such as 8.410 g, the zero must be recorded to indicate that the uncertainty is in the milligram range. Dropping the zero to give 8.41 g would indicate an uncertainty of only 0.01 g.

For equipment which does not display a digital value, estimate the uncertainty by how well you can read the graduation markings on the equipment. In many cases, you can assume that the last digit is uncertain by + 1 unit. For example, suppose the scale below represents graduations of 1 cm and you have determined the length of an object to be at the mark L. You are certain that the length is greater than 2 cm, but you are uncertain about its exact location between 2 and 3 cm; that is, you are uncertain about the final digit. Mentally divide the distance between 2 and 3 cm into tenths (0.1), and then if you estimate the location to be about 7/10’s along the line, the length is 2.7 + 0.1 cm.

1 2 3 4 5

^

L

The scale below is similar to the one above, but suppose this one is further subdivided into millimeters.

 

1 2 3 4 5

^

L

 

Now as you look at the mark L, you are certain that it is at least 3.5 cm, but again, you are uncertain about the final digit. Mentally divide the distance between 3.5 and 3.6 cm into tenths, (0.01 cm = 0.1 mm) and estimate the position of L. Your answer could be 3.52 cm but others might decide 3.51 or 3.53 cm. The measurement you would record is 3.52 + 0.01 cm. Notice that the closer the graduations, the more difficult it is to use an uncertainty of + 1 unit. An easier division to estimate might have been 0.2 or 0.5 cm. With some practice, you can decide what is the most appropriate division.

Relative uncertainty is the ratio of the uncertainty to the magnitude of the measurement.

It is frequently expressed as a percent.

uncertainty

% relative uncertainty  100

magnitude of measurement

Notice that with a constant uncertainty, the relative uncertainty becomes smaller as the magnitude of the measurement becomes larger.

The precision of a group of measurements is often expressed in terms of deviation (d). The deviation of a result is the numerical difference between an experimental value and the average value, (xi — xavg). If the deviations for a set of results are small, as compared to the average value, then the measurement is precise. Although we cannot always express the accuracy of a measurement (because we may not know the true value), we can always express the precision of several measurements.

In the statistical treatment of experimental data, a common measure of precision is the standard deviation, x, which characterizes the average uncertainty in the individual measurements. The standard deviation for a small sample is the “square root of the average of the squared deviations” and can be calculated as follows:

word image 870

where the series is the sum of the squared deviations, di2 =(xi — xavg)2, and n is the number of items in the sample. (For theoretical reasons, the value [n-1] is used rather than n.) The standard deviation, x, is expressed in the same units as the data. (See the sample calculation.) A small standard deviation indicates high precision and a large standard deviation indicates low precision. Generally, high precision indicates high accuracy, though this is not necessarily so. For example, when a constant procedural error is repeated in every measurement, high precision may be observed even though the result is not an accurate determination of the true value.

If we assume that most measurements are corrected to eliminate systematic errors and that random errors are small and as likely to be positive as to be negative, then a graph of the distribution of many measurements approaches a bell-shaped curve. (See Figure 1.) Such a graph can be generated by plotting the value of a measurement versus the number of times (frequency) that particular value is obtained.

The average value is found at the peak maximum and is considered the most probable value (more likely to occur frequently). Those values which depart widely from the average occur less frequently than those for which the deviation is small. If most of the measured values cluster closely around the mean, the curve will be narrow and steep and the measurements will have a high precision. If the measured values are widely dispersed from the mean, the curve will be broad and the measurements imprecise. Inspecting Figure 1, you will notice that although both curves have the same mean (and thus the same accuracy), one set of results is much more precise than the other. This is a very effective depiction of the difference between accuracy and precision.

 

Figure 1

 

Number of times the value occurred

 

Value

 

Often one value in a group of results differs significantly from the others. In such cases, it is likely that some inadvertent error was made in a measurement. There is a simple rule that you may use in deciding whether a value should be discarded. If the deviation of the suspected value from the mean (determined without the suspected value) is equal to three times (or more) the average deviation from the mean of the other values, the suspected value should be discarded.

 

***Some of the common laboratory equipment you will use this year appear in Figure 2. Information on expected accuracy and uncertainty is provided in Table 1.

 

Figure 2

word image 3208

 

 

 

TABLE 1

 

 

Instrument

 

Graduation

 

Uncertainty

 

Top-loading electronic balance centigram

 

0.01 g

 

0.01 g

milligram

 

0.001 g

 

0.001 g

 

Graduated cylinder

10-mL

 

0.1 mL

 

 

100-mL

 

1 mL

 

 

 

Buret, 50-mL

0.1 mL

0.05 mL

Volumetric flask

not graduated —

 

25-mL

fill to mark

 

50-mL

100-mL

 

Pipet, volumetric,

 

 

 

not graduated —

 

 

 

 

10-mL

 

fill to mark

 

 

 

110° thermometer

1° C

0.2° C

 

Mercury barometer

 

0.1 mm

 

0.2 mm

 

Calculating a Standard Deviation

Measured Value

Deviation

Deviation Squared

xi

di=(xi – xavg)

di2=(xi – xavg)2

5.21

+0.01

0.0001

5.18

– 0.02

0.0004

5.23

+0.03

0.0009

5.20

0.00

0

5.17

– 0.03

0.0009

Sum =

 

25.99

 

 

0.0023

 

𝐴𝑣𝑒𝑟𝑎𝑔𝑒

=

 

𝑥

̅

=

𝑥

𝑎𝑣𝑔

 

=

 

(

𝑥

1

+

𝑥

𝑖

)

𝑛

=

25

.

99

5

=

5

.

20

 

𝑆𝑡𝑑

.

𝐷𝑒𝑣

.

=

𝜎

𝑥

=

(

𝑑

1

2

+

𝑑

2

2

+

𝑑

3

2

+

.

+

𝑑

𝑖

2

)

(

𝑛

1

)

=

(

0

.

0023

)

(

5

1

)

=

0

.

02

 

Problems (show work)
  1. Explain why it is correct to speak of the uncertainty of a single result but not the precision of a single result. Is it correct to speak of the accuracy of a single result?
  2. What are the maximum and minimum values for a mass which has been reported as 7.23 ± 0.05 g?
  3. Show that for an accuracy of 1% in weighing a 50 g sample, a balance which is uncertain to only

0.5 g is required.

  1. Calculate the percent error in a molecular weight determination if the experiment yielded values of 121.5, 122.3, 121.9 and the actual molecular weight is 122.1.
  2. The following volumes were measured in a 100-mL graduated cylinder: 74.6, 75.2, 73.7, 74.2,

75.0, 73.9. Calculate the average and the standard deviation for the measurements.

  1. If your calculator has a built-in statistics program, verify the calculations in the Example and in Problem #6. .
  2. Suppose the following percentages were determined for the amount of water in a crystal hydrate: 36.20%, 36.52%, 35.50%, and 35.95%. Determine if one of the values is suspect and may be discarded.
Answers to Selected Problems: #4. 0.2% #5. 74.4, 0.6

THE EXPERIMENT

It is important that you learn how to use common equipment correctly in order that your measurements in the laboratory are as accurate and precise as possible. The validity of your experiment depends on the reliability of your measurements. You will use the equipment from this laboratory exercise to perform your experiments throughout the year. At the beginning of the laboratory period, your instructor will demonstrate the proper use and function of this equipment. During the experiment, you will be expected to acquire certain skills and techniques as you use the equipment.

Today you will practice making mass and volume measurements. Mass measurements will be made on an electronic balance and their precision assessed. All of your mass measurements should be made on the same balance. Volume will be measured with a variety of glassware — graduated cylinder, pipet, and buret. The accuracy of these devices will be determined by comparing a measured volume with the calculated volume. The conversion factor which relates mass and volume, density, is expressed as mass per unit volume. At room temperature, one gram of water occupies approximately one milliliter of volume (that is, the density of water = 1 g/mL).

PROCEDURE

You will need a calculator for this experiment. It may be helpful to review the built-in standard deviation program on your calculator. Although such a program is not necessary for the experiment, it will make your calculations of standard deviation easier. Also, make sure you understand rounding of numbers and the determination of significant figures as explained in your textbook.

Before you begin, place approximately 250 mL of distilled water in a large beaker and allow it to come to room temperature. All of your volume measurements should be done with this water. Place a thermometer in the water and record its temperature. Record the temperature of the air in the room.

Measuring Mass

  1. Weigh a dry 25-mL graduated cylinder on an electronic balance. Record the uncertainty of the measured mass. Repeat twice, zeroing the balance before each trial. Determine the average mass. Calculate the standard deviation.
  2. Weigh a dry 150-mL beaker on an electronic balance. Record the uncertainty, then determine the average mass and the standard deviation.

Assess the precision of the balance. Based on your triplicate mass measurements, how many times do you think you should measure mass in future experiments?

Assessing the Accuracy of Glassware

Your results from the preceding section probably showed that mass measurements using the same balance are very precise (and presumably accurate because the balances have been calibrated with known masses). We next want to answer the question “which glassware most accurately measures the volume of a given amount of water?” To determine this, you will use the mass of water as a basis for comparing and judging the accuracy of the volumes of water you measured. The conversion factor which relates mass and volume is the density. The density of water at various temperatures has been accurately determined in other laboratories. Alternatively, you may use the formula:

density = 1.0028 − (0.000228)T

where T is in oC and the density is in g/mL.

  1. Into the same graduated cylinder you used in the Measuring Mass section above, add about 10 mL of distilled water at room temperature. Do not attempt to adjust the volume in the graduated cylinder to exactly 10.0 mL; this may introduce error. Record the volume of water in the graduated cylinder (this is the measured volume; it should not be exactly 10 mL.). Weigh the graduated cylinder containing the water and determine the mass of water in the graduated cylinder.
  2. Set up a clean buret and rinse with distilled water. Fill the buret with distilled water and drain enough water through the tip to flush out any air bubbles. Bring the water level to below the 0.0-mL mark. Record the initial volume reading. Do not attempt to adjust the initial volume to the 0.00 mL mark. This is unnecessary and may introduce error. Then drain about 10 mL from the buret into the 250-mL dry beaker you used in the previous section. Record the final volume reading from the buret to obtain the measured volume of water delivered from the buret. Weigh the beaker and water to determine the mass of water delivered from the buret.

 

From the known density of water at today’s room temperature and the measured mass of water delivered from each piece of glassware, determine the calculated volume of water delivered from each. Finally, calculate the error and the percent error in the measured volume, taking the calculated volume as the accepted/true value.

mass (g)

Volume (mL) 

density (g/mL)

Compare the accuracy of the three types of glassware. What are the advantages and disadvantages of the three types of glassware?

 

 

Temperature (°C)

 

Density of water (g/mL)

20

 

0.99823

21

 

0.99800

22

 

0.99777

23

 

0.99767

24

 

0.99733

25

 

0.99708

 

 

___________________________________________________________________________________________________

MEASUREMENTS

___________________________________________________________________________________________________

MEASURING MASS

Mass of 10-mL graduated

cylinder __________ (g) __________ (g) __________ (g)

 

Avg. Mass __________ (g) Uncertainty + _______

 

Std. Deviation _______

 

 

 

 

 

Mass of 150-mL

beaker __________ (g) __________ (g) __________ (g)

 

Avg. Mass __________ (g) Uncertainty + _______

 

Std. Deviation _______

 

 

 

 

***Sample data table provided above. The table in your lab notebook should follow the same format.

Assessment of Precision:

GLASSWARE ACCURACY

Temperature of room air:

__________

 

 

Temperature of water:

__________

 

 

Density of water (from table or calculation)

___________

 

 

 

 

 

 

REQUIRED Sample Calculations for the following:

 

Graduated Cylinder:

 

Measured Vol. of H2O

Mass of Cylinder + H2O

Mass of H2O

 

 

 

Calculated Vol. of H2O

Error

Percent Error

 

 

Buret:

 

 

Initial Vol. in Buret

Final Vol. in Buret

Measured Vol. of H2O

 

 

 

Mass of Beaker + H2O

 

Mass of H2O

Percent Error

Calculated Vol. of H2O

Error

Percent Error

Assessment of Accuracy:

 

Experiment 3 QUALITATIVE ANALYSIS OF CATIONS

 

 

 

 

word image 871 word image 872 word image 873 word image 874 word image 875

 

Introduction:

This experiment will analyze known and unknown solutions for the presence of specific cations. The general approach to finding out what ions are in a solution is to test for the presence of each possible component by adding a reagent that will cause that component, if present, to react in a certain way. This method involves a series of tests, one for each component, carried out on separate samples of solution. Difficulty sometimes arises, particularly in complex mixtures, because one of the species may interfere with the analytical test for another. Although interferences are common, many ions in mixtures can usually be identified by simple tests.

word image 876 In this experiment, you will analyze unknown solutions that may contain one or more of the following ions in solution:

Fe3+, Ni2+, Mg2+, Cu2+

First, you will perform various tests to detect the presence of individual ions. Once you have observed these specific reactions, you will perform the same tests on a series of unknown solutions.

 

Safety Precautions:

 

In some of the tests, you will be required to use fairly concentrated acids and bases. When in contact with skin, most of these chemicals can cause severe burns if not removed promptly. Wear goggles when working with any of the reagents required in this experiment.

Experimental Set-Up:

For each experimental procedure, arrange four test tubes according to the picture below, so that the test tubes are organized like the data tables. The first row of test tubes corresponds to row 1 on the data tables. The first column of the test tubes corresponds to column A on the data tables.

Procedure for Cation Tests:

Controls – First, you will test known solutions and observe the results (rows 1-4). For each cation there will be a unique set of positive reactions that indicate the presence of that cation.

 

The test solutions for Cation are sodium iodide in acetone (NaI), sodium ferrocyanide (Na4Fe(CN)6), dimethylglyoxime in ethanol/water solution, and sodium hydroxide (NaOH).

  1. Place 20 drops of the first cation solution (Fe(NO3)3 ) in the first 4 test tubes.
  2. Repeat for test tubes 5-16 using the appropriate cation solution.
  3. Place 4 drops of the first test solution (NaI) in the first test tube of each cation solution.
  4. Repeat for test tubes 5-16 using the appropriate test solution.
  5. Observe and record any reactions in the data table. Be sure to record the color and amount of any precipitates.

Unknowns

  1. Identify the cations in the unknown solutions in the vial.
  2. Repeat steps 3-5 for the unknown vial to identify unknown cation solution.
  3. Clean Up – carefully pour all of the solutions from the test tube into the waste container, then clean and dry the test tubes. If needed, rinse the test tube with soap or acetone to remove residue.

 

Name_______________________ Section___________ Lab Partner_____________

Data Tables:

Cation Identification (descriptions will vary)

 

 

A

B

D

E

 

 

7.5 % NaI

0.25M Na4Fe(CN)6

0.25 M

Dimethylglyoxime

1 M NaOH

1

0.25 M

Fe(NO3)3

 

 

 

 

2

0.1 M

Ni(NO3)2

 

 

 

 

3

0.1 M Mg(NO3)2

 

 

 

 

4

0.1 M

Cu(NO3)2

 

 

 

 

 

  1. Balance and classify the following reactions: combination (C), decomposition (D), single

replacement (SR), double replacement (DR) or combustion (CB).

 

Classification

a. ____C3H8 (g) + ____O2 (g) → ____CO2 (g) + ____H2O (l)

____________

b. ____K (s) + ____H2O (l) → ____KOH (aq) + ___H2 (g)

____________

c. ____MgSO4 (aq) + ____Na2CO3 (aq) → ____MgCO3 (s) + ____Na2SO4 (aq)

____________

    1. ____Cl2 (g) + ____NaI (s) → ____NaCl (s) + ____I2 (s)
    2. ____NaNO3 (s) → ____NaNO2 (s) + ____O2 (g) ________
    3. ____H2 (g) + ____Cl2 (g) → ____HCl (g) ________
    4. ____Ca(C2H3O2)2 (aq) + ____H3PO4 (aq) → ____Ca3(PO4)2 (s) + ____HC2H3O2 (aq) __________
    5. ____H2O2 (l) → ____H2O (l) + ____O2 (g) _________
    6. ____Na (s) + ____S (s) → ____Na2S (s) _________
    7. ____BaCl2 (aq) + ____K2CrO4 (aq) → ____BaCrO4 (s) + ____KCl (aq) ________

 

  1. Balance each of the following equations. Then write the total (complete) ionic equation for each reaction. Finally write the net ionic equation for each reaction. Remember: insoluble substances are not present as separate ions in solution.
    1. ____Pb(NO3)2 (aq) + ____NaOH (aq) → ____Pb(OH)2 (s) + ____NaNO3 (aq)

Complete Ionic equation:

Net ionic equation:

    1. ____Na3PO4 (aq) + ____Ca(NO3)2 (aq) → ____Ca3(PO4)2 (s) + ____NaNO3 (aq)

Complete Ionic equation:

Net ionic equation:

    1. ____NaCl (aq) + ____AgNO3 (aq) → ____AgCl (s) + ____NaNO3 (aq)

Complete Ionic equation:

Net ionic equation:

    1. ____Na2SO4 (aq) + ____AgNO3 (aq) → ____Ag2SO4 (s) + ____NaNO3 (aq)

Complete Ionic equation:

Net ionic equation:

    1. ____Na2CO3 (aq) + ____BaCl2 (aq) → ____BaCO3 (s) + ____NaCl (aq)

Complete Ionic equation:

Net ionic equation:

 

    1. ____Na2CO3 (aq) + ____HCl (aq) → ____NaCl (aq) + ____H2O (l) + ____CO2 (g)

Complete Ionic equation:

Net ionic equation:

 

 

 

Experiment 4 DENSITY OF SOLUTIONS

 

 

 

word image 877 word image 878 word image 879 word image 880

 

Density is a physical property of a substance; it relates the mass of the substance to the volume it occupies and it is thus characteristic of the substance. Densities can be measured for gases, liquids, and solids and are expressed as mass per unit volume. Gas density is usually expressed as g/liter, while for the more dense liquids, the density is expressed as g/mL. Densities of many pure substances have been determined and you can look them up in reference books. For example, the density of water at 4°C is 1.000 g/mL (density changes somewhat with temperature, so it is necessary to specify the temperature to which the density refers).

word image 881 To understand how the densities of substances can vary widely, imagine two trucks, each with the same volume of cargo space. One truck is filled with lead shot, the other with packing peanuts. Although the volumes occupied are identical, the truck filled with lead will be much heavier than the truck filled with packing peanuts. Thus, the density of lead is much greater than the density of packing peanuts since lead has the larger mass-to-volume ratio.

Solutions are homogeneous mixtures of two or more substances in which the boundaries of the component substances can no longer be identified. A common solution is that of a solid-in-liquid, such as sugar in water. Before preparing the solution, you can see the solid and you can see the liquid, that is, you can identify their boundaries. After mixing and stirring them together, you cannot see the sugar anymore, but it is still present, each sugar molecule separated from the others and surrounded by water molecules. If you were to drink this solution, you would know the sugar had not vanished. The substance which is present in minor amount in the solution is the solute and the substance which is present in greater amount is the solvent. The solute is said to dissolve in, or be soluble in the solvent.

A solution itself can be either solid (a metal alloy such as brass), liquid (the sugar-water example above), or gaseous (the air we breathe); solutions can be composed of any combination of solid, liquid, and/or gaseous substances, although some combinations are rare. In the chemistry laboratory, we most often encounter solid-in-liquid or liquid-in-liquid solutions.

Densities of mixtures, as well as of pure substances, can be measured. Knowledge of the density of a solution not only provides useful information about its composition, but can also aid in determining the concentration of an unknown solution. Today, it is a well known fact and accepted principle that the density of a solution varies proportionately with the mass of the solute present.

Archimedes’ cry of “eureka” supposedly occurred when he realized that pure gold and adulterated gold (a solid solution) have different densities. Those who would pay for pure gold wanted to know that it was, in fact, pure.

Considering a solid-in-liquid solution, such as the sugar water example above, you can see how the density should increase as the ratio of sugar to water increases. (It is often the case, as it is here, that the solid solute has a greater density than the liquid solvent.) Suppose you have a liter of pure water, which has a density of about 1 g/mL (what is the mass of water present?). If you dissolve a certain amount of sugar in the water, and then weigh a liter of the resulting solution, there will be a greater mass of solution present in that one liter volume than there was of the pure water, hence a greater density for the solution. As you add more and more sugar, each time you weigh a liter of the solution it will have an increasingly greater mass and greater density.

A convenient aspect of the solution density concept is that the density of a solution and its concentration are linearly related. Concentration is expressed in many different units (e.g.

molarity, weight %, volume %), but it refers to the ratio of the amount of solute to the amount of solution (rarely, it refers to the ratio of solute to solvent, an important distinction). Whenever the solute/solvent ratio changes, the concentration of the solution changes and there will be a proportional change in the density of the solution. Look at the graph to see the linear relationship shown as line A.

A

#

0

.

9

9

1

.

0

1

1

.

0

0

1

.

0

2

1

.

0

3

1

.

0

4

0

1

2

3

4

5

6

D

e

n

s

i

t

y

(

g

/

m

L

)

 

 

 

 

 

 

 

 

 

Concentration (M)

In this graph, density versus concentration is plotted. If each mark on either axis represents one unit of density or concentration, then by inspection of this example, you can see that each two-unit change in concentration results in a one-unit change in density.

How could you exploit this relationship if you do not know the concentration of a solution you have been given? You could prepare a series of solutions using the same solute and solvent, with known, but varying concentrations; these are called standard solutions. After determining the density (by measuring the mass and the volume) of each standard solution, you would plot a graph of density versus concentration. If you then measure the density of the solution of unknown concentration and mark that density on the line of the graph, the corresponding concentration you read from the graph will be that of the unknown solution. The graph illustrates this.

Suppose the density you measure is denoted by the # symbol on the density axis. If you then draw a horizontal line (—–) across the graph until you come to the data line A, and then a draw a vertical line down from that point, the concentration of the solution has the value where the vertical line intersects the concentration axis.

THE EXPERIMENT

In this experiment, you will put together the principles and concepts which were discussed above. Although determining the unknown concentration of a solution is the ultimate goal, efficiently preparing the series of standard solutions and accurately measuring their densities will consume most of the laboratory period.

You will prepare the standard solutions by the method of dilution. This procedure requires taking a known amount of a concentrated stock solution to which is added pure solvent. You will begin by delivering, from a buret into a volumetric flask, a known volume (Vconc) of the stock solution with known concentration (Cconc); then you will add pure solvent until you reach a new known volume

(Vdil).

Calculating the concentrations of the diluted solutions is simple. No matter what the choice of concentration units, all dilutions can be represented by the equation:

(Volume)dil * (Concentration)dil = (Volume)conc * (Concentration)conc

or

Vdil * Cdil = Vconc * Cconc

where the subscript “dil” refers to the dilute solution and “conc” refers to the more concentrated solution. Rearranging the above equation to solve, and substituting known numerical values, gives the concentration of the diluted solution (Cdil). The concentration units will be molarity (M), moles of solute/liter solution. In your measurements and in your calculations, remember to pay attention to uncertainties and significant figures.

As you prepare the graph of your results (whether on graph paper or electronically), keep in mind some features of good graphing technique: label each axis clearly, denoting units; choose unit divisions that can be easily subdivided; divide each axis so that the data fills the graph paper; assign the x- and y-axes so that you avoid drawing your data line with either a steep or a shallow slope; and draw the best straight line of the data points using a transparent straight edge.

Problems (show work)
  1. Calculate the mass of NaCl needed to prepare 100 mL of a 3.00-M sodium chloride solution.
  2. Calculate the volume of solution necessary to make 25 mL of a 2.00-M solution from the 3.00-M solution.
  3. Calculate the molarity of 25 mL of solution, which was diluted from 10 mL of a 2.0-M solution.

Answers to Selected Problems

1). 17.5 g 2). 16.7 mL 3). 0.80-M

PROCEDURE

The solid salt used to make your solutions will depend on what the stockroom has available. They will issue it as a 3.00-M standard solution. Your instructor will inform you of the identity of the salt. Since solid salts are convenient to use, the word “salt” will be used throughout the instructions.

Preliminary Calculations: As a preliminary step, calculate the volumes of 3.00-M stock salt solution required to prepare 10.00 mL each of: (a) 2-M, (b) 1-M and (c) 0.5-M salt solutions. These are “rough” calculated volumes to guide you when you make your actual dilutions, below. The result of the first calculation is shown in the table below.

Vconc = (Vdil  Cdil ) / Cconc

Vdil

 

Cdil

 

Cconc

Vconc

 

10 mL

 

 

 

 

2 M

 

 

3 M

 

 

 6.7 mL

 

1. Before beginning your density measurements, obtain a clean, dry 10.00-mL volumetric flask with stopper. Weigh the volumetric flask and stopper on the electronic balance. For all subsequent mass determinations in this experiment, use the same balance.

2. Dilutions
  1. Rinse your buret with a small quantity (~5 mL) of the 3.00-M stock solution and then pour it out; add about 15 mL of the 3.00-M stock solution to the buret. Record the initial volume of solution in the buret.
  2. Deliver directly into a 10-mL volumetric flask a volume of 3.00-M salt solution which closely corresponds to (but is not exactly the same as) one of the preliminary calculations above.

Record the final volume of solution in the buret.

  1. Add distilled or deionized water until the volumetric flask is half to three-quarters full, then swirl and mix the contents of the flask. Continue adding water, dropwise if necessary, until the bottom of the meniscus just touches the calibration mark on the volumetric flask. Mix well. Measure the mass of the solution. After you have measured the mass of the solution (below), calculate the actual concentration of the solution in the flask.
3. Densities
  1. Weigh the volumetric flask containing the salt solution prepared in Step 2 above. From the mass of the solution and its volume, calculate the density.
  2. Empty and rinse the volumetric flask well with distilled water and shake out excess water. Using the same 10-mL volumetric flask, prepare each of the remaining diluted standard solutions using the same method as in Steps 2b and 2c above. Determine the density of each solution as in Step 3a.
  3. Determine the density of the 3.00-M stock solution. The buret is not required for this step; simply fill the flask to the mark.
  4. To compensate for any errors you might have made in preparing the solutions it is helpful to measure the density of water. This value, which will be the smallest density of all (i.e., it has no salt and so is 0.0-M), will be at the left edge of your graph and will help you decide how to draw the best straight line through your data points.
  5. In your lab notebook, plot a graph of density versus concentration. Review the principles of good graphs before you begin. Include a linear regression line, or draw the best straight line that fits the data points. This experiment is not performed in pairs and each student must estimate his/her own unknown concentration prior to leaving the lab. Your lab instructor will write your true/accepted unknown concentration value in your lab notebook and you may then enter the data into the computer’s database.
  6. Density and concentration of a solution of salt of unknown concentration:
    1. Obtain a sample of salt solution of unknown concentration from your lab instructor (unknown solutions prepared by Chem Preps Room).
    2. Remember to write down the unknown letter.
    3. Determine the density of the salt solution by transferring 10 mL of the solution into the clean, pre-weighed volumetric flask.
    4. Reweigh the volumetric flask and its contents.
    5. Determine the solution concentration based on the data on your graph.

Consider possible sources of error that could account for any difference in your value and the true density value. ______________________________________________________________________________________________________

DENSITY OF SOLUTIONS

______________________________________________________________________________________________________

  1. Mass of dry 10-mL volumetric flask with stopper ___________ (g)
  2. Calculations of Actual Concentrations of Diluted Solutions Vdil  Cdil = Vconc  Cconc

 

Volume of stock solution

in Buret

Volume of stock solution delivered

Concentration of dilute solution

Initial (mL)

Final (mL)

Vconc (mL)

Cdil (M)

 

 

 

 

 

 

 

 

  1. Determination of Densities

Solution Conc.

(Cdil)

Mass Container

& Solution

(g)

Mass

Container

(g)

Mass of

Solution

(g)

Volume

of solution

(mL)

Density of solution

(g/mL)

 

Stock

Solution

~3M

 

~2M

 

~1M

 

~0.5M

 

H2O

 

Unknown

 

 

 

 

 

 

Enter your actual concentrations; above values provided as a guide to organize your table.

 

DATA ANALYSIS SUPPLEMENT

Data: After you have entered all the masses and volumes, enter your own data into the appropriate Web-based data entry form, using a computer in the lab. In your measurements and calculations, remember to pay attention to uncertainties and significant figures.

Results: Outside of class, retrieve your data in spreadsheet form (along with that of the entire class) from: http://chem.gmu.edu/results/. Then use Excel to calculate the densities of the various diluted solutions, stock solution and water.

Your lab report should include the following in addition to the pages from your lab notebook:

  • An Excel sheet “Raw Data” showing the raw data for the entire class (Landscape orientation; appropriate column labels and units; sigfigs; student identifying information).
  • A 2nd Excel sheet “Class Results” showing the calculated results for each student in the class – each in a separate row. The columns should include:
    • the concentrations of the five solutions (including the distilled water).
    • the measured densities of the five solutions
    • the slope and intercept for the density vs. conc. plot. Use the Excel functions =SLOPE(y-range, x-range) and =INTERCEPT(y-range, x-range) o the measured density of the unknown o the calculated concentration of the unknown.
  • A 3rd Excel sheet “My Results” showing your own results, only. Be sure to include your unknown number at the top of this sheet! There will be six rows, one for each of the known solutions and one for the unknown. There should be two columns:
    • density (g/mL) o concentration (M)
  • An Excel chart of density vs. concentration for your own data only, based on the “My Results” sheet described above. Density (dependent variable) should be on the y-axis and concentration on the x-axis. The chart should have five points – one for each of the known solutions. Include a linear regression fit to your points and show the equation of the line on the graph; include the R2 value. Format your charts so that they look like examples on the following pages. Make sure to print graphs and tables in LANDSCAPE orientation.

 

Discussion: In your discussion section address the following questions:

  • Making the simplistic assumption that the dissolved NaCl(s) does not affect the volume of the solvent water, determine the constants m and b in the equation Molarity = mdensity + b that relates the NaCl molarity to the NaCl(aq) density. Take the density of water to be 1.00 g/mL and the molar mass of NaCl to be 58.5 g/mol.
  • How does your experimental value for the slope m and intercept b compare with the values derived from the simple model above?
  • What are the units for slope and intercept for the plot?

 

 

word image 3209 word image 3210

SAMPLE CALCULATIONS

Conc A = C_A (M)

 

=RawData!G39*RawData!B39/RawData!H39

Density A = D_A (g/mL)

=RawData!J40/RawData!K40

Slope

=SLOPE(G3:K3,B3:F3)

Intercept

=INTERCEPT(G3:K3,B3:F3)

[UNKNOWN]

=L3*M3+N3

%ERROR

=ABS((O3-P3)/P3)*100

 

word image 3211

word image 882

 

 

Experiment 5 EMPIRICAL FORMULA OF ZINC IODIDE

 

 

 

word image 883 word image 884 word image 885 word image 886 word image 887

 

One of the simplest type of chemical reactions is the combination of two pure elements to form a binary compound. The term “binary” refers to the number of different elements in the compound. A binary compound is composed of just two elements, although there may be more than one atom of an element in the formula unit of compound. For instance, copper metal reacts with sulfur to form the binary compound copper(I) sulfide, according to the chemical equation:

2 Cu + S Cu2S.

word image 888 Even though the reaction may appear to be simple and straightforward, the morphology of the starting materials and reaction conditions determine the actual stoichiometric resultant product. The above reaction generates the desired product if Cu (powder) reacts with molten sulfur or if copper is reacted with sulfur vapour. However, if Cu (pellets) are used as a starting material, then a much higher reaction temperature is required in order to form copper sulfide. In a likewise manner, it is also possible to produce copper(II) sulfide, CuS, and other stoichiometric copper sulfides based on the reaction route of choice.

The quantities (masses) of the substances consumed in such a reaction obey two important laws. The Law of Conservation of Mass states that the total mass is conserved in the reaction. In the example above, this means that the total masses of copper and sulfur that reacted must equal the mass of copper sulfide produced. The Law of Definite Proportions states that the ratio of the masses of the elements in a given compound is constant. Thus, the ratio of the masses of copper and sulfur that react will be a simple multiple of the ratio of the atomic weights of those elements. For instance, in the example above,

mass of Cu reacted atomic weight of Cu

 X mass of S reacted atomic weight of S

The multiple, X, is 2 in this case because copper combines with sulfur in the ratio of two atoms of copper for every one atom of sulfur. Another way of stating this is that copper combines with sulfur in the ratio of two moles of copper for every one mole of sulfur. This information is conveyed in the empirical formula for copper(I) sulfide, Cu2S, where the formula Cu2S is the conventional way of representing Cu2S1.

 

 

THE EXPERIMENT

The purpose of this experiment is to determine the empirical formula of the binary compound, zinc iodide, which is the product of the reaction of the elements zinc and iodine. This is done by preparing a sample of zinc iodide and measuring the reacting masses of the constituent elements.

Zn + I2 → ZnxIy

The x and y subscripts are to be experimentally determined. In actuality, the absolute values for x and y themselves cannot be determined directly, only their ratio y:x. Then, x and y are chosen as the smallest integers which will give the proper ratio. For example, suppose one finds experimentally that the ratio of y to x is 1.5. The smallest integers which will give this ratio are y = 3 and x = 2. The ratio must be expressed as a ratio of integers since atoms are indivisible and would result in Zn2I3 as the empirical formula.

In the reaction of Zn and I2, one of the two reactants will be present in excess. This means that there will be a “left-over”, unreacted quantity of this particular starting material when the reaction is complete. The other reactant (not present in excess) is called the limiting reagent. This reactant will be entirely consumed in the reaction; none will remain when the reaction is complete. The mass of the limiting reagent in the final product is the same as its starting mass, since all of it reacts. The mass of the reactant present in excess must be found indirectly by measuring the mass that is not consumed and then subtracting the value from its mass before the reaction began.

Example

Consider a reaction in which 1.132 g of copper and 0.620 g of bromine (Br2) were mixed and allowed to react to form the binary compound copper bromide, CuxBry. Suppose that when the reaction was complete, it was found that all of the bromine had reacted, i.e. bromine was the limiting reagent, and that 0.640 g of the copper remained unreacted (excess). The empirical formula for the copper bromide produced can now be determined.

First calculate the moles of each element which reacted to form product.

mass Br

moles Br2  2 0.620 g 0.00388 mol Br2

molecular wt. Br2 279.9 g/mol

Bromine exists naturally as the diatomic molecule Br−Br. We need to conceptually “break” the molecule into atoms so

moles of Br atoms = 2  (moles of Br2 molecules) = 0.00776 mol Br mass Cu consumed (1.132-0.640) g

moles Cu    0.00774 mol Cu

atomic wt. Cu 63.55 g/mol

The mole ratio of atoms in the product can now be determined:

moles Br atoms 0.00776

 1

moles Cu atoms 0.00774

Since the mole ratio is 1, the ratio of the integers in the empirical formula y:x is also 1 and the smallest integers giving this ratio are x = 1 and y = 1. Thus the empirical formula must be CuBr (short for Cu1Br1). When using your experimental data in the calculation, do not be alarmed if the final mole ratio obtained is not exactly an integer. It should be close to an integer however, and your number should be rounded to the nearest integer in order to generate a valid empirical formula.

THE EXPERIMENT

The quantities given below for the reactants are approximate in the sense that they are guidelines. It is not necessary (or even advantageous) to use precisely those quantities in order to obtain a successful outcome for the experiment. However, it is absolutely necessary that the precise quantities that are actually used be recorded in your laboratory notebook for use later in your calculations. For example, when directed to weigh out 2 g of zinc for the reaction, this does not mean that you must use exactly 2.000 g. A mass within a few percent of this is just as good or better, i.e. 1.852 grams. However, you must record your exact mass, for instance 1.852, and use this number in your calculations. One advantage to weighing out a quantity of zinc that is within a few percent of 2 g (not exactly 2.000) is that this can be done much more quickly. Another advantage is that your are less likely to obtain a biased mass.

Safety Precautions

To avoid potential allergic reactions to iodine vapor, the experiment should be performed in the hood up to the point at which the iodine limiting reagent is consumed. Be careful to avoid inhalation of iodine vapor and avoid contact between the iodine and your skin.

Procedure
  1. Record the mass of a clean and dry porcelain evaporating dish (P.E.D). directly into your lab notebook; (a) in the data table.
  2. Weigh approximately 2 grams of zinc granules directly into the P.E.D. Record the mass of the P.E.D. + the zinc granules; (b) in the data table.
  3. Determine the mass of zinc granules by difference; (c) in the data table.
  4. Weigh the vial containing the iodine crystals (LEAVE THE CAP ON WHILE WEIGHING); (d) in the data table.
  5. Transfer the iodine crystals directly into the previously weighed P.E.D. that contains the zinc granules. (Do not put iodine directly on the balance pan.)
  6. Re-weigh the vial + cap assembly; record the mass directly into your lab notebook; (e) in the data table.
  7. Determine the mass of iodine crystals transferred to the P.E.D. by difference; (f) in the data table.
  8. Take the P.E.D. containing the iodine to the fume hood. Observe the physical properties of the substance.
  9. Add 5 mL of distilled water to the P.E.D. containing the iodine crystals and swirl to partially dissolve the iodine, observing the color and appearance of the resulting solution. (Note that the change observed is not a chemical reaction, since the iodine has merely partially dissolved in the water.)
  10. In the fume hood, add a few drops of 6 M acetic acid solution to the iodine and water.
  11. Observe and record any evidence you see that a chemical change is occurring. Stir the solution periodically with a stirring rod. It usually takes 10 minutes for the reaction to go to completion (evidenced by a clear and colorless solution and granules).

The rest of the experiment may be performed outside the hood.

  1. Decant and discard the solution containing the zinc iodide; DO NOT DISCARD THE GRANULES from the P.E.D. Wash the excess zinc in the dish three times with 1 mL portions of distilled water acidified with a drop or two of the 6 M acetic acid solution.
  2. Dry the excess unreacted zinc by placing the P.E.D. on a hotplate (if a drying oven is not available) for at least 20 minutes.
  3. Once the zinc is completely dry, allow the dish to cool completely to room temperature.
  4. Weigh the P.E.D. containing the dried unreacted zinc; (g) in the data table.
  5. Determine, by difference, the mass of unreacted zinc; (h) in the data table.
  6. Determine the mass of zinc that reacted with the iodine; (i) in the data table.

 

Calculate the mole ratio of I:Zn and then determine the empirical formula.

Enter your empirical formula into your data table.

Pay attention to significant figures, using the appropriate number of them to reflect the precision of your mass measurements. Do not round-off until the final calculation of the empirical formula.

Disposal and Cleanup — Dispose of the excess solid zinc in a waste jar (not the sink). _______________________________________________________________________________________________

 

EMPIRICAL FORMULA OF ZINC IODIDE

_______________________________________________________________________________________________

 

  1. Mass of porcelain evaporating dish (P.E.D.) _______________ (g)

 

  1. Mass of P.E.D. + Zn granules at beginning _______________ (g)

 

  1. Mass of Zn at beginning _______________ (g)
  2. = (b) – (a)

 

—————————————————————————————

  1. Mass of screwcap vial + I2 _______________ (g)

 

(e) Mass of screwcap vial

 

_______________ (g)

(f) Mass of I2 at beginning

_______________ (g)

(f) = (d) – (e)

 

—————————————————————————————

(g) Mass of P.E.D. + unreacted Zn

 

_______________ (g)

(h) Mass of unreacted Zn

_______________ (g)

  1. = (g) – (a)

 

  1. Mass of reacted Zn _______________ (g)

(i) = (c) – (h)

 

(j) Moles of reacted I atoms (in product)

 

_______________

(k) Moles of reacted Zn atoms (in product)

 

_______________

(l) Mole ratio (I to Zn)

 

_______________

(m) Empirical formula

_______________

Show calculations for h, i, j, k and l on the corresponding pages of your data table.

DATA ANALYSIS SUPPLEMENT

Data: After you have recorded all the masses, enter your own data into the appropriate Web-based data entry form, using a computer in the lab. In your measurements and calculations, remember to pay attention to uncertainties and significant figures.

Results: Outside of class, retrieve your data in spreadsheet form (along with that of the entire class) from: http://chem.gmu.edu/results/. Then use Excel to calculate the moles of Iodine and Zinc in the product, the Mole ratio and the Empirical Formula.

Your lab report should include the following in addition to the pages from your lab notebook:

  • An Excel sheet “Raw Data” showing the raw data for the entire class (refer to guidelines stated on page 42). Your personal data should be in boldface.
  • A 2nd Excel sheet “Class Results” showing the calculated results for each student in the class – each in a separate row. Your personal data should be in boldface. The columns should include:
    • Mass of unreacted zinc o Mass of reacted zinc o Moles of reacted I atoms o Moles of reacted Zn atoms o Your Mole ratio (I to Zn)
    • Your Empirical formula
  • An Excel chart showing a histogram of the distribution of mole ratio results. Use equally spaced “bins” chosen to display the “spread” in the class results. Leave space between the vertical bars in the histogram.

Discussion: In your discussion section address the following questions:

  • What was your own percent error for the mole ratio?
  • What percentage of the students in your section obtained a mole ratio that was within (a) 10% and (b) 25% of the true value?
  • What percentage of the students in your section obtained a mole ratio that was higher than the true value?
  • What experimental error(s) might have caused the calculated mole ratio to come out too high?
  • What experimental error(s) might have caused the calculated mole ratio to come out too low?

 

 

Experiment 6 SYNTHESIS OF ASPIRIN

 

 

 

word image 889 word image 890 word image 891 word image 892

 

Aspirin is one of the best known and most widely used medicines in modern society. The pharmaceutical formulation, acetylsalicylic acid, was devised by the Bayer Chemical Company more than 100 years ago. Its medicinally active part, salicylic acid, is a naturally-occurring substance that can be extracted from willow tree bark and meadowsweet flowers. It had long been known and used as an herbal fever reducer, but salicylic acid is too irritating to the mucous membranes of the mouth, esophagus and stomach to be administered comfortably as such. Chemically modifying the salicylic acid by attaching an acetyl group minimizes the adverse side effects while retaining the desirable analgesic, antipyretic, and anti-inflammatory properties.

The same strategy of chemical modification was used later to modify morphine, a powerful but addictive painkiller. However, when morphine was converted into diacetylmorphine, the result was heroin. The modification intensified the addictive property instead of reducing it.

word image 893 As you inspect the molecular structures of salicylic acid and aspirin, below, notice the small change in structure as a hydrogen atom in salicylic acid (in boldface) is replaced by the acetyl group, CH3 C

O

 

Note: The vertices of the “hexagons” in salicylic acid and aspirin are occupied by carbon atoms.

H

H

H

H

C

O

O

 

H

H

O

C

H

3

C

O

O

C

O

C

H

3

H

H

H

H

C

O

O

 

C

H

O

C

H

3

O

C

H

3

C

O

O

H

Salicylic acid

C7H6O3

Acetic anhydride

C4H6O3

Acetyl salicylic acid

C9H8O4

Acetic acid

C2H4O2

A substance related to aspirin is oil of wintergreen, or methyl salicylate. Both of these compounds are synthesized from salicylic acid. Notice the similarity of names, which implies a similarity of molecular structure. Oil of wintergreen is largely used as a flavoring agent, but ingestion of relatively small amounts may cause severe poisoning and death. Oil of wintergreen is used in many liniments, however, as it is absorbed through the skin. Once absorbed, it may be cleaved back to salicylic acid; thus it is a source of pain relief, albeit localized pain relief.

H O H O

H

H

 

Salicylic acid

C7H6O3

Methanol Methylsalicylate

CH4O C8H8O3

 

THE EXPERIMENT

Water

H2O

 

C O CH3

C

O

H

O

H

H

H

O

H

H

C

H

H

H

O

H

H O H

H

In this experiment you will synthesize aspirin in much the same way as it is manufactured industrially. You will also calculate the percent yield of the aspirin you synthesized.

In the synthesis of aspirin from salicylic acid and acetic anhydride, the most likely impurity in the final aspirin product is salicylic acid itself, which is present either from incomplete reaction or from hydrolysis (the reverse reaction) of the aspirin product during the isolation steps. The salicylic acid impurity is removed during the various stages of the purification and in the final crystallization of the product. The presence of salicylic acid is detected by its colorful reaction with ferric chloride.

_________________________________________________________________________________________________________

Percent Yield

The theoretical yield for any reaction is the maximum number of grams (or moles) of product that may be obtained from the reaction based on the reaction stoichiometry. Before a yield can be calculated, it is necessary to first balance the chemical equation for the reaction.

Next, the moles of each reactant are calculated from the grams of each reactant (or volume if the reactant is liquid) which were used in the reaction. If the ratio of the moles of reactants in the balanced equation is different from the ratio of the moles of reactants actually used for the reaction, it is necessary to determine the limiting reagent in the actual synthesis.

The limiting reagent is the reactant which theoretically would be completely consumed in the reaction based on the molar ratio calculated from the actual quantities of reactants. Any other reactants are therefore present in excess, and will not be completely consumed. There can be no more moles of product formed than that amount available from the moles of the limiting reagent. The limiting reagent determines the maximum or theoretical yield of product(s). If all of the reactants in the synthesis are actually used in the synthesis in the same molar ratio as in the balanced equation, any of the reactants may be used to calculate a theoretical yield.

 

word image 894

The actual yield is the number of grams (or moles) of product actually obtained in the synthesis reaction. The percent yield describes the efficiency of the reaction:

word image 895

 

Sample Calculation

A sample calculation using an example of the aspirin preparation is shown below. The molecular weights of the compounds can be found in Table 1.

From the balanced chemical equation you can determine the molar ratios among the reactants and products. Here, one mole of salicylic acid reacts with one mole of acetic anhydride to give one mole of aspirin product and one mole of acetic acid. The molar ratio of the reactants is 1:1 as is the ratio of salicylic acid (reactant) to aspirin (product).

Next you must determine which of the two reactants (if either) is the limiting reagent in a laboratory synthesis. Suppose you started with 4.21 grams of solid salicylic acid and 15.0 mL of liquid acetic anhydride (density = 1.08 g/mL).

1 mol

Moles of salicylic acid used  4.21 g   0.0305 mol 138.12 g

1.08 g 1 mol

Moles acetic anhydride used  15.0 mL   0.159 mol 1 mL 102.09 g

Therefore, since salicylic acid and acetic anhydride react in a 1:1 ratio, but the ratio actually used is 0.0305/0.159 or 1 : 5.2, the acetic anhydride used here is clearly in excess and thus salicylic acid is the limiting reagent. The theoretical yield must be calculated using moles of salicylic acid. Theoretical yield (grams) =

1 mol aspirin 180.17 g aspirin

0.0305 mol salicylic acid    5.50 g aspirin

1 mol sal. acid 1 mol aspirin

Suppose the actual yield = 3.92 g aspirin, then:

Percent yield = (3.92 g ÷ 5.50 g)  100 = 71.3%

 

 

Vacuum Filtration

A common laboratory procedure is the separation of a solid from a liquid by filtration. Suction filtration requires a water vacuum aspirator and special equipment, but once set up, it proceeds rapidly.

How to filter with a Buchner Funnel:

  1. Assemble the equipment (ring stand, clamp, filter flask, rubber tubing, Buchner funnel, and filter paper (see Figure). Clamp the filter flask to the ring stand. Attach the flask side arm to the vacuum aspirator with the rubber tubing.
  2. Place the filter paper in the funnel, and moisten it with water or the solvent to be used. Do not allow it to dry before proceeding to the next step.
  3. Turn the vacuum aspirator on.
  4. Pour into the funnel, at a moderate rate, the solution to be filtered.
  5. To interrupt the vacuum, remove the rubber tubing from the flask. Do not turn off the water first.
  6. Rinse the solid in the funnel by pouring into the funnel a small amount of cold solvent. Stir the contents with a stirring rod, being careful not to tear the filter paper. Allow the solid to dry by drawing air through the funnel.
  7. Every time you do another filtration, use a fresh piece of filter paper.

 

Figure 1 ALWAYS SECURE SETUP USING A RING STAND AND UTILITY CLAMP.

word image 3212

 

PROCEDURE

Hazard: Acetic anhydride is a corrosive liquid and is quite flammable. Avoid breathing vapors and carry out all transfers in a HOOD. Concentrated sulfuric acid is also highly corrosive. Handle it with great care.

Before you begin the synthesis, half-fill a 400 mL beaker with tap water and heat it on an electric hot plate. Bring the temperature of the water to 45-50°C. This will serve as the hot bath. It is important that the temperature of the water bath is constant before heating the reactants.

  1. Precisely weigh, using a weigh boat, approximately 2 g of salicylic acid. Carefully transfer the salicylic acid into a clean and dry 125-mL Erlenmeyer flask.
  2. Take your flask containing the salicylic acid crystals to the HOOD and carefully add 5 mL of acetic anhydride. Measure the volume in a small graduated cylinder.
  3. In the hood, carefully add 4 drops of concentrated sulfuric acid (H2SO4) using a dropper. (Sulfuric acid acts as a catalyst in the reaction and is not the limiting reagent.) Gently swirl the mixture.
  4. Heat the flask in the water bath (45-50°C) for 10 minutes, stirring until all the salicylic acid dissolves. (Do not stir with a rubber spatula.)
  5. Remove the flask from the bath and allow it to cool to room temperature. Solid particles of acetylsalicylic acid (aspirin) should begin to precipitate out of the reaction mixture as it cools if it has not done so already. If precipitation does not occur, scratch the inside walls of the flask with a glass rod.
  6. Cool the mixture in an ice bath until no more aspirin precipitates from the liquid solution. While you are waiting, assemble the vacuum filtration apparatus.
  7. word image 3213 After precipitation is complete, add 20 mL of ice cold water to the reaction flask, stir to mix, and again cool the mixture in an ice bath.
  8. Collect the solid product in a Buchner funnel by vacuum filtration while the solution is still cold. The filtrate (the liquid solution that collects in the filter flask) can be used to rinse more product from the Erlenmeyer flask in which you performed the synthesis until all solid has been collected.
  9. Rinse the solid twice with small portions of cold water. Continue drawing air through the crystals in the Buchner funnel by suction until the crystals are relatively dry.

SAVE THE FILTRATE FOR PART B ON THE NEXT PAGE.

At this point, your instructor may choose to have you perform either Part A or Part B which follows. _____________________________________________________________________________

PART A – Recrystallization
  1. word image 3214 Transfer the crude aspirin product to a 150-mL beaker and add 8 mL of ethanol. Warm on a water bath if necessary to completely dissolve the crystals, then pour 17 mL of distilled water into the solution. Cover the beaker with a watch glass and set aside to cool to room temperature. Place the beaker in a container of ice to cool and further precipitate.
  2. Filter the solution by suction, using a Buchner funnel and fresh filter paper. Scrape the beaker to obtain as much of the solid product as possible. Do not wash the beaker or the crystals which are on the filter paper. Dry the crystals by spreading by pulling a vacuum on the crude wet product until it is light and fluffy. Gently separate the product periodically (using a glass stirring rod) to allow more air to pass through and dry the product.
PART B – Ferric chloride test

In order to determine if unreacted salicylic acid is present with your aspirin product, perform a ferric chloride color test.

  1. Into a small test tube containing about 5 mL of water, add a small amount of aspirin (about the size of a pen tip). Stir to dissolve.
  2. Into a second test tube put 5 mL of water and about the same amount of salicylic acid as the amount of aspirin you put in the first test tube. Stir to dissolve the solid.
  3. Into a third tube, put 5 mL of the filtrate you saved from the synthesis procedure.
  4. To each test tube, add 10 drops of 1% FeCl3 and stir.
  5. Compare the color in the three test tubes. Explain your results.

_________________________________________________________________________________________________________

Weigh the aspirin product and then calculate the percent yield. (If your aspirin is still wet, your percent yield may be more than 100%.)

Submit your product to your instructor in a vial, labeled with your name, the product name, the yield in grams, and percent yield.

 

TABLE 1

The molecular weights, formulas, and densities for the substances which have been mentioned in this experiment are given below. [CRC Handbook (92nd Ed.)]

Compound Formula M.W. (g/mole) Density (g/mL) salicylic acid C7H6O3 138.12 acetic anhydride C4H6O3 102.09 1.0820

acetylsalicylic acid C9H8O4 180.17

acetic acid C2H4O2 60.05 1.0492 methanol CH4O 32.04 0.7914

methyl salicylate C8H8O3 152.16

Problems (show work)

  1. What is the purpose of (a) recrystallization; (b) the sulfuric acid in aspirin synthesis?
  2. If you use 12.22 g of salicylic acid and an excess of acetic anhydride in the synthesis of aspirin, what would be the theoretical yield of acetylsalicylic acid in moles? In grams? Determine the percent yield if 10.50 g of aspirin is isolated?

Answers to Selected Problems

2) 8.847 x 10-2 mol; 15.94 g; 65.87% _____________________________________________________________________________________________________

SYNTHESIS OF ASPIRIN

_____________________________________________________________________________________________________

Chemical Equation for synthesis of Aspirin:

Show all calculations. Pay attention to significant figures.

Reagent

volume

mass

mole

 

 

 

 

salicylic acid

———–

____________

___________

acetic anhydride

___________

___________

___________

limiting reagent ______________________________________

Theoretical Yield of Aspirin ____________________________

Mass of aspirin isolated _______________

Percent yield of aspirin _____________

Part B Observations and Conclusion: (These go directly in your lab notebook.)

 

DATA ANALYSIS SUPPLEMENT

Data: After you have recorded all the masses and volumes, enter your own data into the appropriate Web-based data entry form, using a computer in the lab. In your measurements and calculations, remember to pay attention to uncertainties and significant figures.

Results: Outside of class, retrieve your data in spreadsheet form (along with that of the entire class) from: http://chem.gmu.edu/results/. Then use Excel to calculate the moles of Salicylic acid and Acetic anhydride, and the Theoretical yield and Percent yield of aspirin.

Your lab report should include printouts of data and result tables from your spreadsheet in addition to the pages from your lab notebook. Each table should be formatted as discussed in earlier experiments.

  • An Excel sheet “Raw Data” showing the raw data for the entire class (properly formatted column headings, units, sigfigs, etc.). Row and column labels should always be printed for each Excel sheet.
  • A 2nd Excel sheet “Class Results” showing the calculated results for each student in the class – each in a separate row. Each table should be formatted as in previous classes with lines separating the table headings from the contents. A line below the contents should also be drawn. Make sure to show an example calculation for each column. This can be done by copying the equation to a blank cell below the elements in the table. An apostrophe in front of it makes it text instead of an equation. The columns should include:

o Moles of salicylic acid o Moles of acetic anhydride o Theoretical yield of aspirin o Mass of aspirin isolated o Percent yield of aspirin

  • An Excel chart showing a histogram of the distribution of percent yield results. Use equally spaced “bins” chosen to display the “spread” in the class results.

Discussion: In your discussion section address the following questions:

  • Which reactant, Acetic anhydride or Salicylic acid is the limiting reagent? Why is it called so?
  • What is the most likely impurity in the final aspirin product? How is its presence detected?
  • What are the possible sources of error in this experiment? Refer in your answer to the class histogram as well as to your own results.

Printing Tables:

  • Large tables should be printed in the landscape mode and sometimes it might be necessary to go to “page setup” and select “Fit to 1 page” to make the page fit on one page.
  • Include column and row headings on the printout so that the instructor can more easily understand equations used. This can be done by going to File → Page Setup → Sheet tab → Select “Row and Column Headings”.

 

 

 

Experiment 7 VOLUMETRIC ANALYSIS: TITRATION OF VINEGAR

 

 

 

 

word image 896 word image 897 word image 898 word image 899 word image 900

Volumetric analysis is the quantitative determination of substances by titration. Titration is a laboratory procedure in which one solution is added, by means of a buret, to a second solution until all of the reactant in the second solution has been consumed in a chemical reaction. The solution added from the buret is the titrant. A standard solution is a solution whose concentration is known accurately. The standard solution may be used to titrate a solution of unknown concentration, and by means of appropriate calculations, the concentration of the second solution can be determined.

Acid-Base Titrations

When a solution of an acid is mixed with a solution of a base the reaction that occurs is called a neutralization reaction. When a strong acid (symbolized HA) and a strong hydroxide base are mixed the reaction can be represented by the partial ionic equation:

word image 901 H+A(aq) + OH(aq) A(aq) + H2O (l) (1)

When a diprotic acid (H2A) is neutralized by a strong hydroxide base, the acidic protons are removed in a stepwise fashion :

H2A(aq) + OH(aq) HA(aq) + H2O (l) (2a)

HA (aq) + OH(aq) A2−(aq) + H2O (l) (2b)

H2A(aq) + 2OH(aq) 2H2O (l) + A2−(aq)

In aqueous media, the overall reaction for the diprotic reaction can be represented in the following manner:

2H+ (aq) + A2− (aq) + 2OH(aq) 2H2O (l) + A2−(aq) (2c)

In both cases, the reactions are rapid and essentially quantitative. Such acid-base neutralization reactions are commonly used in volumetric analysis to determine the concentrations of acidic or basic solutions. The equivalence point (or neutralization point) of the titration is often recognized by the color change of an indicator which is added to the titrated solution.

A common laboratory problem is to determine the concentration of acid in a solution. To do this, the acidic solution is titrated with a standard base, usually NaOH of known concentration. However, when solutions of NaOH are prepared, the concentration of the base solution is not accurately known because of the difficulty of handling solid NaOH. In order to establish the concentration of the basic NaOH which is to be used in a titration, the NaOH solution concentration must first be determined by titration with an acid. The acid which is used to titrate the NaOH solution is called a primary standard acid solution. The primary standard is easily handled and solutions of very accurately known concentration can be prepared. The overall procedure can by represented by:

primary standard

 

secondary standard

 

 

acid

titrate

>

base

titrate

>

acid

[known conc.]

 

[calculate conc.]

 

[calculate conc.]

 

 

word image 3215

THE EXPERIMENT

Vinegar is an aqueous solution of acetic acid, a simple organic acid. The expanded structural formula for acetic acid is

H O

H

C

C

O

H

H

 

Only the hydrogen bonded to the oxygen is acidic. The other three hydrogens bonded to carbon do not react with a base such as NaOH. Acetic acid is a weak acid, which reacts with NaOH according to the equation:

CH3CO2H + Na+(OH) CH3CO2 Na+ + HOH (3)

In this experiment, the ultimate goal is to determine the concentration of acetic acid in vinegar purchased from the grocery store by titrating it with a standard NaOH solution (a secondary standard). For the standardization of the NaOH solution, oxalic acid (in the form of its hydrate H2C2O4•2H2O), will be used as a primary standard. Both hydrogens, which are bonded to oxygens in the oxalic acid molecule, are acidic and thus 2 moles of NaOH will be required to neutralize every 1 mole of H2C2O4 during the titration:

H2C2O4 + 2 Na+(OH) 2 Na+ (C2O4)2− + 2 H2O (4)

Oxalic acid is also a simple organic acid and is used frequently as a bleaching agent, especially in marine settings. It is very corrosive, so be careful when you handle it.

Phenophthalein, which is a suitable indicator for the titration of a weak acid with a strong base, is used in the titration. Phenolphthalein is another weak organic acid. Its color change is from colorless in acidic solution to pink in basic solution upon removal of the acidic proton present in the molecule. A more exhaustive treatment of indicators and their roles during titration will be addressed during the second semester laboratory course (LeChatelier’s Principle and Equilibrium).

The molarity of the acetic acid solution can be determined from the volume and molarity of the NaOH used in the vinegar titration. Finally, the weight% of acetic acid in the vinegar solution will be calculated and compared with the 5wt% allowed by Federal Law.

Calculations

The fundamental equation used in titration calculations is the same as the one previously used for dilutions:

VA MA = VB MB

The volumes and concentrations pertain to an acid-base neutralization reaction rather than dilution of a concentrated solution. The only adjustment needed for the acid-base titration calculation is for those cases where the reaction stoichiometry is not 1:1. In a 1:1 neutralization such as in equations (1) and (3), the number of moles of acid which transferred H(+) and the number of moles of base (OH) which accepted them must be the same at the neutralization point:

VAMA = VBMB (5)

 

VAMA = molesA = nA ; VBMB = molesB = nB

 

nA = nB

At the neutralization point in reactions, such as those in equation (2) and (4), the two moles of

H(+) from one mole of the oxalic acid have combined with 2 moles of OH(-) from 2 moles of the base.

2(VH2AMH2A) = VBMB (6)

2 • molesA = molesB

 

2nA = nB

 

 

Example 1: Calculate the molarity of KOH if 20.10 mL was used to neutralize 13.62 mL of 0.75-M HCl.

  1. KOH (aq) + 1 HCl (aq) 1 KCl (aq) + 1 H2O (l)

VHClMHCl = VKOHMKOH

MKOH VHClMHCl (13.62)(0.75) 0.51-M VKOH (20.10

Note: Either mL or L may be used for the volumes (V).

 

Example 2: How many milliliters of 2.1-M NaOH are required to neutralize 42.85 mL of 3.6-M H2SO4?

  1. NaOH (aq) + H2SO4 (aq) Na2SO4 (aq) + 2 H2O (l)

2 (VH2SO4MH2SO4) = VNaOHMNaOH

VNaOH  2(VH2SO4MH2SO4) 2(0.04285)(3.6)  0.1469 L 146.9 mL

2.1

M

NaOH

 

PROCEDURE

CAUTION: Oxalic Acid is corrosive.

1. Preparation of Oxalic Acid Standard Solution

  1. Weigh ~3 g of pure oxalic acid dihydrate crystals, H2C2O4•2H2O, to ± 0.001 gram into a weigh boat. (Mol. Wt. oxalic acid dihydrate = 126.1 g/mol)
  2. Carefully transfer the weighed crystals, without loss, to a clean 100-mL volumetric flask, and then add 50-75 mL of distilled water (avoid filling the flask into the neck), and gently mix the solution by swirling it in the flask until solution is complete. (Save time here by preparing your NaOH solution for Part 2 while these crystals are dissolving.)
  3. Allow time for all of the crystals to dissolve, then fill the 100-mL volumetric flask carefully to the mark with additional distilled water, adding the last portion by drops. Stopper the flask and mix the solution well by repeated inversion and swirling.

Calculate the number of moles of oxalic acid weighed into the weigh boat (step a above). Calculate the molarity of the primary oxalic acid standard acid solution you prepared in the volumetric flask.

2. Standardization of the NaOH Solution

  1. Rinse two burets thoroughly with tap water, then once with distilled water, and allow to drain.
  2. Rinse one buret with a 5 mL portion of approximately 0.5-M NaOH solution and discard the rinsing solution. Repeat this rinsing procedure. Run at least a portion of each rinse through the stopcock and the tip so that all parts inside the buret will have come in contact with the NaOH solution.
  3. Fill the buret with the NaOH solution above the 0.00-mL mark. Allow the solution to drain out through the stopcock and tip to ensure that they are full. Make sure there are no air bubbles in the tip. The liquid level may be stopped at the 0.00-mL mark or at any point somewhat below this.
  4. Rinse the second buret with oxalic acid, then fill it with the standard oxalic acid solution. Label the burets.

Touch off the remaining drop at the tip of each buret on the side of a beaker. Always keep a waste container underneath burets whenever experimental solutions are in them.

 

Now you are ready to conduct your first titration. The first titration is considered as your “scout titration”.

  1. Deliver approximately 10 mL of the oxalic acid solution from the burtet into a clean, rinsed 125-mL Erlenmeyer flask. (The flask need not be dry since the water will not affect the amount of oxalic acid delivered from the buret.) Touch the last drop on the buret tip to the inside of the flask, and rinse it into the flask with distilled water from a squeeze bottle.
  2. Add 2 drops of the phenolphthalein indicator solution.
  3. Begin titrating the oxalic acid with the NaOH. Add the base solution rapidly at first and mix the solution by swirling the flask. As the endpoint is approached, the faint pink color existing in parts of unmixed NaOH solution will persist for several seconds. At this point, decrease the rate of NaOH addition until you are adding the drops one at a time and thoroughly mixing the solution before the addition of the next drop.

The endpoint is reached when a very faint pink color persists throughout the solution after it has been swirled, thereby indicating that the solution is basic. If the pink color fades after the solution has been standing, it will be because carbon dioxide from the air has been absorbed, forming carbonic acid.

If you should add too much base and go past the end point by more than 1 drop, add enough additional oxalic acid solution from its buret to remove the pink color (remember to record the volume added). Continue the titration with the NaOH, adding the base more slowly this time.

 

  1. Read and record the final buret reading for each solution.
  2. Repeat the titration.
  3. Calculate the molarity of the NaOH solution from the two titrations. If the molarities agree within 1%, proceed to the next part. If they do not, perform a third titration.

In calculating the actual molarity of the sodium hydroxide solution, you must remember that there are two titratable hydrogens per oxalic acid molecule (equation 6), therefore, 2 moles of NaOH are required to titrate one mole of H2C2O4.

 

3. Determination of Acetic Acid in Vinegar

  1. Empty the buret containing the oxalic acid.
  2. Rinse the buret and fill it with vinegar, using proper rinsing procedures.
  3. Carry out at least two titrations of acetic acid using NaOH as the titrant, as you did in Part 2 (using vinegar instead of oxalic acid), and phenolphthalein as indicator.

Calculate the molarity of acetic acid. The molecular weight of acetic acid is 60.05 g/mol.

Calculate the %(w/w) of acetic acid in the vinegar solution.

The % (w/w) of a solution is the ratio of grams of solute to grams of solution, expressed as a percent. In this experiment, acetic acid is the solute and vinegar is the aqueous solution.

acetic acid (g)

%(w/w) 

vinegar (g)

Up to this point, you have determined the molarity of acetic acid in vinegar solution (moles acetic acid/L vinegar). In order to calculate % (w/w) you need the following conversion factors:

grams acetic acid

  • the molecular weight of acetic acid 

moles acetic acid

1.005 g vinegar

  • the density of vinegar solution  mL

Show ALL calculations in the appropriate space near your data. Pay attention to SIGNIFICANT FIGURES. _________________________________________________________________________________________________________

 

TITRATION OF VINEGAR

_________________________________________________________________________________________________________

 

1. Preparation of Standard Oxalic Acid Solution

 

Weight of oxalic acid dihydrate

 

 

g

Molecular weight of oxalic acid dihydrate

 

 

g/mol

Moles of oxalic acid dihydrate

 

 

moles

Molarity of oxalic acid solution H2C2O4

 

2. Standardization of the NaOH Solution with Oxalic Acid

 

M

 

 

H2C2O4 buret:

Trial 1

Trial 2

Trial 3

 

 

 

(if necessary)

Initial Reading

 

________(mL)

________(mL)

________(mL)

Final Reading

 

________(mL)

________(mL)

________(mL)

Volume of H2C2O4 delivered

 

________(mL)

________(mL)

________(mL)

Moles of H2C2O4 delivered

________ moles

________ moles

________ moles

 

NaOH buret:

Initial Reading

 

________(mL)

________(mL)

________(mL)

Final Reading

 

________(mL)

________(mL)

________(mL)

Volume of NaOH used

 

________(mL)

________(mL)

________(mL)

Moles of NaOH used

(based on moles of H2C2O4 titrated)

 

________ moles

________ moles

________ moles

Molarity of NaOH

 

________ M

________ M

________ M

Average Molarity of NaOH

 

________ M

 

3. Determination of Acetic Acid in Vinegar

 

 

Trial 1

Trial 2

Trial 3

 

Vinegar buret:

 

 

 

(if necessary)

Initial Reading

 

________(mL)

________(mL)

________(mL)

Final Reading

 

Volume of vinegar

________(mL)

________(mL)

________(mL)

(acetic acid solution)

 

NaOH buret:

 

________(mL)

________(mL)

________(mL)

Initial Reading

 

________(mL)

________(mL)

________(mL)

Final Reading

 

________(mL)

________(mL)

________(mL)

Volume of NaOH

 

________(mL)

________(mL)

________(mL)

Average Molarity of NaOH calculated from Part 2

 

 

______ M

 

Molarity of acetic acid ________ M

________ M

________ M

 

Average Molarity of acetic acid ________ M

 

Calculate the grams of CH3COOH per liter of vinegar solution, using the average value of the molarity of CH3COOH.

g CH3COOH/L vinegar solution

Calculate the grams of acetic acid per gram of vinegar solution, using the density of vinegar solution = 1.005 g/mL.

g CH3COOH/g vinegar solution

Calculate the weight percent (%wt/wt) CH3COOH in your vinegar sample.

% % Error

 

DATA ANALYSIS SUPPLEMENT

Data: After you have recorded all of the masses and volumes, enter your own data into the appropriate Web-based data entry form, using a computer in the lab. In your measurements and calculations, remember to pay attention to uncertainties and significant figures.

Results: Outside of class, retrieve your data in spreadsheet form (along with that of the entire class) from: http://chem.gmu.edu/results/. Then use Excel to calculate the moles of oxalic acid dihydrate, molarity of oxalic acid solution, moles of H2C2O4 delivered, molarity of NaOH, average molarity of NaOH, molarity of acetic acid, average molarity of acetic acid, grams of CH3COOH per liter of vinegar solution, grams of acetic acid per gram of vinegar solution, weight percent CH3COOH in the vinegar sample and %error.

Your lab report should include the following in addition to the pages from your lab notebook:

  • An Excel sheet “Raw Data” showing the raw data for the entire class (properly formatted). An unformatted example of the raw data sheet is shown below:

Raw Data

 

 

A

B

C

D

E

F

G

H

I

J

K

L

M

N

O

P

1

ID

 

number

_ox

 

number

_acet

 

wt_oxalic

g

V1_A mL

V1_B mL

V2_A mL

V2_B mL

V3_A mL

V3_B mL

V1_Acet mL

V1_B2 mL

V2_Acet mL

V2_B2 mL

V3_Acet mL

V3_B2 mL

 

2

1

3

3

3.000

20.00

35.00

20.50

36.10

20.20

35.50

25.00

25.00

25.00

25.00

25.00

25.00

3

2

2

2

3.021

20.00

19.20

20.50

19.50

 

 

20.00

32.30

20.10

32.60

 

 

4

3

2

2

3.033

20.18

18.99

20.11

18.98

 

 

19.99

32.79

19.89

32.56

 

 

 

Explanation:

      • number_ox: number of titrations to standardize NaOH o number_acet: number of titrations to determine the vinegar concentration.
      • wt_oxalic: mass of oxalic acid used to make up the standard oxalic acid solution (all in g) o V1_A, V2_A, V3_A: volume of oxalic acid used for titration 1,2,3 (all in mL)
      • V1_B, V2_B, V3_B: volume of base (NaOH) used for titration 1,2,3 in the standardization of NaOH (all in mL)
      • V1_Acet, V2_Acet, V3_Acet: volume of acetic acid used for titration 1,2,3 of the determination of acetic acid in vinegar for part 2 (all in mL)
      • V1_B2, V2_B2, V3_B2: volume of base (NaOH) for titration 1,2,3 of part 2 – the determination of acetic acid in vinegar for part 2 (all in mL)
  • A 2nd Excel sheet “Results: Oxalic acid” showing the calculated results for each student in the class – each in a separate row. Make sure to show an example calculation for each column. This can be done by copying the equation from the formula bar to a blank cell below the elements in the table. An apostrophe in front of it makes it text instead of an equation. Each table should be formatted as in previous classes with lines separating the table headings from the contents. A line below the contents should also be drawn. A typical example of the results sheet is shown below:

Results: Oxalic acid

 

 

A

B

C

D

E

F

G

H

I

J

K

L

1

mol

Oxalic moles

[Ox]

M

n1_Ox used moles

n2_Ox used moles

n3_Ox used moles

n1_NaOH used moles

n2_NaOH used moles

n3_NaOH used moles

[NaOH]1

M

[NaOH]2

M

[NaOH]3

M

[NaOH] ave

M

2

0.0238

0.238

0.0048

0.0049

0.0048

0.0095

0.0098

0.0096

0.272

0.270

0.271

0.271

3

0.0240

0.240

0.0048

0.0049

 

0.0096

0.0098

 

0.499

0.504

 

0.502

4

0.0241

0.241

0.0048

0.0048

 

0.0097

0.0096

 

0.511

0.510

 

0.151

 

Explanation:

    • mol Oxalic: moles of oxalic acid dihydrate used to prepare standard solution (moles) o [Ox]: molarity of oxalic acid solution (M)
    • n1_Ox used, n2_Ox used, n3_Ox used : moles of oxalic acid dihydrate used for standardizing sodium hydroxide (all in moles)
    • n1_NaOH used, n2_NaOH used, n3_NaOH used: moles of sodium hydroxide used to neutralize oxalic acid (all in moles) o [NaOH] 1, [NaOH] 2, [NaOH] 3: molarity of sodium hydroxide (all in M) o [NaOH] ave: average molarity of sodium hydroxide (M)

 

  • A 3rd Excel sheet “Results: Acetic acid” showing the rest of the calculated results. A typical example of the results sheet is shown below:

 

 

Results: Acetic acid

 

 

A

B

C

D

E

F

G

H

I

J

1

n1_NaOH used

Moles

n2_NaOH used moles

n3_NaOH used moles

[Acetic]1

M

[Acetic]2

M

[Acetic]3

M

[Acetic]ave

M

g Acet/L soln g/L

wt%

 

%Error

 

2

0.0068

0.0068

0.0068

0.271

0.271

0.271

0.271

16.278

1.62

61.90

3

0.0162

0.0164

 

0.810

0.814

 

0.812

48.758

4.85

2.97

4

0.0167

0.0166

 

0.837

0.836

 

0.836

50.256

5.00

0.00

5

 

 

 

 

 

 

 

Average =

3.82

 

6

 

 

 

 

 

 

 

StdDev =

1.91

 

 

Explanation:

      • n1_NaOH used, n2_NaOH used, n3_NaOH used: moles of sodium hydroxide used to neutralize vinegar solution (all in moles) o [Acetic] 1, [Acetic] 2, [Acetic] 3: molarity of acetic acid (all in M) o [Acetic] ave: average molarity of acetic acid (M)
      • g Acet/L soln: grams of CH3COOH per liter of vinegar solution using average molarity of acetic acid (g/L) o wt%: weight percent of CH3COOH in the vinegar sample o %Error: difference between legal value and experimental value o Average: average weight percent of acetic acid in vinegar solution o StdDev: standard deviation observed in weight percent of acetic acid in vinegar solution
  • A 4th Excel sheet “Histogram” showing a histogram of the distribution of the weight percent results. The histogram (inclusive of borders) should cover at least 80% of the printed page. DO NOT INCLUDE BINS ON THE HISTOGRAM PAGE. Use equally spaced “bins” chosen to display the “spread” in the class results.

Printing Tables:

  • Large tables should be printed in the landscape mode and sometimes it might be necessary to go to “page setup” and select “Fit to 1 page” to make the page fit on one page.
  • Include column and row headings on the printout so that the instructor can more easily understand equations used. This can be done by going to File → Page Setup → Sheet tab → Select “Row and Column Headings”.

Discussion: In your discussion section address the following questions:

  • What is equivalence point and how is it recognized in this experiment?
  • How well does your weight percent acetic acid in vinegar solution agree with the legal standard?
  • How well does the class average agree with the legal value?
  • Are the deviations systematic or random?
  • What are the possible sources of error in this experiment? Refer in your answer to the class histogram as well as to your own results.

 

 

Experiment 8 THE IDEAL GAS LAW

 

 

 

word image 902 word image 903 word image 904 word image 905 word image 906

IDENTIFICATION OF AN UNKNOWN PURE METAL SAMPLE

Purpose

The purpose of this experiment is to use the modern Ideal Gas Law to quantify and determine the identity of a sample metal from the volume of gas produced over water in a reaction of the metal with an aqueous strong acid.

Introduction

Gases are everywhere – atmosphere, environmental processes, industrial processes, bodily functions.

  • Gases have unique properties from liquids and solids.
  • word image 907 Gases are compressible (very important for storage).
  • Gas particles are widely separated (dispersed) and move at very fast speeds.
  • Most gases have relatively low densities.
  • Gases have relatively low viscosity (resistance to movement) allowing them move freely through pipes and even small orifices.
  • Gases are miscible, mixing together in any proportion.
  • The physical behavior of all gases is similar.
  • Chemical behavior of gases depends on composition.

The primary metric for the measurement of gases is pressure, which is defined as Force per Unit Area. Force, in turn, is a function of a mass under acceleration of gravity.

Force (F) = mass (m) x acceleration of gravity (a) = m x a

Pressure (P) = Force (F)/unit area (A) = F / A

where pressure has common units of:

1 Pascal (Pa) = 1 kg/m.s2

1 atmosphere (atm) = 101,325 Pa = 101.325 kPa

1 Bar = 100 kPa

1 atmosphere (atm) = 1.013 bar

1 atmosphere (atm) = 14.7 psi

1 atmosphere (atm) = 760 mm Hg = 760 torr

1 torr = 1 mm Hg = 133.3 Pa

Gas behavior can be described by pressure (P), temperature (T), volume (V), and molar amount (n). Holding any two of the variables constant allows relations between the other two. There are three primary laws that govern the relationship between these variables.

1) Boyle’s Law: The volume of a sample of gas at a given temperature varies inversely with the applied pressure.

V  1/P PV = constant P1V1 = P2V2

2) Charles Law: The volume occupied by any sample of gas at constant pressure is directly proportional to its absolute temperature

V  Tabs (V/T) = constant (V1/T1) = (V2/T2)

3) Avogadro’s Law: The volume of a sample of gas is directly proportional to the number of moles of gas, n

V  n (V/n) = constant (V1/n1) = (V2/n2)

The above three laws can be “combined” into the modern Ideal Gas Law:

PV = (constant)(nT) = nRT n = (PV/RT)

where the “constant” term is referred to as the Ideal Gas Law Constant, R, with units of:

8.314 kg •m2

R = = 8.314 Joules / mol•K mol•K

2

s

L•atm

R = 0.0821

mol•K

The Ideal Gas Law takes advantage of a special set of physical conditions called collectively standard temperature and pressure, or “STP” for short. Standard temperature is 273.15 Kelvin (0oC, 32oF) and standard pressure is 101.325 Pa. Using SI (metric) units, one mole of any gas sample (regardless of formula or mass) collected at STP occupies an equally standard volume in liters, 22.4 L.

 

Design of the Experiment and Background Chemistry

A sample of an unknown metal will be allowed to react with aqueous hydrochloric acid (HCl). Hydrogen gas will be produced as one of the reaction products, and a water-soluble metal chloride salt will be the other product. The unknown metal assigned by your lab instructor could be either a divalent (M2+) metal or a trivalent (M3+) metal.

The unknown metal will dissolve as the reaction proceeds, and the gas produced will be collected over water and quantified by volume (mL). Using the Ideal Gas Law, the value n, number of moles of gas produced, can be determined. From the known mass and number of moles, the molecular weight of the metal can be determined and associated with the applicable element from the Periodic Table.

The atmospheric pressure to be used in the experiment will be obtained from the laboratory barometer or supplied to the class by the lab instructor or prep room staff. The value obtained will probably be in units of millimeters (mm) of Mercury and will have to be converted to atmospheres using an appropriate conversion factor (see table above).

Volume measurements will be in milliliters (mL) and will have to be converted to a volume unit compatible with the value of R used in the calculations, generally, liter (L). Likewise, temperature values will be measured in degrees Celsius and will have to be converted to degrees Kelvin.

The gas collected during the experiment is not just hydrogen. Since it is being collected over water, the pressure exerted is a combination of the hydrogen gas produced from the reaction and the pressure of the water vapor. Thus, the total pressure of the gas collected is dictated by Dalton’s Law of Partial Pressures:

Ptotal = P(H2) + P(H O2 )

The vapor pressure of water needs to be subtracted from the total pressure. It is the pressure due to just the hydrogen gas that must be used to determine the number of moles, n, using the Ideal Gas Law.

P(H2) = P(total) – P(H O2 )

The vapor pressure of water is temperature dependent and can be obtained from the lecture text or the table below. Values not precisely stated (e.g., the 23.1oC value used in the example below) can be determined by extrapolating between two appropriate values in the table.

Vapor Pressure (VP) interpolation example

Room Temp: 23.1oC

VP of water at 22oC: 19.8 torr

VP of water at 24oC: 22.4 torr

VP difference: (22.4 – 19.8) torr = 2.6 torr / 2 oC = 1.3 torr/oC

Difference * 0.1: 1.3 * 0.1 = 0.13

VP of water at 23.1oC: (19.8 + 1.3 + 0.13) torr = 21.2 torr = 21.2 mmHg

Table 8.1: Vapor Pressure of Water vs. Temperature

Vapor Pressure of Water

Temp

            

(oC)

18

19

20

21

22

23

24

25

26

27

28

29

Press (Torr)

15.5

16.5

17.5

18.7

19.8

22.1

22.4

23.8

25.2

26.7

28.3

30.0

Reaction of Divalent Metals with Acid

M(s) + 2 HCl(aq)  MCl2(aq) + H2 (g)

Molar Equivalence: 1 mole metal (M(s) produces 1 mole gas (H2 (g))

Mass M s 

Moles M s = n = 1 =  

Mol Wgt M s 

Moles H = n2 gas = 1 = PV

RT

 

Mass M s 

 

Mol Wgt 1

=

PV 1

RT

 

Mass M s  1 PV 1

= × = ngas Mol Wgt M s  1 RT 1

 

RT Mass M s 

Mol Wgt M s = Mass M s ×    PV = ngas

Reaction of Trivalent Metals with Acid

2M(s) + 6 HCl(aq)  2MCl3(aq) + 3H2 (g)

Molar Equivalence: 2 moles metal (M(s) produces 3 moles gas (H2 (g))

Mass M s 

Moles M s = n = 2 =  

PVMol Wgt M s 

Moles H = n2 gas = 3 =

RT

 

Mass M s 

Mol Wgt 2

=

PV 3

RT

 

Mass M s  2 PV 2

= × = × ngas

Mol Wgt M s  3 RT 3

 

3 RT 3 Mass M s 

Mol Wgt M s = × Mass M s ×  2 PV = ×2 ngas

Sample Data and Calculations

Mass of unknown #1 (divalent) sample: 0.111 g

Volume of gas produced = 43.6 mL (0.0436 L)

Temperature = 23.1 oC (296.3 K)

Atmospheric Pressure = 761.4 torr ( or mmHg)

Interpolated Vapor Pressure of water @ 23.1 oC = 21.2 torr

Corrected VP = (761.4 – 21.2) torr = 740.2 torr

Conversion of torr to atm = 740.2 torr x (1 atm/760 torr) = 0.9755 atm

Solving for moles (n)

PV = nRT

PV 0.9755 atm0.0436 L

n = = = 0.00175 mol H2

RT 0.08206 L • atm 296.3 K

 mol • K 

From Molar Ratio of metal to gas (1:1):

0.00175 mol unknown metal (M(s)) produces 0.00175 mol H2

Mass of Sample = 0.111 g

Mass 0.111 g

Mol Wgt = = = 63.4 g / mol moles 0.00175 mol

From Periodic Chart: Unknown metal = Copper (Cu) Mol Wgt = 63.55 g / mol

 

 

Equipment and Chemical Reagents:

Materials

Equipment

Unknown metal, pre-cut,

400-mL Beaker

Unknown Number

large test tube (18-20 mm diameter)

10 mL of 6M HCl

50-mL graduated cylinder

(enough for all 3 tests)

ring stand and utility clamp

 

assembled stopper/gas delivery system

(one-hole rubber stopper w/glass tubing outlet)

 

5-mL-sample vial

 

rubber tubing

 

250-mL beaker (to hold reaction tube)

You will perform 3 trials of the full experimental procedure for your assigned unknown metal. For each trial, you will determine n, relate this value for n to moles of unknown metal using the given stoichiometric relationships, and determine the mass of unknown metal-to-moles of unknown metal ratio. You will then average the three mass/moles ratios before attempting to identify your unknown metal sample.

______________________________________________________________________________ Experimental Procedure

  1. Determine the atmospheric pressure in the lab from the barometer or obtain the value from the instructor.
  2. Convert the pressure in millimeters (mm) to atmospheres (atm).
  3. Obtain the temperature in the lab.
  4. Convert the temperature in degrees Celsius (oC) to Kelvin (K).
  5. word image 3216 Fill the 400-mL beaker with about 300 mL of tap water.
  6. Fill the 50-mL graduated cylinder as completely as possible with tap water.
  7. Cover the open end of the cylinder with the latex glove.
  8. Quickly pick up the cylinder, invert it, and insert it into the 400-mL beaker.
  9. Remove the glove, taking care to lose as

little water as possible. There will be a small air space at the top of the inverted cylinder, this is normal.

  1. Record the volume of air in the cylinder by reading the meniscus point on the cylinder scale. This value is the Volume (initial) in your calculations .
  2. Clamp the cylinder in place using the clamp. A complete assembly will be available on your lab instructor’s desk.
  3. Thread the rubbing tubing of the pre-assembled stopper/gas delivery system under the water line and into the mouth of the inverted graduated cylinder; it should be far enough inside the cylinder to ensure that it will not come out while gas is being collected.
  4. word image 3217 Place the medium test tube into the 250 mL beaker, and test fit the stopper end of the gas/delivery system into the top of the test tube.
  5. Carefully insert the end of the rubber tubing far enough into the open end of the graduated cylinder to ensure it will stay in place during the reaction.
  6. Obtain a sample of unknown metal, noting the unknown number and determine its mass to the nearest 0.001 g on the electronic balance.
  7. Prepare the sample for reaction by carefully rolling or folding it, so that it will drop into the test tube.
  8. Remove the rubber stopper from the top of the test tube and drop the sample into the bottom of the test tube.
  9. Fill the 5-mL sample vial half full with 6.00 M hydrochloric acid (HCl) solution
  10. Carefully slide the vial containing the acid into the test tube making sure no acid spills out and the vial comes to rest at the bottom of the tube.
  11. Replace the rubber stopper into the top of the test tube.
  12. Being careful to keep the rubber tubing inside the cylinder, tilt the test tube containing the vial of acid and metal sample until the acid runs out of the vial. Tilt back and forth until all of the acid is removed from the vial.
  13. Return the test tube to the 250-mL beaker and allow the reaction to proceed to completion.
  14. When the reaction is over (bubbling stops), read the volume reading on the graduated cylinder and record the value as Volume (final).

24.

Rinse the test tube and vial;

 

waste is always required..

dispose of waste as instructed—proper disposal of

 

  1. Repeat the process for two (2) more trials.
  2. At the end of the experiment, return the reaction vessel/gas collection tube and plumbing, clean, and shaken dry, to the instructor’s desk.

Test fit the stopper onto the top of the test tube, being careful not to allow the open end of the rubber gas delivery tube to slip out of the graduated cylinder–the fit should be snug to prevent gas escaping from around the stopper/gas delivery assembly during the reaction in a manner similar to Figure 1.

Waste Disposal Information

There should be no unreacted pieces of unknown metal remaining in the test tube; but if any such pieces do remain, dissolve them in a small amount the hydrochloric acid solution, dilute, and flush down the drain.

Data Tables and Calculated Quantities

(Note: values with an asterisk (*) are values for which you should show at least one completely worked example in your lab report).

_____________________________________________________________________________ Number of the unknown metal: _____

 

Trial #1 Trial #2 Trial #3

Mass of unknown metal sample: __________ __________ __________

Initial volume of trapped air (if any): __________ __________ __________

Final volume (evolved H2, trapped air):__________ __________ __________

Total volume of H2 produced: __________ __________ __________

Room temperature (°C): __________ *K: __________

Room atmospheric pressure (given in lab): __________

unit: __________

 

*Room atmospheric pressure (atm)

__________

 

*Vapor pressure of water, interpolated for room temperature:

__________

 

*Pressure of H2 gas:

__________

 

*Moles of H2 gas evolved __________

__________

__________

 

*Molar equivalent of unknown metal: __________

__________

__________

*Mass metal-to-moles metal ratio: __________

__________

__________

 

*Average mass-to-moles ratio:

__________

 

Based upon the value for average mass-to-moles ratio, the unknown metal is:

______________________

Further Questions for Thought and Review

(include detailed answers, with calculations when appropriate, in your lab report)

  1. Explain why it is necessary to bring the water level inside the graduated cylinder equal to the water level outside the graduated cylinder, after gas collection is complete. Which of the Ideal Gas Law’s variables are likely to be affected if you don’t (and how)?
  2. If calcium metal was used as an unknown in an experiment of this type (and it may!), what would be a reasonable value for ‘n’, assuming the sample mass was 0.063 g? Write the balanced chemical equation for the reaction of calcium metal with aqueous hydrochloric acid. Using the value for n determined above, determine the amount of hydrogen gas produced, in mL. Assume temperature and pressure to be the same as they were the day you did the experiment in lab.
  3. Explain how failing to account for any trapped air inside the graduated cylinder at the start of the experiment affects the calculation of: (a) Volume of H2 gas produced (b) Number of moles of H2 gas (c) Molar mass of the unknown metal sample.
  4. What changes (if any) would you make to the experimental procedure if you decided to use aqueous sulfuric acid (H2SO4) instead of aqueous hydrochloric acid (HCl)? Write a balanced chemical equation for the reaction of the unknown metal you were given with aqueous sulfuric acid.

 

 

 

DATA ANALYSIS SUPPLEMENT

Data:

After you have measured the volume of gas collected for all three trials, enter the volumes (mL), the weights (grams), the temperature, atmospheric and water pressures into the appropriate Webbased data entry form, using a computer in the lab. In your measurements and calculations, remember to pay attention to uncertainties and significant figures.

Results: Outside of class, retrieve your data in spreadsheet form (along with that of the entire class) from: http://chem.gmu.edu/results/. Then use Excel to calculate the results as follows:

Your lab report should include the following in addition to the pages from your lab notebook:

  • An Excel sheet “Raw Data” showing the raw data for the entire class (properly formatted). A typical example of the unformatted raw data sheet is shown below. The column headings are also explained.

 

A

B

C

D

E

F

G

H

I

J

1

Number

w1

Vg1

w2

Vg2

w3

Vg3

Temp

Pat

Pw

2

2

0.100

250.0

0.100

260.0

 

 

25.0

760.0

21.0

3

3

0.043

44.0

0.042

45.0

0.044

44.0

24.5

755.8

21.5

 

Explanation:

  • Number = The number of trials: 1, 2, or 3
  • w1 ,w2, w3 = Weight of sample (in g) for each trial (given to 3 decimal places).  0.002 g absolute uncertainty
  • Vg1, Vg2, Vg3 = Volume of gas (in mL) collected in the measuring cylinder for each trial (given to 1 decimal place). 2.0 mL with a 50.0 mL Cylinder
  • Temp (C) = Temperature of water over which the gas is collected (given to 1 decimal place). 1.0 deg C
  • Pat = Measured atmospheric pressure (in torr).
  • Pw = Vapor pressure of water at temperature Temp. Use the interpolated vapor pressure of water or literature value taken from handbook.

 

  • A 2nd Excel sheet “Class Results” showing the calculated results for each student in the class – each in a separate row. A typical example of the unformatted results sheet is shown below:

 

 

A

B

C

D

E

F

G

H

I

J

K

1

Pcorr

(torr)

Pcorr

(atm)

n1

n2

n3

AW1,+

2

(g/mol)

AW2,+

2

(g/mol)

AW3,+

2

(g/mol)

AvgAW,+

2

(g/mol)

AvgAW,+

3

(g/mol)

Metal ID

2

739.0

0.9723

0.0099

0.0103

 

10.1

9.7

 

9.9

14.8

Al

3

740.2

0.9740

0.0018

0.0018

0.0018

23.8

23.8

23.8

23.8

35.8

Mg

 

Explanation:

  • Pcorr = Corrected pressure (in torr). Hydrogen gas pressure corrected for the vapor pressure of water. (= Pat – Pw).
  • Pcorr = Corrected pressure (in atm). (= Pat/760) , (1 atm = 760 torr)
  • n1, n2, n3 = moles of hydrogen produced for each trial =PV/RT = (Pcorr*Vg/1000) / (0.08216*(273.16+Temp))
  • AW1, AW2, AW3 = (w1/n1), etc. This is based on the assumption that the metal takes charge +2, in which case moles H2 = moles of metal.
  • AvgAW,+2 = Average atomic weight of metal, assuming a +2 charge.
  • AvgAW,+3 = Average atomic weight of metal, assuming a +3 charge. (=AvgAW,+2 multiplied by 3/2)
  • To understand the effect of uncertainties in your data on your final results, you may change the weight of the sample and the corresponding volume by their respective absolute uncertainties. In this experiment, for instance, the absolute uncertainty in the weight measurement is 0.002 grams and that in the volume is 2.0 mL. Thus, if you change the weight of your sample, for instance, from 0.045 to 0.047 or to 0.043 grams, then the calculated atomic weight will change by some corresponding amount. Comment on how your final atomic weights are affected by uncertainties in your weight and volume measurements.

Discussion: In your discussion section address the following points:

  • Answer the first three questions on the experiment procedure handout (or any that your instructor assigns).
  • How many of the unknowns are are the same metal as yours?
  • Find the standard deviation of all the samples in the class that are the same metal as yours.
  • Is your own average atomic weight values within experimental error?
  • Which measured parameter (weight, volume, temperature, or pressure) caused the greatest uncertainty in the final result?

 

 

Experiment 9 HEAT OF REACTION: HESS’S LAW

 

 

 

word image 908 word image 909 word image 910 word image 911

Purpose

Using Hess’s Law and simple calorimetry, the enthalpy of formation for magnesium oxide [Hf(MgO)] will be determined to a first approximation.

Background

An understanding of the heat changes that accompany a chemical reaction is of fundamental importance in chemistry. Energy changes that occur during ordinary chemical reactions can be complex. Chemical reactions involve the breaking of chemical bonds in a given set of reactants, and the formation of other different chemical bonds in a given set of products. A thorough examination of the many roles that energy changes have in chemical processes can lead to insights into other chemical phenomena, such as:

  • word image 912 Chemical Kinetics – – the physical mechanisms whereby reactants are converted to products.  Reactions Rates – – the “speed” of a given process that converts reactants into products.
  • Chemical Thermodynamics – – the changes in energy that occur during chemical processes as a function of absolute temperature.

A common measurement used when discussing the relationship between energy and chemical changes is the enthalpy change for a chemical process, H. The enthalpy change of a chemical reaction is defined as the amount of heat exchanged by that chemical reaction conducted at a constant pressure. It is a measure of the difference between the “heat contents” of the products and the reactants.

ΔHrxn = ΔHfinal – ΔHinitial = ΔHproducts – ΔHreactants

A chemical reaction occurs within a focused finite part of the universe, an environment called a “system.” Any energy released in the reaction or absorbed in order to initiate the reaction comes from an external environment called the “surroundings.” Thus, the reactants and products constitute the system and everything else, such as the reaction container, room, etc. make up the surroundings.

The sign of the quantity H indicates the general direction of the energy flow into or out of a reaction system. If H has a negative sign, heat has been transferred from the system to the surroundings in an “exothermic” reaction. If H has a positive sign, heat has been transferred from the surroundings to the system in an “endothermic” reaction.

Enthalpy, along with “internal energy (E)” and “entropy (S)” are referred to as being “state” functions. In thermodynamics, a state function, function of state, state quantity, or state variable is a property of a system that describes quantitatively the equilibrium of the system irrespective of how the system arrived in that state. Thus, the temperature, pressure, and volume of a system would also be considered state functions.

Enthalpy is also an “extensive” property, which means the amount of energy exchanged in the reaction is a function of the amount of substance being investigated. Enthalpy is itself a general term for heat transfer. In the context of chemical reactions, it is common to be able to measure the following types of heat transfer.

  • Enthalpy of Formation (Hf): The quantity of heat involved when forming 1 mole of a given substance in its standard state, directly from the elements that comprise the substance in their standard states (STP).
  • Enthalpy of Combustion (H): The quantity of heat transferred per mole of a combustible substance, upon its reaction with excess oxygen, i.e., burning.
  • Enthalpies of Solution (Hsol), Vaporization (Hvap), Fusion (Hfus), or Sublimation (Hsub): The quantity of heat involved when a substance changes its physical state (phase) as it melts, vaporizes, dissolves in water, or changes directly from a solid to a gas.
  • Enthalpy of Neutralization (Hn): The quantity of heat involved when 1 mole of water is produced by the reaction of an aqueous acid and an aqueous base, in water.

Enthalpy, Internal Energy, and Heat

Enthalpy is a measure of the total energy of a thermodynamic system; the energy transferred between the system and its surroundings. It includes the internal energy, E, which is the energy required to create a system, and the amount of energy required to make room for it by displacing its environment and establishing its volume and pressure, i.e. the work (w) involved.

ΔH = ΔE + w = ΔE + PV

The internal energy of a system, E, is precisely defined as the heat at constant pressure (qp) plus any work (w = PV) done by the system.

ΔE = q + w = q + (- PV)p p

Internal Energy used to expand volume by increasing pressure is lost to the surroundings, thus the negative sign in the work term.

Thus, q = p ΔE + PV and ΔH = qp at constant pressure.

“Heat content” is the chemical energy possessed by a substance due to the nature of its chemical bonds and its state (pressure, temperature, volume). In an exothermic reaction, the reactants have a larger heat content than the products and t he excess heat is liberated during the reaction; thus Hrxn is a negative quantity. Conversely, in an endothermic reaction, the reactants have a smaller heat content than the products, heat is absorbed during the reaction, and Hrxn is positive. See Figure 1.

word image 913

This experiment is designed to test and verify Hess’s Law; which stated in words says: The total energy change for any process that consists of a series of steps is equal to the sum of the enthalpies of the individual steps. Thus for a given reaction that can be expressed as the sum of four intermediate reactions, Hess’s Law would look something like this:

reaction #1 ———– > H1 reaction #2 ———– > H2 reaction #3 ———– > H3 reaction #4 ———– > H4 Htot = sum of H values for individual reactions = H1 + H2 + H3 + H4

Heat of Reaction and Heat of Formation

The total enthalpy change for a reaction involving several steps or multiple reactions requires the determination of the individual H values. This can be done in two ways:

  1. Measuring the heat of reaction using a calorimeter.
  2. Tabulating the heat of reaction for a compound formed from its uncombined elements.
Heat of Reaction (experimental)

In order to measure the heat of reaction of a substance utilizing non-standard amounts of reagents, it is necessary to determine the actual heat generated in the reaction. The total heat change (qp) for a reaction carried in a calorimeter at constant pressure is defined as the product of the mass of material present (m, in grams), the specific heat capacity of that material (Cp, in J/g oC), an the change in temperature (T) during the reaction.

 

𝑞𝑟𝑥𝑛 = 𝑚𝐶𝑝ΔΤ

The heat value, qrxn, is based on the actual amount of material present. The enthalpy change of the reaction (Hrxn) is based on the amount of heat generated per mole of substance. Thus, under conditions of constant pressure, as in this experiment:

qrxn

ΔHrxn =

# moles limiting reagent

Example:

How much heat is generated when 0.500 g calcium oxide (CaO) is placed in 50.0 mL of water (H2O) in a Styrofoam cup at 23.5oC and the temperature increased to

27.0oC?

balanced chemical equation:

CaO(s) + H O(l) 2  Ca(OH) (aq)2

mass of calcium = 0.500 g

mass of water is calculated from its volume and density (0.998 g/mL) mass of water = 50.0 mL * 0.998 g/mL = 49.9 g

mass of the system is the mass of calcium plus the mass of water

mass of system = (0.500 + 49.9) g = 50.4 g temperature change (T) = 27.0oC – 23.5oC = 3.5oC specific heat of water = 1.00 cal/g deg = 4.184 J/g deg heat of reaction: qrxn = m Cp T

4.184 J

qrxn = 50.4 g  3.5 deg = 738  J = 0.738 kJ g•deg

The “Heat of Reaction” (Hrxn in kJ/mol), is the “Heat Evolved” adjusted to reflect the amount of heat that would have been generated by 1 mole of the reactant. heat evolved (kJ)

ΔHrxn = × mol wgt calcium oxide (g / mol) = kJ / mol

mass of calcium oxide (g)

0.738 kJ

= × 56.08 g / mol = 82.8 kJ / mol

0.5000 g CaO

Since heat was liberated to the surroundings, an exothermic reaction, the products must contain less energy than the reactants. By convention then, Hrxn must be written as 82.8 kJ/mol.

Heat of Formation (tabulated)

If one or more of the intermediate reactions involves the formation of a compound from its uncombined elements, the heat of formation of the compound can be computed from Standard Heats of formation, symbolized by H𝑓𝑜, that have been derived and tabulated in reference lists (see the back of your general chemistry text book). Recall that the heat of formation for an uncombined element is by definition, zero (0).

The standard enthalpy of formation or standard heat of formation of a compound is defined as the change of enthalpy from the formation of 1 mole of the compound from its constituent elements, with all substances in their standard states at 1 atm (101.3 kPa) and 25oC (298 K).

For example, the heat of reaction to produce calcium oxide from its elements calcium and oxygen is just the heat of formation of calcium oxide.

Cao(s) + 1O2(g) ® CaO(s)

2

∆Horxn = å∆Hof (product – CaO) – å∆Hof (reactants – Ca + O)

∆Horxn = -635.1kJ / mol – (0+0) = -635.1 kJ / mol

The negative sign for the heat of formation value indicates the reaction is exothermic; energy is released to the surroundings. The other compound needed in our original reaction is water. Its heat or formation from hydrogen and oxygen is -285.8 kJ/mol. Thus, the heat of formation for Ca(OH)2 (aq) can now be computed.

CaO(s) + H O(l) 2  Ca(OH) (aq)2

H𝑓𝑜 (kJ/mol) 635.1 285.8 1002.8

H word image 914 (kJ/mol) = [(1002.8)  (635.1  285.8)] kJ/mol = 81.8 kJ/mol

The difference between the tabulated and experimental heats of reaction is due to experimental error. A percent error can be calculated.

The Experiment:

The goal of this experiment is to determine the heat of formation of magnesium oxide (MgO), using calorimetry and Hess’s Law. The experimental result will be compared to the known heat of formation word image 915 ) and the percent error calculated.

The Hess’s Law components will consist of three reactions. The first two will involve the calorimetric determination of the heats of reaction for the reaction between magnesium (Mg) metal and hydrochloric acid and the reaction between magnesium oxide (MgO) and hydrochloric acid.

Mg(s) + 2HCl(aq)  MgCl2 (aq) + H2(g) Hrxn (A) (1)

MgO(s) + 2HCl(aq)  MgCl2(aq) + H2O(l) Hrxn (B) (2)

The 3rd reaction requires the standard heat of formation, H𝑓𝑜 (kJ/mol) for the formation of water from hydrogen and oxygen.

H2 + ½ O2  H2O Hf(H2O) (3)

The net reaction for the formation of magnesium oxide (MgO) from magnesium (Mg) and oxygen (O) is obtained from the summation of reactions 1 ,2, and 3. Before the summation can be done, the equations must balanced and, if necessary, reversed in order to cancel out all intermediate components that do not participate in the net equation. The only change in this regard is to reverse equation 2. As a result of this reversal, the sign of the Hrxn (B) value must also be reversed.

Mg(s) + 2HCl(aq)  MgCl2 (aq) + H2(g) Hrxn (A) (1)

MgCl2(aq) + H2O(l)  MgO(s) + 2HCl(aq) Hrxn (B) (2)

H2 + ½ O2  H2O(l) H𝑓𝑜 (H2O) (3)

Mg(s) + ½ O2 (g)  MgO(s) H𝑓𝑜 (MgO) (4)

 

The Heat of Reaction for the formation of MgO, H𝑓𝑜 (MgO), from oxygen and magnesium is computed from the individual heats of reaction from reactions 1, 2, 3.

(H)f (MgO) = H1 + H2 + H3

Note: The reactions of both Mg and MgO with HCl result in the release of hydrogen and heat; thus, they are exothermic reactions and the Hrxn values are negative. However, since reaction (2) was reversed, the original Hrxn (2) value must also be reversed, i.e., it is now positive.

The standard heat of formation for water, H𝑓𝑜 (H2O) (3), is obtained from standard reference tables. Its value is -285.8 kJ/mol

Since H usually does not change significantly with temperature and the data will be obtained at close to standard conditions (1atm, 25oC), H’s and Ho’s can be used interchangeably.

Sample Calculation:

Assume that the H(1) and H(2) values from the calorimeter measurements for the experiment were -496 kJ/mol and -195 kJ/mol, respectively. Taking into account the reversal of the 2nd reaction, the overall Hess’s Law expression for the Hf of magnesium oxide from its elements would look like the following:

 

Hf (MgO) = H1 + (-H2) + H3

Hf (MgO) = 496 kJ/mol + 195 kJ/mol + (285.8 kJ/mol)

Hf (MgO) = 586.8 kJ/mol (actual value = 601.8 kJ/mol)

 

_______________________________________________________________

Materials and Equipment:

Materials Equipment

Styrofoam cups (2), with plastic cover thermometer

Weigh boat

 

metal spatula or glass stirring rod

magnesium metal turnings

 

electronic balance

magnesium oxide powder

 

Ring stand assembly and clamps

hydrochloric acid (include concentration)

 

 

 

 

Procedure:

  1. Obtain 2 Styrofoam cups and plastic cover with hole.
  2. Form a calorimeter by placing one cup into the other cup.
  3. Add about 100 mL (precisely measured to nearest 0.1 mL) of 1.00 M HCl to the calorimeter.

Note: Hydrochloric acid is in excess and 100 mL of the acid should be sufficient for all four samples (two samples of magnesium metal and two samples of magnesium oxide).

  1. Cover the calorimeter and place a thermometer through the hole in the cover.
  2. Record and continue to monitor the temperature of the solution.
  3. Weigh approximately 0.2 g of Magnesium (Mg) metal precisely measured to the nearest 0.001 g.
  4. Add the metal to the calorimeter all at once and quickly cover the calorimeter.
  5. Stir the mixture with gentle swirling.
  6. Monitor the temperature of the solution until a temperature maximum has been reached.
  7. Record the final temperature.
  8. Repeat this process for a 2nd weighed sample of Magnesium .
  9. Weigh out a sample of about 0.5 g of Magnesium Oxide (MgO) precisely measured to the nearest 0.001 g.
  10. Add the MgO to the HCl solution.
  11. Stir the mixture with gentle swirling.
  12. Monitor the temperature of the solution until a temperature maximum has been reached.
  13. Record the final temperature.
  14. Repeat the process for a 2nd sample of MgO.
  15. Clean out the reaction vessel, flushing the solution down the drain with water.

 

Calculations:

Compute the heat of reaction (qrxn) for each trial.

𝑞𝑟𝑥𝑛 = 𝑚𝐶𝑝ΔΤ

Where: m = mass of the system (metal + HCl soln)

Cp = Specific Heat of HCl soln (water) = 4.184 J/g-deg

T = Change in temperature

Note: The aqueous hydrochloric acid solution is the absorbing mass in the calorimeter, but it can be assumed that the specific heat is the same as for water, i.e., both water and HCl absorb the same amount of heat per gram of their mass for each degree of temperature change.

Compute the moles of each reactant.

mass moles =

mol wgt

Verify that the limiting reagent in each reaction is magnesium and magnesium oxide, respectively.

  1. Reaction molar ratio: Mg : HCl = 1:2
  2. Actual molar ratio

moles magnesium actual molar ratio =

moles HCl

  1. If the actual molar ratio < reaction molar ratio; then Mg is limiting

Compute the enthalpy of the reaction (Hrxn) for each trial

qrxn

ΔHrxn =

# moles limiting reagent Compute the average Hrxn for magnesium and magnesium oxide.

Compute the heat of reaction for the formation of MgO, H𝑓𝑜 (MgO),

H𝑓𝑜 (MgO), = sum of Hs of individual reactions = H1 + H2 + H3

H𝑓𝑜 (MgO) = H (1) + H(2) + H𝑓𝑜 (3)

 

  

TABLE

 

 

H°f

 

H°f

Compound

kJ mol−1

Compound

kJ mol−1

HCl (aq)

−167.2

MgCl2(aq)

−801.2

H2O (l)

−285.8

CaO (s)

−635.1

CO2 (g)

−393.5

CaCl2 (aq)

−877.1

CH4 (g)

−89.0

NO2 (g)

+33.9

C8H18 (l)

−250.0

NO (g)

+90.4

 

__________________________________________________________________________________________________

HEAT OF REACTION: HESS’S LAW

__________________________________________________________________________________________________

Density 1.0-M HCl Solution 1.016 g/mL Specific heat of HCl Solution ____________

 

A 1st

A 2nd

B 1st

B 2nd

vol. HCl solution

 

 

 

 

mass HCl Solution

 

 

 

 

mass Mg or MgO

 

 

 

 

Tinitial

 

 

 

 

Tmax

 

 

 

 

Tchange

 

 

 

 

Heat evolved (kJ)

 

 

 

 

H (kJ/mol)

 

 

 

 

Average

 

Hrxn(A)

 

Hrxn(B)

 

Chemical Equation for Reaction A:

Chemical Equation for Reaction B:

Chemical Equation for Reaction A – Reaction B:

Hrxn = Hrxn(A) – Hrxn(B) = kJ/mol

 Hf (H2O l) [Table]

Experimental Hf (MgO s ) (Show calculation)

Hf (MgO s) [Accepted value] % Error __________

DATA ANALYSIS SUPPLEMENT

Data: After you have noted down all the masses, volumes and temperatures, enter your own data into the appropriate Web-based data entry form, using a computer in the lab. In your measurements and calculations, remember to pay attention to uncertainties and significant figures.

Results: Outside of class, retrieve your data in spreadsheet form (along with that of the entire class) from: http://chem.gmu.edu/results/. Then use Excel to calculate the mass of HCl solution, Tchange, Heat evolved (kJ), H (kJ/mol), Average Hrxn (A), Average Hrxn (B), Hrxn, Experimental Hf (MgO s), %Error.

Your lab report should include the following in addition to the pages from your lab notebook:

  • An Excel sheet “Raw Data Sheet” showing the raw data for the entire class. A typical example of the raw data sheet is shown below:

Raw Data Sheet

 

 

A

B

C

D

E

F

G

H

I

1

ID

V1_Mg mL

M1_Mg

g

initT1_Mg

°C

FnlT1_Mg

°C

V2_Mg mL

m2_Mg

g

initT2_Mg

°C

FnlT2_Mg

°C

2

1

100.0

0.206

21.5

29.9

100.0

0.220

21.0

30.4

3

2

100.0

0.236

22.0

32.5

100.0

0.222

21.9

31.1

 

 

 

 

J

K

L

M

N

O

P

Q

1

V1_MgO mL

M1_MgO

g

initT1_MgO

°C

fnlT1_MgO

°C

V2_MgO mL

m2_MgO

°C

initT2_MgO

°C

fnlT2_MgO

°C

2

100.0

0.501

21.0

24.5

100.0

0.535

21.0

24.7

3

100.0

0.500

21.5

24.8

100.0

0.504

22.0

25.0

 

Explanation:

  • V1_Mg, V2_Mg=volume of HCl in reaction A (trial 1 and trial 2).
  • m1_Mg, m2_Mg=mass of magnesium in reaction A (trial 1 and trial 2).
  • initT1_Mg, initT2_Mg=initial temperature in reaction A (trial 1 and trial 2).
  • fnlT1_Mg, fnlT2_Mg=final temperature in reaction A (trial 1 and trial 2).
  • V1_MgO, V2_MgO= volume of HCl in reaction B (trial 1 and trial 2).
  • m1_MgO, m2_MgO= mass of magnesium oxide in reaction B (trial 1 and trial 2).
  • initT1_MgO, initT2_MgO= initial temperature in reaction B (trial 1 and trial 2).
  • fnlT1_MgO, fnlT2_MgO= final temperature in reaction B (trial 1 and trial 2).
  • A 2nd Excel worksheet “Results Mg Sheet” and a 3rd Excel worksheet “Results MgO Sheet” showing the calculated results for each student in the class – each in a separate row. Make sure to show an example calculation for each column. This can be done by copying the equation from the formula bar to a blank cell below the elements in the table. An apostrophe in front of it makes it text instead of an equation. Each table should be formatted as in previous classes with lines separating the table headings from the contents. A line below the contents should also be drawn. A typical example of the results sheet is shown below:

Results Mg Sheet

 

A

B

C

D

E

F

G

H

I

J

K

L

1

ID

mHClsoln1

g

molMg1 mol

dt1,

°C

mHClsoln2

g

molMg2

Mol

dT2

°C

heat1 kJ

heat2 kJ

dH1 kJ/mol

dH2 kJ/mol

dH.rxn.A kJ/mol

2

1

101.6

0.0085

8.4

101.6

0.0091

9.4

3.6

4.0

-421.9

-442.1

-432.0

3

2

101.6

0.0097

10.5

101.6

0.0091

9.2

4.5

3.9

-460.5

-428.8

-444.7

Explanation

  • mHClsoln1, mHClsoln2 =mass of HCl solution for trials1 and trial2
  • molMg1, molMg2 =moles of Mg
  • dT1,dT2=temperature change in trial1 and trial2.
  • heat1 =heat evolved in kJ
  • dH1,kJ/mol = heat evolved in kJ/mol
  • dH.rxn.A, kJ/mol = average heat of reaction for reaction A

 

Results MgO Sheet

 

A

B

C

D

E

F

G

H

I

J

K

L

M

N

O

1

ID

mHClsoln

1

G

molMgO1

moles

dT1

°C

mHClsoln2

g

molMgO

2

moles

dT2

°C

heat1

kJ

heat2

kJ

dH1

kJ/mo

l

dH2

kJ/mo

l

dH.rxn.B

,

kJ/mol

dH.rx n

kJ/mo

l

dH.f.MgO

,

KJ/mol

%

Error

2

1

101.6

0.0124

3.5

101.6

0.022

3.7

1.5

1.6

-120.2

-71.8

-96.0

-336.0

-621.8

3.43

3

2

101.6

0.0124

3.3

101.6

0.0207

3

1.4

1.3

-113.5

-61.8

-87.7

-357.0

-642.8

6.92

 

Explanation

  • mHClsoln1, mHClsoln2 =mass of HCl solution for trials1 and trial2
  • molMgO1, molMg2 =moles of MgO
  • dT1, dT2=temperature change in trial 1 and trial 2
  • heat1 = heat evolved in kJ
  • dH1,kJ/mol = heat evolved in kJ/mol
  • dH.rxn.B, kJ/mol = average heat of reaction for reaction B
  • dH.rxn, kJ/mol = difference between heat of reaction A and heat of reaction B
  • dH.f.MgO, kJ/mol = experimental heat of formation for MgO
  • % Error = Error % of the heat of formation of MgO compared with the literature value.

 

  • An Excel chart showing a histogram of the distribution of experimental Hf (MgO s) results. Use equally spaced “bins” chosen to display the “spread” in the class results. Refer to the histogram instructions on the website.

Printing Tables:

  • Large tables should be printed in the landscape mode and sometimes it might be necessary to go to “page setup” and select “Fit to 1 page” to make the page fit on one page.
  • Include column and row headings on the printout so that the instructor can more easily understand equations used. This can be done by going to File → Page Setup → Sheet tab → Select “Row and Column Headings”.

Discussion: In your discussion section address the following questions:

  • What does the Hess’s Law state?
  • Why is the calorimeter assembled by placing one cup inside the other?
  • What could be a reason for the difference between the theoretical and experimental heats of reaction?
  • Is your answer within the expected experimental error relative to the theoretical results?
  • What are the possible sources of error in this experiment?
  • From the histogram explain how the lab section’s results compare to the theoretical values.
  • How do your results compare with the lab section’s results?

 

Experiment 10 SOLVENT POLARITIES FROM A SOLVATOCHROMIC DYE[2]

word image 916 word image 917 word image 918 word image 919 word image 920

 

word image 921 The color of solutions of certain compounds depends upon the solvent into which it is dissolved. Brooker’s merocyanine (Figure 1) is one of them. This is due to the effect the solvent has on which resonance form is dominant in the equilibrium below. The change in the color, commonly known as solvatochromism, is useful in studying the polarity of different solvents. Small changes

word image 922

Figure 1 Two resonance forms of Brooker’s Merocyanine in the polarity of the solvent can have observable changes in the color. The reason for this color change can be understood by studying the energy levels in Figure 2.

Looking first at the energy level on the left (nonpolar), an electron is normally found in the low energy state. When a photon that has an energy equal to the difference in energy between this state and one of the excited states, such as the one shown (ET,n), the energy is absorbed and the electron moves to the higher energy state. If the same compound is placed in polar solvent (right energy level diagram), the energy levels change as a result of the interaction with the solvent. The ground state energy changes the most. As a result, the energy required to move an electron in this solvent is higher (ET,p). Since ET,p > ET,n, the color corresponding to that energy shifts to a shorter wavelength.

word image 923

Energy

 

word image 924

Nonpolar

 

word image 925

E

T, n

 

word image 926

Polar

 

word image 927

E

T, p

 

word image 928

Figure

2

 

Comparison of the

Transition Energies, E

T

 

for a

nonpolar and polar solvent

 

In this experiment, the effect of different solvents on the color of the solution will be studied two ways: (1) by visually observing the color of the solution and (2) using a spectrophotometer. A spectrophotometer is an instrument that is used to measure how much light is absorbed at different wavelengths.

Light, which is characterized by its wavelike behavior, has three components which must be taken into consideration: wavelength ( ), frequency ( ),and the speed (c) at which it travels.. The energy of each photon in light is given by equation (1)

ℎ ∙ 𝑐 (1)

word image 929

This equation tells us that the energy of a photon is directly proportional to the frequency and inversely proportional to the wavelength. High frequency or short wavelength photons have a high energy, whereas low frequency or long wavelength photons have a low energy. In this experiment, you will use wavelength to determine the relative energy of a photon. Since the energy of a photon is inversely proportional to the wavelength, the energy of photons at the low end of our scan range (400 nm) will have a higher energy than photons at the high end of the wavelength scan range (700 nm).

Figure 3 shows the range of wavelengths and frequencies for light. Photons at high frequency (on the left in the figure) have a very high energy and are the result of light being emitted

 

 

 

word image 930

 

 

 

word image 931

 

 

word image 932

 

 

word image 933

 

 

word image 934

 

 

word image 935

 

 

word image 936

 

 

 

word image 937

 

 

word image 938

 

 

word image 939

Figure 3 Electromagnetic spectrum. Visible portion is the narrow band in the middle.

from high energy processes. Gamma rays, shown on the far left, are produced from nuclear events, such as fusion; x-rays are produced from high energy electronic transitions in atoms. Lower energy transitions in the visible, infrared, radio, etc. are derived from much lower electronic transitions in molecules.

word image 940

Figure

4

 

Absorption of radiation

through a solution. The color we

observe is the one that is not

word image 3218 When light passes through a solution, some of the light can be absorbed by the solution. This will happen when the energy of that light matches the energy required for an electron to move from a ground state to a higher energy state. The light is used up and thus the number of photons traveling through the solution decreases – light intensity decreases. As a result, the light that exits appears to be a different color than the light that entered. Consider the diagram in Figure 4. Blue and red light go into the solution, but only red makes it through; so the solution appears to be red. In this experiment, we will use a white light source and measure how much light is absorbed at several wavelengths to determine which wavelength is mostly absorbed. Our instrument measures what is absorbed, whereas our eyes observe the relative amount that remains. We can roughly determine what color was absorbed by using the color wheel in Figure 5. The color absorbed will be the observed color’s complement. For example, if we see green, then red will be absorbed by the solution.

The color absorbed is characterized by a

wavelength of absorption as represented in equation (2) below. In this setup, we have written the equation for

one mole of photons by including Avogadro’s number. Figure 5 A traditional color wheel. Each

color serves as the complement of the

word image 941opposite color across the wheel. E.g. if

(2) purple is absorbed, ellow is observed.

where NA is Avogadro’s number, h is Planck’s constant, and c = 3.0×108 m/s (the speed of light in a vacuum).

An instrument that measures the amount of light absorbed is known as a spectrometer. There are a variety of designs that can do this, with the more expensive ones allowing a more automated operation. The controls can be adjusted to allow only a certain wavelength of light pass through. The amount of light passing through the cell can then be measured and displayed as % transmittance (T) – the relative amount of light coming through the spectrometer – or absorbance (A), which is given by the following equation:

word image 942 (3)

word image 3219 word image 943

Wavelengt

h

 

word image 944

Absorbance

 

word image 945

 

word image 946

max

 

word image 947

Figure

6

 

Example of absorbance vs wavelength. The point where there is

maximum absorbance is referred to as

max

 

where Po is the power going into the sample and P is the power coming out.

When the wavelength is varied over a range and absorbance (A) or % transmittance (T) measured, a plot can be made showing how the absorbance changes with wavelength, commonly referred to as an absorption spectrum. The figure below shows how this is done. The wavelength of maximum absorption will be in this is experiment will be measured in this way for the different solutions.

Experimental

The Genesys 20 Spectrometer

96

 

word image 3220 word image 948

Figure

7

 

Genesys 20 Spectrophotometer

 

The most commonly used spectrometer for general chemistry laboratories is the Genesys 20. Your instructor will demonstrate how to use the the instrument. You will use this instrument to determine the wavelength of maximum absorption of the solutions for this experiment.

 

Cuvettes

Cuvettes (or cells) are sample tubes of precise dimensions, made from special glass to ensure uniform transmittance of light. You will need at least two — one for the solution and the other for a reference blank. The composition of glass may vary from one manufacturer to another or from batch to batch. Check to see that both cuvettes are the same brand, such as Pyrex or Kimax.

Wash and rinse cuvettes with distilled water. Rinse a cuvette with a small amount of the sample solution and then add fresh sample solution to measure the absorbance. Wipe the outside of the cuvette with a tissue and handle the cuvette only on its top sides.

Calibrating the Instrument
  1. Let the instrument warm up for at least 20 minutes.
  2. Use the Wavelength and buttons to adjust the wavelength to the appropriate one for the experiment
  3. Insert the cuvette that ½ full with your “blank” solution into the instrument
  4. Not sure what your blank solution is?
  5. A blank solution is a solution containing little to no analyte of interest
  6. In this course, the blank solution is usually the solvent you will be using
  7. Still not sure? Ask your instructor
  8. Press the “0 ABS/100% T” Button
  9. The instrument is calibrated
  10. Press the “A/T/C” Button to select the proper data output for the experiment. Absorbance (A) is the recommended setting for this experiment
Taking an Absorbance reading

Place a cuvette containing your sample in the instrument, close the lid, and record the absorbance and wavelength.

The instrument must be calibrated every time you change the wavelength. This is not necessary if you merely change the sample but not the wavelength. A substance’s “absorptivity factor” is a constant only as long as the wavelength of light is constant.

Solutions of 1.0×104 M Brooker’s merocyanine have been prepared in each of five solvents. The solvents you will use are methanol (CH3OH), 1-propanol (CH3CH2CH2OH), 2-propanol

(CH3CHOHCH3), acetonitrile (CH3CN), and acetone (CH3COCH3). You will determine the wavelength of maximum absorption for merocyanine in each solution. Based on your experimentally determined value for the wavelength of maximum absorption, also known as lambda max (max ), you will determine the energy of the transition. The end result will lead to a determination of the relative polarity of each solvent.

  1. Obtain 5 cuvettes, each containing merocyanine dissolved in one of the solvents. Make sure the cuvettes are properly labeled indicating which solvent is being used.
  2. Observe the color of each solution and qualitatively determine the relative polarity of each solvent. Record the observed color of each solution directly into your lab notebook
  3. Using a UV-Vis spectrophotometer, determine the wavelength of maximum absorption, max, for each solution by scanning from 400 nm to 700 nm.
    1. Set the wavelength to 400 nm.
    2. Zero the spectrometer using water in the cuvette as a blank.
    3. Measure the absorbance of each of the merocyanine solutions.
    4. Increase the wavelength by 25 nm and repeat the absorbance readings.
    5. Continue increasing the wavelength by 25 nm and taking absorbance readings until you have reached 700 nm.

 

Results:

  • Make tables in your notebook like the ones below.
  • Use Excel to plot the absorbance at different wavelengths for each solution. Find the max for each solution as shown in Figure 4
  • Calculate the ET of the transition for each solution using the max from equation (2). Since the energy of the transition indicates the effect of the polarity of the solvent, use the magnitude of this value to decide upon the relative polarity of each solvent. Does this result agree with your estimate based upon the color you observed in step 2 above? Explain the cause of any discrepancy.
  • Use the visually observed color and the color wheel (Figure 5) to determine the color absorbed. Based upon the colors absorbed for each solution and using the electromagnetic spectrum shown in Figure 3, determine the relative energies of the absorption occurring in the different solutions and rank the polarities from 1 for the most polar to 5 for the least polar.
  • ET in table 3 should be calculated using equation (2). Make sure to express answers in kJ/mol.
  • Using your experimentally determined values for ET , rank the solvents in order of relative polarities ranging from 1 for the most polar to 5 for the least polar.
Table 1 Observed color and estimated relative polarity

Solvent

Observed Color

Color absorbed

Relative Polarity (based on color absorbed)

Methanol, CH3OH

 

 

 

1-propanol, CH3CH2CH2OH

 

 

 

2-propanol, CH3CHOHCH3

 

 

 

Acetonitrile, CH3CN

 

 

 

Acetone, CH3COCH3

 

 

 

 

Table 2 Absorbance of the solutions at different wavelengths

Wavelength, nm

Absorbance for the solution with the following solvent

 

methanol,

CH3OH

1-propanol, CH3CH2CH2OH

2-propanol, CH3CHOHCH3

Acetonitrile,

CH3CN

Acetone,

CH3COCH3

400

 

 

 

 

 

425

 

 

 

 

 

450

 

 

 

 

 

475

 

 

 

 

 

500

 

 

 

 

 

525

 

 

 

 

 

550

 

 

 

 

 

575

 

 

 

 

 

600

 

 

 

 

 

625

 

 

 

 

 

650

 

 

 

 

 

675

 

 

 

 

 

700

 

 

 

 

 

 

Table 3 Transition energies and relative polarities of Merocyanine

Solvent

max

ET (kJ/mol)

Relative solvent polarity (1-5)

methanol, CH3OH

 

 

 

1-propanol, CH3CH2CH2OH

 

 

 

2-propanol, CH3CHOHCH3

 

 

 

Acetonitrile, CH3CN

 

 

 

Acetone, CH3COCH3

 

 

 

 

 

Conclusions:

  1. Explain if the ordering from visual analysis of the relative energies agree with the energy ranking from actual measurement of the energies using the Genesys 20. Explain why there might be a discrepancy.
  2. Compare your results with what the literature reports (see reference at the beginning of the experiment).
  3. A shift to longer wavelengths upon absorption in some solvents is called a bathochromic shift. A hypsochromic shift occurs when the shift is to shorter wavelength. Which of the two categories did you observe during your experiment? What else is this shift often called?
  4. Hand draw an energy level diagram in your laboratory book illustrating a hypsochromic shift.
  5. Compare the relative polarities you obtained for your solvents with the values shown in the table at https://sites.google.com/site/miller00828/in/solventpolarity-table
  6. Let’s say that you have a solvatochromic compound that appears red in a solvent. You dissolve the compound in another solvent and it changes to an orange color. Would this be considered a hypsochromic or bathchromic shift? Would the solvent be more polar or less than the first solvent? Explain.

 

 

Experiment 11 VSEPR AND MOLECULAR MODELING

 

 

 

 

word image 949 word image 950 word image 951 word image 952 word image 953

The “experiment” described here is designed to help you understand molecular shapes, that is, the spatial arrangement of the atoms in a molecule. This is an important concept in chemistry because many properties of a compound can be understood in terms of its molecular geometry. In particular, the molecular geometry, along with the electronegativities of the atoms in the molecule, determine whether a compound is polar or non-polar. This, in turn, largely determines how a molecule will interact with its environment. For example, the polar character of a molecule affects many of its physical properties such as solubility, melting point, and boiling point.

word image 954 One of the difficulties you may face in understanding molecular geometries and learning to predict them is that molecules are three-dimensional. This means that representations of molecules on the (two-dimensional) page are inherently misleading. Unfortunately, we are unable to actually “see” the spatial arrangement of the atoms in an individual molecule; atoms are much too small. Molecular shapes can be deduced experimentally, but only indirectly, using rather sophisticated methods such as spectroscopy.

Three different representations of the simple molecule methanol, CH4O, are shown below: a line drawing, a ball-and-stick model, and a space-filling model.

 

H

 

word image 3221 word image 3222

C

O

H

H

H

 

 

Today you will study the geometry of molecules by constructing ball-and-stick models. Modelbuilding is an important aspect of science; often the best way to understand a complicated, real system is to study an idealization of it. In this case we are using simple, solid balls to represent complex atoms and “sticks” to represent directional covalent bonds. This sort of model is useful, not only to the novice trying to visualize the three dimensional shapes of simple molecules, but even to the experienced chemist approaching an unfamiliar, complex molecule.

Valence Shell Electron-Pair Repulsion Theory, or VSEPR Theory is one method to predict the shape of a molecule from a knowledge of its chemical formula. An atom is designated the “central” atom and the theory predicts the arrangement of the atoms (or groups of atoms) about the central atom to which they are all covalently bonded. The atoms or groups of atoms that are bonded to the central atom are called the ligands. In the methanol example above, if the geometry around the carbon atom is to be determined, the three hydrogens and the OH group are the attached ligands.

The basic idea of VSEPR theory is that the arrangement of the ligands about the central atom is determined by the repulsions among the electron pairs around the central atom. Electron pairs include both the non-bonding electron pairs belonging to the central atom as well as the bonding electron pairs connecting the ligands to the central atom. Because of the mutual repulsion of the electrons’ negative charges, these electron pairs will distribute themselves about the central atom in such a way as to be as far from one another as possible. This gives rise to distinctive geometries that depend solely on the number of electron pairs around the central atom.

The notation used to distinguish the bonding and non-bonding pairs is ALxNy, where A stands for the central atom, L for the ligands, x for the number of ligands (and thus bonding pairs — ignore any multiple bonding pairs such as double- or triple-bonds) around the central atom, N for the non-bonding electron pairs, and y for the number of non-bonding pairs on the central atom. The geometric arrangement of the electron pairs about the central atom is determined by the sum (x+y).

The possible geometric arrangements of the electron pairs (bonding and non-bonding pairs) around a central atom and the number of electron pairs required to give rise to a particular arrangement are shown on the next page.

For example, :NH3 (ammonia) has a nitrogen atom in the center of the molecule connected to three hydrogen atom ligands. Nitrogen also possesses a pair of non-bonded electrons. Its notation is therefore AL3N1. The geometry of the ammonia molecule will result from the mutual repulsion of 4 pairs of electrons — the 3 bonding and 1 non-bonding pairs of electrons. Ammonia will therefore have a tetrahedral geometry where the three hydrogen atoms and the non-bonding pair are located at the vertices of a tetrahedron. In the line drawings below, the solid wedge line indicates that the ligand is pointing out of the plane toward the viewer and the dashed line means the ligand is pointing out of the plane away from the viewer. The simple solid line represents a ligand in the plane of the paper.

H

N

N

H H H or H H

The tetrahedron is perhaps the most difficult of the geometric forms to represent on paper. This is one of the reasons you will be constructing the three-dimensional models.

TABLE 1

(x + y)

# of Electron Pairs on Central Atom

Geometric Arrangement of Electron Pairs around the

Central Atom

______________________________________________________________________________

word image 3223

  1. linear A

 

 

 

 

A

word image 3224

  1. trigonal

planar

 

 

 

 

A

word image 3225

  1. tetrahedral

 

 

 

 

A

word image 3226

  1. trigonal

bipyramidal

 

 

6

 

octahedral

 

A

 

 

 

word image 3227

 

Look back at the methanol models presented in the introductory material. Recall the geometry around the carbon is tetrahedral. What is the geometry around the oxygen as the central atom? Neither the ball-and-stick nor the space-filling model explicitly show the non-bonding electrons, but it is clear that they influence the bond angles between the atoms.

For linear, trigonal planar, tetrahedral, and octahedral geometries, all vertices in the geometrical figure are equivalent. This means that the size of the ligand or lone pair has no effect on which vertex it will occupy. Every vertex provides the same amount of space. However in the case of trigonal bipyramidal, the five vertices are not equivalent. Instead, the vertices are of two different classes — axial and equatorial. There are two axial vertices, one directly above the central atom and one directly below. Their positions correspond to the locations of the north and south poles on the globe, with the central atom corresponding to the center of the earth. The remaining three vertices, which are equatorial, can be viewed as lying on an imaginary equator, hence the name. The three equatorial positions form the corners of an equilateral triangle (equatorial plane).

An understanding of molecular geometry is also necessary to understand so-called geometrical isomers. Geometrical isomers are compounds with the same structural formula but with different relative spatial positions of the ligands. For example, there are two possible geometric isomers for a trigonal bipyramidal molecule of the type AL4L’ where L and L’ are nonidentical ligands. One isomer has the L’ ligand in the equatorial position; the other isomer has the L’ in the axial position. Geometrical isomerism can also be demonstrated for the trigonal bipyramidal molecule (AL3L2‘) and the octahedral molecule (AL4L2‘).

Molecular models can be used to determine whether two compounds are geometrical isomers. If the two models have the same ligands attached to the central atom, but their models cannot be superimposed so that all of the ligands on each molecule match, then the two are geometrical isomers.

Pre-Laboratory Assignment

In order to successfully perform this modeling exercise, you must know how to draw Lewis electron dot structures for simple molecules. Before coming to class, calculate the number of valence electrons and a draw Lewis structure (disregarding geometry) for each of the nine compounds listed in Parts II and III of the experiment. Consult your textbook for help.

Note: These pages give only a brief summary of the concepts of molecular geometry and the VSEPR theory. For additional discussion of this topic, you should consult your textbook.

THE EXPERIMENT

 

You will use VSEPR theory to construct accurate three-dimensional models of a number of molecules. During the process, you should gain familiarity with VSEPR theory as well as an ability to visualize the three-dimensional structure of molecules. In the last part of the experiment, you will use the models to predict how many different isomers of a given compound possibly exist.

PROCEDURE
Part I

Construct the five basic geometrical shapes on which VSEPR theory is based: linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral.

  1. Select the proper ball, containing the proper number and placement of holes, from your molecular model set to represent each central atom.
  2. Insert rods into the holes to construct the figure. The rods represent electron pairs (either lone pairs or bonding pairs) around the central atom. Refer to the drawings in your manual.
  3. Sketch the molecular models you constructed.
  4. Use a protractor to measure the angles between all adjacent electron pairs (represented by the rods) in each geometrical model.
  5. For the trigonal bipyramidal shape, differentiate between the axial and equatorial positions and angles.
  6. For the octahedral shape, differentiate between the two sets of angles.
  7. Record the measurements in your notebook data sheet.
  8. When you have completed all five figures, have your instructor check them.
Part II

Construct models of the molecules listed below using VSEPR theory to predict the geometry.

(The central atom is written first in the molecular formula.)

BeH2 NH3 XeF4 OH2 PCl3F2 SF4Cl

For each case:

  1. Calculate the number of valence electrons in the molecule (if you did not do this before lab).
  2. Draw a Lewis electron dot structure for the molecule (if you did not do this before lab).

 

  1. Determine the values of x and y in the ALxNy notation. (Remember, y refers to the number of non-bonded electron pairs on the central atom only; do not include non-bonded pairs on the ligands.)
  2. State the geometric arrangement of electron pairs corresponding to (x + y).
  3. Build the molecular model using different colored balls to represent different atoms or ligands. Represent the single bonds by using a single connecting rod between the ligand and the central atom. Use a rod, but no ball to represent non-bonded electron pairs on the central atom.

1. Draw the molecule in its correct geometry using the figures in Table 1 as a guide. Explicitly write the symbol for each atom or non-bonding pair of electrons in your drawing.

Part III

To investigate geometrical isomerism, build molecular models of the two possible isomers of PCl4Br. Convince yourself that, even though the isomers contain the same atoms and bond types, the two cannot be superimposed and are different compounds, that is, geometrical isomers. Sketch these isomers in your notebook.

Next, draw Lewis electron dot structures and construct models of all possible isomers for each of the molecules or ions listed below. Use your models and sketches to decide how many isomers there may be for each case and record this on your notebook data sheet, along with the geometry and a sketch of each isomer. To decide whether your models represent isomers or identical molecules, remember that models of identical molecules are mutually superimposable, while models of isomers are not superimposable.

Although isomers may be drawn or models may be built for many molecules, all of the possible isomers may not exist. There are theories which explain why one isomer is more stable than another and thus why it may be the only one obtainable in a laboratory synthesis.

TeCl3Br SF4Cl2 XeF2

TABLE 2

________________________________________________________________________________________________________

(x + y) ALxNy Geometric Arrangement

x y of Electron Pairs

________________________________________________________________________________________________________

3 3 0 trigonal planar

  1. 2 1 trigonal planar
  2. 4 0 tetrahedral

4 3 1 tetrahedral

4

 

2

2

tetrahedral

5

 

5

0

trigonal bipyramid

5

 

4

1

trigonal bipyramid

5

 

3

2

trigonal bipyramid

5

 

2

3

trigonal bipyramid

6

 

6

0

octahedral

6

 

5

1

octahedral

6

 

4

2

octahedral

6

 

3

3

octahedral

6

 

2

4

octahedral

 

Part IV

Now that you have an understanding of VSEPR and the Lewis Octet Rule, apply the concepts and see how many isomers you can propose for C3H6O.

Even though you are not familiar with the classification of organic compounds, you should still be able to derive plausible structures.

 

Hint:

Ketone (acetone)

Ether

Aldehyde

Alcohol

 

Experiment 12 ESTIMATION OF AVOGADRO’S NUMBER

 

Estimation of

Avogadro’s Number

word image 955 word image 956 word image 957

Background

Avogadro’s number refers to the number of molecules in one mole of a substance. As we have learned, one mole is in grams equal to the molecular weight. We could easily determine Avogadro’s number if we knew the mass of a single molecule; but seeing as they are so small, that is nearly impossible. However, in this experiment, we have devised a way to measure the length of a molecule. By making an assumption about the molecule’s shape, we can then find the volume of a molecule, and then using the density, find its mass. Knowing the mass of a molecule and the molecular mass will lead us to Avogadro’s number.

word image 958 To understand this lab, we must first understand the difference between a hydrophilic and a hydrophobic molecule. Hydrophilic molecules have a functional group (a set of atoms in a molecule that give it certain reactivities) that tends to mix well or dissolve in water. On the other hand, a Hydrophobic molecule has a functional group that does not mix well or dissolve in water. Certain organic molecules have a unique property where they are hydrophobic on one of end of the molecule and hydrophobic at the other end and we will put this to use during this laboratory.

Procedure

word image 3228 word image 959

Figure

8

 

How to read the syringe

 

Start by calibrating a medical syringe, which will be used as a dropper in this experiment. Add as close to 1 milliliter of Oleic Acid to a 10mL graduated cylinder until the meniscus is tangent to the 1 or 2 ml line as shown below. Count the number of drops required to fill the cylinder to the next milliliter line. Repeat once or twice, recording the number of drops each time. It should be ~45 drops average.

In the hood, fill a small 50 mL beaker with the Lycopodium powder. Do not go over the 10 mL mark, as very little is needed. Take parafilm wax and cover the beaker opening, sealing the powder in. With a pencil or scissors, poke a few tiny holes in the parafilm wax to act as a salt/pepper shaker. This will make spreading the powder easier. Using a pan or cookie sheet, fill it until there is a small layer of water covering the bottom. This next step is crucial. Carefully and gently, sprinkle the Lycopodium powder over the pan. If you can clearly see through the layer, it is too thin; or if you cannot see the bottom, it is too thick. If the layer is too thick, the film will not spread. If it is uneven, it will not be circular.

Pull about 1 mL of Oleic Acid into the syringe. Push a few drops out, and then let a drop of Oleic Acid fall into the center of the Lycopodium film. If a circular hole is not produced, the procedure should be repeated. Record the diameter of the film in several directions with a ruler in centimeters and take the average.

Data Collection

It is important to keep one’s data in order, so a proper data table should be set up and should look like the following example.

Volume of Oleic Acid in solution __________ (µL/mL)

Syringe calibration (drops) __________ Average __________

__________

__________

Film Diameter (cm) __________ Average ___________

 

__________

__________

__________

Calculations

  1. Calculate the volume of oleic acid in the film from one drop of the solution. The stockroom prepared the solution so that 0.125 mL of oleic acid is dissolved in a water soluble solvent. The film then is totally from oleic acid.

word image 960

  1. Determine the area of oleic acid on the water. Remember the film will contain only oleic acid since the solvent dissolves in water.

word image 961

  1. Calculate the height of the film. Remember, 1 mL = 1 cm3

word image 962

  1. Assuming the molecule is a rectangle with a height 10 times its width, calculate the volume of one molecule of oleic acid. Here we are also assuming the film is only one molecule thick.

word image 963

 

  1. Calculate Avogadro’s Number, word image 964

word image 965

 

 

 

Experiment 13 GRAPHING

 

 

 

word image 966 word image 967 word image 968 word image 969

 

Many physical and chemical properties are linearly related to each other. As the numerical value for one property changes, there is a proportional change in the other property’s value. A mathematical equation which expresses the linear relationship is

y = mx + b

where x and y are the values of the two properties of interest. The slope, m, and the intercept, b, are constants.

word image 970 It is often helpful or necessary to visualize the linear relationship by graphing the values for x and y. Our example will highlight the relationship between the Celsius and Fahrenheit temperature scales. The data which is plotted on the graph is shown to the below. Refer to the graph on the following page while you read about good graphing techniques below.

Degrees

 

Celsius

Fahrenheit

5

38

22

67.3

32

95.5

51

120

74

165.2

100

205.6

Graphing data:

  • Determine which properties will be plotted on the horizontal (x) axis and on the vertical (y) axis. Often you will be told which property is x and which is y. Label the x– and y-axes on the graph and give a title to the graph.
  • Decide on the scale of each axis. Generally, the completed graph should fill the page. The range of the x axis should have as its minimum value a number which is equal to or somewhat smaller than the smallest x value to be plotted; the maximum x-axis value should be somewhat larger than the largest x value to be plotted. Likewise for the y-axis and its values to be plotted. It is not necessary for the minimum x and y values to be zero, although either or both may be zero.
  • Divide the x– and y-axes into units according to the range of values to be plotted and the number of divisions on the graph paper.
  • Plot the data. In a hand-drawn graph, circle the data points so they are emphasized against the grid background.
  • Using a transparent straight-edge, draw the best straight line. Many times, the best straight line will not go through every point on the graph. If you have determined the data yourself, you may have more confidence in some values than others; these points should lie on the line. Otherwise, try to put as many points on the line as you can, and have the remainder of the points lie nearly equally both above and below the line. Check for any values which are so far off a reasonably straight line that you suspect they are outliers and represent inaccurate data.

Determining the constants, m and b:

  • The intercept, b, is that point on the graph for which the x value equals zero, that is, where the x = 0 point intercepts the y-axis. In some cases, this intercept has physical significance and in others it does not. To find b from the graph, with your straight-edge draw an extension of the best straight line across the region where x = 0. Then carefully read the corresponding value from the y-axis.
  • Calculate the slope using two points from the best straight line. Try to use two points which are widely spaced on the graph. You can use plotted data points only if these fall on the best straight line; otherwise, first determine the x, y point values anywhere on the best straight line you drew. The calculation of the slope is:

y2 y1 m

x2 x1

To clearly specify which data points you used to calculate the slope, draw dashed horizontal and vertical lines from the points to their intersections with the x and y axes, respectively. Use several different points to verify the slope calculation is essentially the same.

Using a calculator:

If you have a calculator or computer capable of linear regression analysis, feel free to use it.

You should always plot your data first. Using your data, calculate the regression constants m and b. If b has physical significance and if the intercept would be a point on the graph, plot the calculated value of b and label it as such. Using the “trend” function on the electronic device, calculate a value for two or more of the dependent variables (y-axis) using your x-value data. Plot these on your graph and label them as such. Now draw the best straight line through the calculated values on the graph. With some practice, you can get a good idea of how to draw a best straight line without using the electronic device.

Problems (show work):

Solve the problems below on a separate page. Show all of your calculations. Follow the guidelines above as you mark points on the attached graph.

  1. On the Fahrenheit-Celsius graph, extrapolate the line to find b. What is the value? Circle it.
  2. Using two points on the best straight line, calculate m.
  3. Using the calculated values of b and m, write the equation which linearly relates all oF and oC temperature values, i.e. oF = m(oC) + b.
  4. What is the Fahrenheit temperature which corresponds to 62oC? Use both the graph and your answer to question #3 above to determine the value.
  5. Absolute zero (0 K) is −273 oC. What is the value in Fahrenheit?

 

Degrees Fahrenheit vs. Degrees Celsius

0

50

100

150

200

250

0

20

40

60

80

100

120

Degrees Fahrenheit

 

Degrees Celsius

 

-114-

 

Experiment 14 THIN-LAYER CHROMATOGRAPHY

 

 

 

 

 

word image 971 word image 972 word image 973 word image 974 word image 975

Chromatography is one of the most powerful modern methods for separating, isolating and/or identifying components of a mixture. The term “chromatography” is derived from the Greek word, chromatos, for color. The technique was so named because the original discovery occurred during the separation of mixtures of natural plant pigments. Although the pigment mixture appeared green to the eye (because of the preponderance of green chlorophylls), the separate chlorophyll and carotenoid pigments, which are green, yellow and orange, could be easily seen as distinct bands of color.

word image 976 Current day chromatographic methods employed do not necessarily depend on natural color because the components can be detected in various other ways. These chromatographic techniques are widely used in a variety of scientific areas — criminology and forensics, environmental and biology labs, and in medical research and analysis.

Chromatographic separation of the components of a mixutre depends on their differential distribution between two phases: one is a stationary phase and one is a mobile phase. The different components in the sample mixture will move at different rates as they pass through the stationary phase because each component will have a different physical interaction with the stationary phase. Therefore, some components will move faster or slower than others, and eventually the components in the mixture will be separated from each other.

The phases may be almost any combination of gas, liquid, or solid. Most practical chromatographic techniques involve a solid-liquid, liquid-liquid, or gas-liquid phase pair. However, all of the techniques have one common principle: the sample mixture is dissovled, usually in a liquid or a gaseous solvent (the mobile phase), and is passed through an adsorbent (the stationary phase).

In thin-layer chromatogaphy (TLC), a solid adsorbent is thinly coated on a glass or plastic rectangular plate. Silica gel or alumina coatings are the stationary phases used most often in TLC. The mixture which is to be chromatographed (separated into its constituent components) is applied as a concentrated liquid solution near one edge of the stationary phase. The edge of the TLC plate is immersed in a solvent and this moving liquid phase (mobile phase) engulfs the sample spot, thereby carrying it through the stationary phase. The sample components are partitioned between the stationary and moving (mobile) phases; separation results from the differences in the rates at which the individual components advance upwards on the plate. As separation of the mixture takes place, the components will be distributed in space over the stationary phase behind the mobile phase (moving liquid front). See Figure 1.

The ratio of distances traveled by a sample component and the moving phase is called the Rf value, which is a physical property of that substance under the specified conditions. Thus, TLC can be used to separate and identify (by comparison with known Rf values of a standard) the components in a mixture.

word image 3229

 

word image 977

If the sample components absorb light in the visible wavelength region, their presence can be detected by observing the colored spots on the developed chromatogram. Fluorescing materials are usually incorporated in the TLC adsorbent coating and, for substances which do not absorb in the visible region, the solute’s location will appear as a dark spot on a fluorescent background when viewed under ultraviolet light. The surface of the thin layer can be scanned with a radiation detector when radioactive tracers are present in the sample components. In preparative large-scale TLC, the desired zones can be removed from the plate by scraping off the part of adsorbent coating which contains the separated solute.

THE EXPERIMENT

Certified red, yellow, and blue dyes are used to color many of the foods we eat and cosmetics we use. The molecular structures of some of the Food, Drug, and Cosmetics dyes are shown in Figure 2. Four combinations of F,D,& C dyes can be bought as food coloring in grocery stores. In the laboratory we have available samples of pure F,D,& C dyes that may be used as Rf comparison standards for analyzing other food or cosmetic products that contain these dyes.

In the experiment today, the coloring agents in various Kool-Aid flavors will be separated by chromatography and their Rf values calculated to ascertain which F, D, & C dyes are present. If you measure the physical property of the Rf value for pure dyes and then compare them with the Rf value for an unknown dye, you may be able to tentatively identify the unknown on the basis of this physical property. The wavelength of maximum absorption (max) is another identifying physical property of the pure dye.

PROCEDURE

CAUTION: There should be no flames in the laboratory during this lab because of the flammable solvents in use.

For the chromatography of pure F,D,&C dyes and Kool-Aid, use Silica TLC chromatography plates, approximately 20 x 5 cm, and a developing solvent that consists of a 1: 1 : 1 mixture of ethanol, 1-butanol, and 2-M ammonium hydroxide.

  1. First, pour some of the solvent mixture into the development chamber (a glass container with a tightly fitting cover) to just cover the bottom of the jar, close the lid, and let the atmosphere become saturated with the solvent vapor while you prepare your chromatography plate. To speed up the vaporization process, place a couple of strips of filter paper in the solvent about half-way up the sides of the jar.
  2. When preparing your chromatogram, handle the TLC plate with your fingertips at the edges only. Obtain two TLC plates and place them on a paper towel on a clean desk to avoid contamination.
  3. On each plate, draw a light pencil line about 1.5 to 2 cm above the bottom edge of the plate where the sample is to be applied. (The position of the line is determined by the amount of solvent used in the developing chamber.
    1. The spots must be above the level of the solvent in the development chamber.) Be careful not to scrape the solid coating.
    2. Lightly draw 4 x’s evenly spaced on the line. The samples will be applied on the x’s as concentrated spots, using toothpicks as described below. Identifying labels are also made (in pencil) below the line.
  4. Spotting a Sample—To spot a Kool-Aid sample, immerse the small end of the toothpick into the sample solution and then touch the tip of the toothpick to the pencil “x” on the paper to form a tiny spot. If it is necessary to spot more sample at the same position, allow the first sample spot to dry; otherwise, the spot will be too large and the separation will be poor.
    1. Spot the same flavor of Kool-Aid on each of the four positions on your plate.
    2. Your instructor will spot a TLC plate with the four pure F ,D, & C dyes. This will be the set of standard dyes that you compare with the Kool-Aid dyes.
  5. Developing the ChromatogramAfter spotting, carefully place the chromatography plate inside the development chamber with the spots at the bottom, but above the solvent surface.
    1. Be sure the plates do not touch the sides of the chamber or each other and that they are as upright as possible.
    2. Cover the chamber and allow the solvent to move up the plate. Do not move the chamber.
    3. Do not let the solvent ascend all the way to the top edge of the plate. The development can be stopped when (a) the solvent comes within 1 cm of the top of the plate; (b) the solvent front stops ascending by itself; or (c) no more color separation appears to take place. The development usually takes less than an hour.
    4. Remove the wet plate carefully, quickly mark the solvent front with a light pencil line and let the plate air dry.
    5. Measure the distance the solvent front traveled from the pencilled line at the bottom of the plate to the line you marked at the end of development.
    6. Determine as closely as possible the center of each component spot, measure its distance from the starting point, and calculate the Rf values.

Identify the component dyes (by their Rf values) in the Kool-Aid samples by comparing them with the standard dyes. Explain your reasoning. _________________________________________________________________________________________________________

THIN-LAYER CHROMATOGRAPHY

_________________________________________________________________________________________________________

Pure F, D, & C Dyes

Distance traveled by solvent front ________ cm

 

Visible color Distance traveled Rf

 

 

 

 

 

 

 

_________________________________________________________________________________________________________

 

Kool-Aid Flavor

 

Distance traveled by solvent front ________ cm

 

Separated color Distance traveled Rfmax

 

_____________________________________________________________________________

Identity of food dyes in Kool-Aid flavor based on Rf.

(In your notebook, record a table of Absorbance vs. Wavelength (400-700 nm). Attach graph of absorption spectra.)

 

Experiment 15 CHEMICAL INTERACTIONS

 

 

 

 

word image 978 word image 979 word image 980 word image 981 word image 982 word image 983

INTRODUCTION
Salts

Materials commonly called salts are those that are composed of a metal ion and a non-metal ion

(or, non-metallic polyatomic ion). Most common salts dissociate in water to form positive ions (cations) and negative ions (anions):

NaCl(s) + H2O (l) → Na+(aq) + Cl(aq)

The resulting ions are said to be aquated or aqueous, which simply means closely associated with water molecules in an aqueous media. This behavior is symbolized by “(aq)”, found after the ions of interest.

word image 984 Some water-soluble compounds are not salts. Common table sugar (sucrose) is an often-cited example. Sucrose is a molecular compound, which dissolves in water, but does not dissociate into ions. To summarize these differences, study the examples below, which include water-soluble examples of ionic and molecular compounds.

NaI (aq):

 

Na+(aq) and I(aq)

Cu(NO3)2 (aq):

 

Cu2+ (aq) and 2NO3 (aq)

C12H22O11 (s):

 

No dissociation, simply C12H22O11(aq)

(NH4)2SO4 (aq):

 

2NH4+(aq) and SO42– (aq)

 

Once dissociated, the positive and negative ions act independently from each other. This allows their aqueous solution behavior to also be studied independently. Any aqueous solution containing the silver(I) ion will react with any other solution containing the chloride ion (Cl) to produce the compound AgCl(s), which does not dissolve in water. The silver chloride thus formed is insoluble in water, and forms a solid precipitate of AgCl(s).

Ag+(aq) + Cl(aq) → AgCl(s), white solid precipitate

Acids and Bases

Acids can be defined as compounds that will increase the aqueous concentration of hydronium ion (H3O+, or more simply H+) as they dissolve in water. Common acids are nitric acid (HNO3), hydrochloric acid (HCl), and sulfuric acid (H2SO4). In general, this increase in aqueous hydronium ion concentration can be represented by a relatively simple dissociation equation:

H2SO4(aq) → 2H+(aq) + SO42– (aq)

For the most part, an acid is most easily recognized by a hydrogen atom in a formula that either begins a formula, or seems to be set apart from the other hydrogen atoms in a formula. This is meant to convey the relative ease with which that hydrogen atom is lost to the aqueous media as hydronium:

HNO3(aq) =

nitric acid

HC2H3O2(aq) =

acetic acid

H3PO4 (aq) =

phosphoric acid

In any case, an acid increases the hydronium ion concentration. When an acid dissolves in water, a convenient measurement of this increase is the pH scale. Acidic solutions have a pH value less than

7.

Bases, by contrast, increase the concentration of hydroxide ion (OH) as they dissolve in water. Any salt containing the hydroxide ion (NaOH, KOH, etc.) in its formula is, by definition, a base. The following ions are produced when sodium hydroxide dissolves in water:

NaOH(s) → Na+(aq) + OH(aq)

Other bases include aqueous ammonia (NH3) and aqueous sodium sulfide (Na2S) for reasons not so obvious. The reason they are classified as bases has to do with the behavior of the ions these compounds produce, after dissolving in water. S2 (aq) + H O(l) → HS (aq) + OH (aq)

2

NH3(aq) + H2O (l) → NH4+(aq) + OH (aq)

PERFORMING THE EXPERIMENTS, AND INTERPRETING YOUR RESULTS

The design and aim of this experiment is to identify unknown aqueous solutions. To do this, you must carefully follow directions, meticulously record what you see when two given solutions are mixed, and be able to deduce what might happen if solutions you do NOT actually mix are mixed. There is a general series of things you can watch for, to bring a degree of organization to your results, and help you deduce the identity of your unknowns.

Color and Odor

You should carefully examine the unknowns you are given, and if a perceptible color or odor are present, describe it to the best of your ability. Specifically:

•Solutions of NH3 and Na2S have characteristic odors. You should carefully direct the vapors from the CLOSED TEST REAGENT BOTTLE toward your nose. Describe the odor in your own words.

•Aqueous solutions of compounds of the transition metals have identifying colors. Copper compounds are usually blue, compounds of nickel are usually green, compounds of iron are orange, yellow, or purple.

•Any precipitates (solids) that form from mixing aqueous solutions should be characterized by color. A chart in this procedure gives you some common insoluble materials and their colors. There may be some trends in that information that could be useful to you.

Acidity and Basicity

The four primary classifications for the strength of acids and bases are listed below in Table 1. Table 1. Acid and Base Strength

Strength

pH range

Characteristics

Strong Acids

<3

Form solutions of the common aqueous acids.

Weak

Acids

3 to 7

Salts (usually of the transition metals) dissolved in water form solutions that test weakly acidic to pH paper.

Weak Bases

7 to 11

Salts of the main group metals (IA, IIA or IIIA) dissolved in water will usually form solutions that test weakly basic to pH paper. NH3 and Na2S form weakly basick solutions.

Strong Bases

>11

Form when hydroxide salts are dissolved in water.

 

The Evolution of Gases

A few of the possible unknown anions in the experiment will produce gases when reacted with aqueous acids. You may or may not see bubbles. The three possibilities include bicarbonate (HCO3), carbonate (CO32-) and sulfide (S2-). The first two produce odorless CO2(g), while sulfide produces a small amount of H2S(g), which smells like rotten eggs.

Aqueous Solubilities

Insoluble precipitates will be formed when certain combinations of aqueous solutions are mixed. It is useful to be aware of what products two solutions will make when mixed, and whether or not any of those products will be insoluble in water. Being aware of these solubility rules can be of a great help in identifying unknowns.

1). All common salts of the nitrate ion (NO3) are soluble in water.

2). All common salts of the Na+, K+, and NH4+ cations are soluble in water.

3). Most common salts of the Cl, Br, and I anions are soluble in water EXCEPT when the accompanying cation is Ag+, Pb2+, and Hg22+.

4). All common salts of the sulfate ion (SO42) are soluble in water EXCEPT when the accompanying cation is Ba2+, Sr2+, and Pb2+.

5). Salts of the S2- ion are insoluble in water EXCEPT when the accompanying cation is Na+, K+, NH4+, Ba2+, and Ca2+.

6). Most of the transition metal cations (Ag+, Cu2+, Ni2+, etc.) form complexes with aqueous ammonia. Sometimes, ammonia (NH3) is used to “dissolve” otherwise insoluble precipitates of these metal cations.

7). Most salts of the hydroxide ion (OH) are insoluble. The exceptions are as follow: the alkali metal (Group I) ions Na+, K+, Rb+, Cs+; the ammonium ion NH4+; and the alkaline earth (Group II) ions Ca2+, Sr2+, and Ba2+.

Table 2. Properties of Some Common Precipitates

 

Category

Compound

Formula

Precipitate Color

Solubility in water

Carbonates

CO32

Silver carbonate

Ag2CO3

Yellow

0.032 g/L (25 °C)

 

 

Barium carbonate

BaCO3

White

24 mg/L (20 °C)

 

Copper(II) carbonate

CuCO3

Pale blue

Insoluble

 

Calcium carbonate

CaCO3

White

24 mg/L (20 °C)

 

Chlorides,

Cl

Silver chloride

AgCl

White

520 μg/100 g at 50

°C

 

Copper(I) chloride

CuCl

White

0.0062 g/100 mL

(at 20°C)

 

Hydroxides,

OH

Silver hydroxide

AgOH

Grey-brown

Insoluble

 

Iron(III) hydroxide

Fe(OH)3

Rust red

Insoluble

 

Copper(II) hydroxide

Cu(OH)2

Pale blue

Insoluble

 

Nickel(II) hydroxide

Ni(OH)2

Pale green

0.013 g/100 mL

 

Zinc(II) hydroxide

Zn(OH)2

White

 

 

Iodides,

I

Silver iodide

AgI

Yellow

3×10−7g/100mL

(20 °C)

 

Copper(I) iodide

CuI

Brown

0.0042 g/100 mL

 

Sulfides,

S2

Silver sulfide

Ag2S

Black

6.21×10−15 g/L

(20 °C)

 

 

Iron(III) sulfide

Fe2S3

Black

 

Nickel sulfide

NiS

Black

Insoluble

 

Copper(II) sulfide

CuS

Black

0.000033 g/100 mL (18 °C)

 

Sulfates,

SO42

Barium sulfate

BaSO4

White

0.0002448 g/100 mL (20 °C)

0.000285 g/100 mL (30 °C)

 

EXPERIMENTAL PROCEDURE

You will be assigned 5 unknowns to test, and then to identify. You must read the test reagent bottles carefully, to ensure you are doing the correct tests, and not doing any unnecessary tests. You will be assigned 1 unknown from each of the following categories:

a). A strong acid — HCl, HNO3, or H2SO4

b). A strong base — KOH, NH3, or Na2S

c). A salt of the chloride ion — NaCl, CuCl2, or BaCl2

d). A salt of the nitrate ion — AgNO3, Cu(NO3)2, or Fe(NO3)3

e). A salt of the sodium ion — NaI, Na2SO4, or Na2CO3

 

Include in your lab reports a completely recopied chart (see sample on following page) describing the results of your tests as you interpret them. Also be sure and identify your unknowns by number and chemical name.

PERFORMING EACH QUALITATIVE TEST

All solutions necessary to perform this experiment are in dropper bottles. You will not need much, so use the bottles where you find them and do not take the screwcaps off. When you are done, make sure that the dropper cap is replaced and that the solutions are in place where you found them. Your fellow lab students will need to know where to find these materials.

For every test: Add approximately 5 drops of an assigned unknown from the dropper bottle to a clean, small test tube. Add approximately 10 drops of the test solution from the dropper bottle (watch those LABELS!!) to the test tube containing the unknown, one drop at a time. Describe carefully what you see happen, if anything, as soon as possible after performing each test. RECORD YOUR OBSERVATIONS IN THE CHART, taking care to note the colors observed, if a precipitate formed, if a gas or odor was given off as a result of the reaction.

Keep your test tubes very clean, and DO NOT touch the tip of any dropper bottle to any other test or unknown solution. This will cause contamination, and will drastically upset your coworkers, the instructors, and the entire lab preparation staff. Points will be deducted ON THE SPOT for anyone being careless in the handling of these solutions.

All tests will end up in the same waste bottle. Any clinging solids should also be rinsed into this waste bottle. Empty as much spent test materials into the waste bottle as possible, before cleaning your test tubes. This minimizes the discharge of anything harmful to the environment into the waste water system.

There are ONLY 8 aqueous test solutions. They are:

(1) KOH, (2) AgNO3, (3) BaCl2, (4) NaCl, (5) NH3, (6) Na2S, (7) Na2SO4, and (8) H2SO4. All test solutions are of sufficient concentration to minimize “false positive” tests.

CAUTION

 

If you have trouble telling the difference between NH3 and KOH (think separately about WHY it might be hard to tell between these two substances) using this general procedure:

  • Use 5 drops of AgNO3 plus 10 drops of NH3 test solution, note result.
  • Use 5 drops of AgNO3 plus 10 drops of KOH test solution, note result.
  • Use 5 drops of AgNO3 plus 10 drops of your unknown solution, compare this to the two above results.

The only potential “false positive” is a mixture of Ba2+ and S2-. It gives a cloudy whiteness, even though it should not. It comes about from tests that are not definitive for either ion anyway.

 

Summarize your results in a table like this one IN YOUR LAB NOTEBOOK

Unknown #

 

 

 

 

 

Color

 

 

 

 

 

 

pH

 

 

 

 

 

 

Odor

 

 

 

 

 

 

KOH

 

 

 

 

 

 

AgNO3

 

 

 

 

 

 

BaCl2

 

 

 

 

 

 

NaCl

 

 

 

 

 

 

NH3

 

 

 

 

 

 

Na2S

 

 

 

 

 

 

Na2SO4

 

 

 

 

 

 

H2SO4

 

 

 

 

 

 

 

Experiment 16 ABSORPTION SPECTROSCOPY

 

 

 

word image 985 word image 986 word image 987 word image 988

Purpose:

The purpose of this experiment is to learn the use of the Spectronic 20 spectrophotometer to investigate the absorption of visible light by an aqueous solution of a transition metal ion (Cu2+, Co2+, or Ni2+) by determining the wavelength of maximum absorption and demonstrating the relationship between absorbance at this wavelength and solution concentration as defined by the Beer-Lambert Law.

Background:

word image 989 Scientists spend much time studying the interactions between matter and energy. Because there are several forms of energy and there is a seemingly inexhaustible supply of matter, these studies have provided many measurements and much knowledge about the universe and its behavior. A study of the interaction of electromagnetic energy with matter provides valuable and important information about molecular structure and properties. In fact, some of the most convincing evidence about atomic and molecular structure has been obtained in this way.

Light is electromagnetic radiant energy. Depending upon the point of view, we can think of light as having either wave properties or particle properties. When we refer to photons, we are speaking of the particle-like aspects of light. According to the wave theory, electromagnetic radiation is composed of both electric and magnetic fields whose waves vibrate in mutually perpendicular planes. It is the electric field of light interacting with the electrons of matter which produces observable properties of a substance.

Two fundamental characteristics are associated with any wave — the wavelength and the frequency. The length,  (Greek letter, lambda), is measured from the crest of a wave to the adjacent crest (see Figure 1). Wavelengths range from very small (measured in nanometers, 1×10-9 m) to quite large. The frequency is the number of waves passing a point in unit time and is designated  (Greek letter, nu). Frequency can be expressed in cycles (or vibrations) per second. Wavelength and frequency are inversely proportional to each other: the shorter the wavelength, the more waves pass a point in a given time and thus the higher the frequency. The longer the wavelength, the lower the frequency. Wavelength and frequency are related by the speed of light (c = 3 x 1010 cm/sec) such that that we have c = .

 

 

 

Figure 1

 

word image 3230

Another important relationship is that between these quantities and the energy of light: E = h, or the equivalent E = hc/, where h is Planck’s constant (6.626 x 10-34 J.sec). The greater the energy of the radiation, the shorter its wavelength and the higher its frequency; the smaller the energy of radiation, the longer its wavelength and the lower its frequency.

Light from the sun is composed of a continuum of energies and thus of a continuum of wavelengths and frequencies. Although much of the electromagnetic radiation continuum is invisible to us, we perceive a portion of it as visible light. The wavelengths of visible light extend from about 800 nm (8.0 x 10-5 cm) to about 400 nm (4.0 x 10-5 cm). The continuum, or electromagnetic spectrum, extends from the visible region outwards in both directions, as shown in Figure 2.

For convenience, the electromagnetic spectrum is arbitrarily divided into regions called bands. Some commonly used regions are the ultraviolet (UV), visible, infrared, microwave, and radio and TV waves. It is against the high energy Ultraviolet (UV) radiation that sunscreen lotions are formulated to protect our skin in the outdoors. Our skin perceives the low energy infrared (IR) radiation only as heat. The lamps placed over food in cafeterias emit most of their energy in the red part of the Infrared region, keeping the food warm.

If little or none of the visible light striking our eyes is not absorbed prior to striking the eye, the color appears white. If this band of visible light is separated into narrow bands of wavelengths by a prism, we perceive the component colors. Isaac Newton divided the visible spectrum into seven color bands ranging in sequence starting with the shortest wavelengths (left to right in Figure 2 : violet, indigo, blue, green, yellow, orange, red.

When light impinges on a substance, one or more combination of things can happen to the light. The light can be scattered, reflected, transmitted, or absorbed by the substance. The absorbed light energy causes such changes as atomic and molecular rotation, vibrations, and electron transitions to higher energy levels. As a result of this absorption, our eyes or specially designed instruments may sense phenomena such as heat, fluorescence, phosphorescence, or color. Modern instrumentation can record these phenomena to a very high degree of precision.

 

Figure 2 Electromagnetic Spectrum

 

 

 

 

 

 

 

 

word image 3231 word image 990

 

 

 

word image 991

 

 

word image 992

 

 

word image 993

 

 

word image 994

 

 

word image 995

 

 

word image 996

 

 

 

word image 997

 

 

word image 998

 

 

 

Let’s consider what happens when light in the visible range interacts with matter. If a substance absorbs all wavelengths in the visible range, none of the light is reflected back to our eyes and the substance appears black; if the substance absorbs none of the visible range wavelengths, it appears white (all light reflected) or colorless (all light transmitted). Colorless substances usually absorb in the UV and IR regions of the spectrum, on either side of the visible range.

If a substance absorbs light principally in one wavelength range (generally, a number of wavelengths on both sides of the principal absorption wavelength are also absorbed, so a broad absorption band results), the color you see will be a mixture of all the wavelengths which are not absorbed. For example, the indigo dye in blue jeans has its maximum absorbance in the 500-650 nm range. Because this absorbance is in the red-to-green region, the wavelengths which are not absorbed are observed to be blue-violet. Chlorophyll reflects primarily yellow and green wavelengths and absorbs mainly in the blue-violet and red regions (which is why plant “grow-lights” are designed to emit more radiation in these color bands). The mix of reflected wavelengths appears green — the color your eye “sees”. The exact color will depend on the relative intensities of the non-absorbed wavelengths as well as on a somewhat subjective interpretation by the observer. (See Figure 3.) Table 1 is a summary of the relationship between the wavelengths of colors observed and colors absorbed.

 

 

Table 1

 

 

Wavelength Absorbed (nm)

Color Absorbed

Color Observed

410

violet

yellow-green

430

blue-violet

yellow

480

blue

orange

500

blue-green

red

530

green

purple

560

yellow-green

violet

580

yellow

blue-violet

610

orange

blue

680

red

blue-green

720

red-purple

green

 

 

A more fundamental grouping of colors than Newton’s is the RGB (red, green, blue) group. If any of these three colors is absorbed from white light, the complementary color which is observed is what remains. Complementary colors are pairs of colors which, when combined in the right proportions, produce white or black. When placed next to each other, they create the strongest contrast and reinforce each other. They are widely used in art and design and especially in video monitors, such as television screens. In painting, which uses subtractive colors, the traditional primary–secondary complementary color pairs, described since at least the early 18th century, were red–green, yellow– violet, and blue–orange.

The more accurate RGB (red, green, blue) color model, invented in the 19th century and fully developed in the 20th century, uses additive color combinations of red, green, and blue light against a black background to make the colors seen on video screens. In the RGB color model, the light of two complementary colors, such as red and cyan, combined at full intensity, will make white light, since two complementary colors contain light with the full range of the spectrum. If any of these three colors is absorbed from white light, the complementary color is observed. The proper combination and intensities of the three primary colors create every conceivable color and shade. Table 2 shows a summary of the primary-secondary complimentary colors used in the RGB model.

 

Figure 3

 

word image 999 word image 1000

400 500 600 700 400 500 600 700

Wavelength (nm) Wavelength (nm)

 

word image 1001 Absorption spectrum of chlorophyll a Absorption spectrum of chlorophyll b The absorption spectra on this page are for three of the colored constituents of spinach leaves. Approximately nine pigments were isolated by thin layer chromatography (see a later experiment) and most of them were identified from literature reports of their max. However, cryptoxanthin and zeaxanthin have nearly identical absorption spectra and could not be positively identified. Their structures are:

 

 

CH3 CH3 H3

H3C

C

H

3

 

Absorption spectrum of either cryptoxanthin or zeaxanthin

 

.16.1.1 Table 2

Complimentary Colors and Wavelengths of Maximum Absorption (nm)

Primary

 

Secondary

 

Red 690

Cyan (Green-Blue)

488

Green 520

Magenta (Red-Blue)

(non-spectral)

Blue 480

Yellow (Red-Green)

580

 

 

Spectroscopy is a basic analytical technique and research tool that utilizes the interaction between matter and electromagnetic energy. A spectrometer is an instrument that separates electromagnetic radiation according to wavelengths, passes these separated wavelength bands through a sample, and detects the intensity of the transmitted light.

The absorption spectrum is characterized by the wavelength of maximum absorption (λmax) at which the absorbance is the greatest (see Figure 4 below. The value of λmax is important for several reasons. It is used in order to obtain the highest sensitivity and to minimize deviations from Beer’s Law. It is characteristic of each unique compound providing information on the electronic structure of the compound.

Figure 4

 

 

 

 

 

 

 

 

 

word image 3232 word image 1002

Wavelength

 

word image 1003

Absorbance

 

word image 1004

 

word image 1005

max

 

All spectrometers have the following fundamental parts: a source or radiant energy, a prism or grating to isolate radiant energy to narrow wavelength regions, a sample holder, and a detector for measuring light intensity. Sophisticated instruments include automatic recorders, digital readouts, computer interfaces, and arrays of detectors that allow the user to analyze a wider range of wavelengths.

The graphical plot of the intensity of absorption of radiation versus the wavelength absorbed is called a substance’s absorption spectrum. Depending upon the spectral region of the incident measure both the percent of transmittance (%T) and absorbance (A) of light passing through the sample. The readings given by the meter or display can be either read in absorbance (A) or percent transmittance (%T).

The percent transmittance is the ratio of the amount of light transmitted (passed through), I, to the amount that initially strikes the sample, Io:

 

%T = (I/Io) • 100

Absorbance is the amount of light absorbed expressed in logarithmic terms and is defined as the negative logarithim (base 10) of the Transmittance:

Absorbance= A = -log T = log (Io/I)

Which of these two is used depends on the operator and the experiment.

The amount of light absorbed by a sample is related to the concentration of the absorbing species. In comparing two colored solutions, the darker colored solution appears darker because it absorbs more of the light striking it, and it absorbs more of the light because there are more lightabsorbing molecules present. Thus, the darker solution is also more concentrated. In both common sense and spectroscopy, the intensity of color is a measure of the amount of a material in solution. And, as is the case between density and concentration, the concentration vs. absorbance relationship is linear.

 

 

C

o

n

c

e

n

t

r

a

t

i

o

n

 

(

M

)

6

5

4

3

2

1

0

#

0

.

0

2

.

0

1

.

0

A

b

s

o

r

b

a

n

c

e

Problems (show work)

  1. For each of the wavelengths in the first column in the Table 2 above, calculate the frequency in Hertz (cycles/sec) and the energy in kJ.
  2. Looking at the range of colors in Figure 2, explain why magenta does not have one wavelength of maximum absorption.
  3. In Figure 2, what are the wavelengths of maximum absorption for chlorophylls a and b? (Do not choose either 400 nm or 700 nm as max.)

Answers to Selected Problems

#1. 690 nm = 4.35 • 1014 sec−1 = 174 kJ/mol

 

THE EXPERIMENT

In this experiment you will use a spectrometer to determine the wavelengths of visible light absorbed by substances in solution and plot an absorption spectrum. After identifying the wavelength of maximum absorption and preparing a series of standard solutions you will plot a calibration line to relate the intensity of absorbance and the concentration of a solution. Then, as an unknown, you will determine the concentration of the substance in solution.

The Spectronic 20 Spectrometer

The most common spectrometer for general chemistry laboratories is the Spectronic 20. Your instructor will demonstrate how to use the Spectronic 20 (Spec 20).

Cuvettes

Cuvettes (or cells) are sample tubes of precise dimensions, made from special glass to ensure uniform transmittance of light. You will need at least two — one for the solution and the other for a reference blank. The composition of glass may vary from one manufacturer to another or from batch to batch. Check to see that both cuvettes are the same brand, such as Pyrex or Kimax.

Wash and rinse cuvettes with distilled water. Rinse a cuvette with the sample solution and then add fresh sample solution to measure the absorbance. the outside of the cuvette with a tissue and handle the cuvette only on its top sides.

 

Calibrating the Instrument

  1. Let the instrument warm up for at least 20 minutes.
  2. Set the desired wavelength with the wavelength selector knob.
  3. With no cuvette in the instrument, set the readout display to 0% transmittance using the left hand knob.
  4. Place a cuvette approximately half filled with the reference blank (distilled water) in the instrument sample holder. Align the front of the cuvette with the mark on the front of the sample holder. The top of the holder must be shut except when you are loading or unloading a sample to prevent stray light from entering the instrument.
  5. Set the readout display to 100% transmittance (zero absorbance) using the right hand knob.
  6. Remove the reference blank. The display should read zero percent transmittance. If it does not, repeat steps 3 through 6 or obtain the instructor’s help.

Taking an Absorbance reading

Place a cuvette containing your sample in the instrument, close the lid, and record the absorbance and wavelength.

The instrument must be calibrated every time you change the wavelength. This is not necessary if you merely change the sample but not the wavelength. A substance’s “absorptivity factor” is a constant only as long as the wavelength of light is constant.

PROCEDURE

Before you begin, check your cuvettes. The clean cuvette that will hold your sample when filled with water should read 100% transmittance (or zero absorbance) against the cuvette with which you calibrate the instrument (water is used as the blank). Use the same cuvette for all parts of the experiment.

Part 1 The Absorption Spectrum

  1. Obtain a 15-mL sample of known concentration of absorbing species. Record its name, molecular formula, and concentration (M).
  2. Set the wavelength to 375 nm and calibrate the spectrometer according to the directions above. Read the absorbance of your sample at that wavelength. Remove the cuvette from the spectrometer.
  3. Increase the wavelength by 25 nm, calibrate the spectrometer, and take an absorbance reading of your sample. To obtain the absorption spectrum, continue to take absorption readings of your sample at 25 nm intervals up to 700 nm.

Remember to calibrate the spectrometer every time you change the wavelength. Save your solution to use in Part 2.

4 Identify the maximum absorbance and its corresponding wavelength (max) from your data. (Do not choose as your max a value at either 400 nm or 700 nm.) Take additional absorbance readings at 10 nm intervals larger and smaller than your initially identified max to increase the likelihood that you have found the true max.

  1. Plot wavelength (x-axis, long edge of graph) versus absorbance (y-axis, short edge of graph) for your sample. Try to draw a smooth curve which includes all of your data points. Your graph should nearly fill the paper.
  2. Identify the principal wavelength(s) absorbed and relate this to the color you observe.

When you are finished with this part, your plotted spectrum should resemble one of the traces shown below in the combined spectra of the three brightly colored solutions (but with the graph filling the page and a different max).

word image 3233

 

PRINT IN LANDSCAPE ORIENTATION; INCLUDE ALL REQUIRED INFORMATION.

Part 2 Absorbance and Concentration

Use the same absorbing substance investigated in the first part of this experiment.

  1. To prepare two of standard solutions with known concentrations from your sample stock solution:

Use a clean buret to deliver into two labeled test tubes approximately 3 and approximately 7 mL respectively of your stock solution. (Do not attempt to deliver an exact volume, but whatever volume is delivered should be recorded precisely.) Deliver appropriate amounts of distilled water from another buret to each of the two tubes to obtain a total of ~10 mL of solutions. We will assume that the volumes delivered from the buret to the test tube are additive, although this is not always the case.

Calculate the concentration of each of these two solutions you prepared. (Remember, Vconc.Mconc. = Vdil.Mdil..) These two solutions, a sample of distilled water, and the stock solution, will be your four standard solutions.

  1. Set the Spec 20 to the max you found from Part I and calibrate the instrument.
  2. Read the absorbance of each of the four standard solutions at the max. Use the same cuvette for each solution, rinsing the cuvette first with distilled water and then with the solution to be used before filling it to take a reading.
  3. Plot the concentration (x-axis) versus absorbance (y-axis) for each standard solution.
  4. Draw the best straight line for your data which shows a linear relationship between absorption and concentration of the absorbing substance.

Part 3 Determination of the Concentration of an Unknown

  1. Obtain a sample (the same absorbing substance as you used in Parts 1 and 2) of unknown concentration from your lab instructor.
  2. Measure its absorbance at the wavelength of maximum absorption and determine its concentration using your graph from Part 2.
  3. After your instructor reports the true concentration, calculate your percent error.

If you compare the visual intensity of the color of your unknown to known reference solutions, can you approximate its concentration? _________________________________________________________________________________________________________

 

ABSORPTION SPECTROSCOPY

_________________________________________________________________________________________________________

Part 1

Absorbing Species _________________________________

Concentration (M) _______

Wavelength Absorbance

Wavelength of Maximum Absorption _________ nm

Color Observed Color(s) Absorbed

Show all calculations in your notebook; use the area below the table.

Part 2

Sample

Number

Vol. stock sol’n. (Vconc.)

Vol. H2O added

Diluted

Vol. (Vdil.)

Conc.

(Mdil.)

Absorbance

1 (pure water)

 

…..

…..

…..

 

 

 

2

 

 

3

 

 

4 (stock sol’n)

 

 

 

 

 

…..

 

 

 

 

 

…..

 

 

 

 

 

…..

 

 

 

 

 

 

 

 

 

 

 

 

 

Part 3 Unknown Number Absorbance _________ nm

Experimentally Determined Concentration _________

True Concentration _________ (value obtained from lab instructor)

Percent Error _________

DATA ANALYSIS SUPPLEMENT

Data: After you have recorded all the wavelengths and absorbances, enter your own data into the appropriate Web-based data entry form, using a computer in the lab. In your measurements and calculations, remember to pay attention to uncertainties and significant figures.

Results: Outside of class, retrieve your data in spreadsheet form (along with that of the entire class) from: http://chem.gmu.edu/results/. Then use Excel to calculate the concentrations of the two standard solutions.

Your lab report should include the following in addition to the pages from your lab notebook:

 An Excel sheet “Raw Data” showing the raw data for the entire class (as downloaded). A typical example of the raw data sheet is shown below:

Raw Data

 

 

A

B

C

D

E

F

G

H

I

J

K

L

M

1

ID

 

Metal

 

Lambda _max nm

Conc_

stock M

Vi_A mL

Vf_A mL

Vi_B mL

Vf_B mL

Abs_stock

 

Abs_A

 

Abs_B

 

Abs_water

 

Abs_unknown

 

2

1

Nickel

660

0.200

3.18

9.92

6.98

10.00

0.31

0.09

0.21

0.31

0.12

3

2

Cobalt

510

0.200

3.10

9.98

7.10

10.25

1.07

0.30

0.73

0.00

0.50

 

Explanation:

  • Lambda_max: wavelength of maximum absorption o Conc_stock: concentration of stock solution o Vi_A: volume of stock solution in standard solution 2 o Vf_A: diluted volume of solution 2 o Vi_B: volume of stock solution in standard solution 3 o Vf_B: diluted volume of solution 3
  • Abs_stock: absorbance of stock solution o Abs_A: absorbance of standard solution 2 o Abs_B: absorbance of standard solution 3 o Abs_water: absorbance of water o Abs_unknown: absorbance of unknown solution

 

 A 2nd Excel worksheet “Results” showing the calculated results for each student in the class – each in a separate row. Make sure to show an example calculation for each column. This can be done by copying the equation from the formula bar to a blank cell below the elements in the table. An apostrophe in front of it makes it text instead of an equation. Each table should be formatted as in previous classes with lines separating the table headings from the contents. A line below the contents should also be drawn. A typical example of the results sheet is shown below:

Results

 

A

B

C

D

E

F

G

H

I

J

K

L

1

.16.1

[1]

M

[2]

M

[3]

M

[4]

M

Abs_

1

 

Abs_

2

 

Abs_

3

 

Abs_

4

 

Slop

e

 

Interce pt

 

[Unk

]

M

 

2

1

0.000

0.060

0.140

0.200

0.31

0.09

0.21

0.31

2.25

0.0360

0.113

3

2

0.000

0.060

0.467

0.200

0.00

0.30

0.73

1.07

0.56

0.0990

0.180

 

Explanation:

o [1]: concentration of water o [2]: concentration of standard solution 2 o [3]: concentration of standard solution 3 o [4]: concentration of stock solution o Abs_x: absorbance of solution x (x = 1-4) o Slope: slope of the absorbance on y- axis vs. concentration on x- axis o Intercept: intercept of the absorbance on y- axis vs. concentration on x- axis o [Unk]: experimentally determined concentration of unknown solution. The equation that relates absorbance to concentration is Abs_unknown = slope * [Unk] + intercept. Use the straight line slope, intercept and the unknown’s absorbance to calculate this value.

 

  • A 3rd Excel sheet “Spectrum” showing an XY (Scatter) chart between the wavelengths on xaxis and absorbance on y-axis of the absorbing species. Use the appropriate line option of the graph to plot a smooth line joining all points. Check whether the max is the same as you have determined in lab.
  • A 4th Excel chart “Calibration Curve” showing an XY (Scatter) chart between the concentration on x-axis and absorbance on y-axis for each standard solution. Use the trendline option to obtain the best straight line and the equation. Refer to http://chem.gmu.edu/Graphing with Excel.doc

Printing Tables:

  • Large tables should be printed in the landscape mode and sometimes it might be necessary to go to “page setup” and select “Fit to 1 page” to make the page fit on one page.
  • Include column and row headings on the printout so that the instructor can more easily understand equations used. This can be done by going to File → Page Setup → Sheet tab → Select “Row and Column Headings”.

Discussion: In your discussion section address the following questions:

  • Why do you need to calibrate the instrument every time you change the wavelength? What precautions do you need to take while calibrating the instrument?
  • What could be the reasons for the difference between the experimentally determined concentration and the true concentration for your unknown?
  • Why is it necessary to measure the absorbance at max?
  • Can you visually estimate the concentration from the relative color intensity? Explain.
  • What is the max from the Spectrum graph and is it the same as what you used for the absorbance measurements of solutions?
  • According to Beer’s Law, absorbance is proportional to concentration. Does your calibration curve support this?

 

 

 

Experiment 17 PERIODIC PROPERTIES: CATION ANALYSIS

 

 

 

 

word image 1006 word image 1007 word image 1008 word image 1009 word image 1010

word image 1011 In the nineteenth century, chemists attempted to classify and organize their knowledge of the chemical elements in a useful way. The Russian chemist, Dimitri Mendeleev, developed a scheme for classifying the elements according to similarities in their observed physical and chemical properties; this scheme was the first periodic table of the elements. Later, it was observed that if the elements are arranged in order of increasing atomic number (number of protons), elements with similar properties tend to appear at somewhat regular intervals in the list. The modern Periodic Table of Elements is constructed of vertical columns, each of which contains one “family” of elements. The families are often called by their Table column, or group, numbers, but some of the oldest known groups have common names (such as the Alkali Metals which form alkaline, or basic, solutions when mixed with water).

We now know that the similar physical and chemical properties of the families of elements are determined by their similar electronic structures, that is, by the arrangement of their electrons around the nucleus. A very stable arrangement of these electrons occurs when the outermost level contains eight electrons, as is the case with the noble gas elements. The driving force for many chemical reactions is to acquire an octet (eight electrons) in the atom’s outermost energy levels and thus to achieve a “noble gas configuration.”

Table 1 lists the physical properties of some members of two families in the Periodic Table– the Alkaline Earth metals (Group II) and the Halogens (Group VII). For both groups, as in other families, there are general trends: atomic and ionic radii increase with increasing atomic number within a family, and ionization energies and electronegativities decrease with increasing atomic number. Refer to your textbook for an explanation of the correlation of these trends with atomic structure.

The elements within each family group all have the same electronic configuration in their outer electron shells, that is, the highest main energy level. The Alkaline Earth metals have two electrons in an s atomic orbital; removing these two s electrons results in a noble gas configuration for the resulting divalent-cation. Therefore, the most stable oxidation state of the metal is +2.

M M2 + 2e

The halogens have seven electrons in their outermost electron shells: two in an s-orbital and five in the p-orbitals. Gaining one electron gives a halogen a noble gas configuration and thus the resulting mono-anion is usually quite stable. (Other oxidation states of the halogens are also known.)

 

Table 1

PHYSICAL PROPERTIES OF THE ALKALINE EARTH ELEMENTS (GROUP II)

 

Property Magnesium Calcium Strontium Barium

ElectronicConfiguration [Ne]3s2 [Ar]4s2 [Kr]5s2 [Xe]6s2

 

Atomic

Radius (nm) 0.136 0.174 0.191 0.198

 

Ionic Radius

M2+ (nm) 0.065 0.099 0.113 0.135

 

Ionization

Energy 1st 738 589 550 503

(kJ mol-1) 2nd 1451 1148 1061 965

 

Electronegativity 1.2 1.0 1.0 0.9

 

 

PHYSICAL PROPERTIES OF THE HALOGENS (GROUP VII)

 

Property Chlorine Bromine Iodine

 

Electronic

Configuration (Ne)3s23p5 (Ar)3d104s24p5 (Kr)4d105s25p5

 

Atomic

Radius (nm) 0.099 0.114 0.133

 

Ionic Radius

(X) (nm) 0.181 0.195 0.220

 

Ionization

Energy

(kJ mol-1) 1254 1141 1007

 

Electron Affinity

(kJ mol-1) 348 324 295

 

Common Oxidation States -1, +1, +3, +5, +7 for all three

Electronegativity 3.0 2.8 2.5

X + e

X

 

Like other metals, the Alkaline Earth metals have low electronegativities and tend to form cations. The di-cation members of this group form ionic salts with many negatively charged ions. The Group II cations can be distinguished from other cations and from each other by their reactions and solubility properties with the polyatomic ions:

Sulfate, SO42− carbonate, CO32− oxalate, C2O42− chromate, CrO42−

Halogens exist in their elemental form as diatomic molecules, for example Br2. Because halogens tend to gain an electron and become a halide anion, they have the chemical property of being good oxidizing agents. The relative oxidizing power of an element is determined by its ability to gain an electron from another element or ion. Among the elements of the halogen family, some are better oxidizing agents than others. Thus we say that bromine is more reactive or a better oxidizing agent than iodine (i.e. it has a greater tendency to gain electrons).

THE EXPERIMENT

During this laboratory period you will explore the chemistry of the Alkaline Earth cations; during the next period, you will learn about the oxidation-reduction chemistry of the Halogens and Halide anions. The Group II and Group VII elements can combine to form ionic salts of the general chemical formula MX2 (M = metal, X = halogen). You will be given a solution of one of the MX2 salts and asked to identify which Group II cation and which Group VII anion you have on the basis of the chemical properties determined in each part of the experiment.

For this laboratory experiment, you will classify the solubility properties of salts of the Group II cations. In order to express the information, you need to become proficient in writing ionic equations (chemical equations which involve cations and anions).

Many ionic compounds do not dissociate appreciably in aqueous solution; although composed of ions, they are confined in their crystal lattice. To represent these solid ionic compounds, we write the chemical formula, leaving out charges. For example, copper sulfide is composed of copper cations and sulfide anions, but it is undissociated in aqueous solution. Its formula would be written as CuS. The formula for a crystalline compound which is present in its undissociated form is usually followed by the notation (s) which means it is in the solid state.

Any compound which is significantly ionized (dissociated) in water will be written as the formula for the ion with a numerical superscript indicating its charge. For example, solid sodium chloride (NaCl, table salt) will dissolve in water forming sodium cations and chloride anions.

H2O

NaCl(s) Na+ + Cl

It is understood that the ions are surrounded by water molecules which stabilize the ions and keep them separated. These solvated ions may also be written as Na+(aq) and Cl (aq).

If we want to write an equation for compounds which have dissociated in aqueous solution, we will enclose in parentheses the symbols of all ions derived from the same compound formula. For example, potassium bromide and lithium hydroxide when dissolved in water form solvated ions:

H2O

KBr(s) + LiOH(s) (K+ + Br) + (Li+ + OH)

Inspect the following chemical equation:

(Ag+ + NO3) + (Na+ + Cl) AgCl(s) + Na+ + NO3

This equation should be interpreted as a reaction taking place when an aqueous solution of dissociated silver nitrate is mixed with an aqueous solution of dissociated sodium chloride. When the Ag+ and the Cl ions encounter each other in solution, they combine into an undissociated solid ionic compound, silver chloride. The ions are confined to their crystal lattice and the AgCl, which is insoluble in water, precipitates out of the aqueous solution. Sometimes a vertical arrow is written beside the compound to indicate this precipitation reaction. Notice that the nitrate and sodium ions are not involved in the overall precipitation reaction. They are merely “spectators” in the solution.

To write a net ionic equation, cross out the spectator ions on both sides of the chemical equation and rewrite the resulting equation:

Ag+ + Cl AgCl(s) 

PROCEDURE

 

NOTE: Please do not use more than the indicated amount of reagent. Do not insert pipets, medicine droppers, or stirring rods into the reagent bottles. Instead, pour small amounts of reagents into a beaker or flask for your own use.

In each part of this experiment, you will be mixing together aqueous salt solutions and observing whether any insoluble substance is formed. If an insoluble substance is formed, the mixture will appear cloudy or the solid may fall to the bottom of the test tube (precipitate out of solution). All of the mixtures eventually form precipitates, so record your results within the first minute after mixing and stirring. Also, make some qualitative comparison of the relative amount and color of each precipitate formed with each reagent.

Relative Solubilities of Some Alkaline Earth Salts

  1. On the reagent bench are 0.1-M solutions of the nitrate salts of Group II cations Mg2+, Ca2+, Ba2+, and Sr2+. (Remember, all nitrate salts are soluble in aqueous solution.) Transfer about 1 mL (20 drops) of each nitrate salt solution to each of four small test tubes which have previously been cleaned and rinsed thoroughly with distilled water (the tubes need not be dried). Label your test tubes.
  2. To each of these test tubes add 1 mL of 1-M sulfuric acid and mix well. (H2SO4 is soluble in water and acts as a source of the sulfate anion). Remember to rinse your stirring rod well between use in each of the test tubes.

If any of these alkaline earth cations form an insoluble sulfate salt with sulfuric acid, then a cloudiness or precipitate will form in that test tube. If the contents of the test tube remain transparent, then the sulfate salt of the alkaline earth cation is soluble. Record your results in a table in your notebook. Write IP (Insoluble Precipitate) or S (Soluble) in the appropriate columns. State the color of the solid that is formed.

  1. Clean your test tubes by rinsing well first with a copious amount of tap water, then with distilled water.
  2. Repeat the above solubility experiment testing the same Group II cations with 1-M sodium carbonate solution instead of the sulfuric acid. Na2CO3 is a soluble salt and is the source of the carbonate anion.
  3. Continue the experiment testing for the solubility of the Group II oxalates and chromates using 0.25-M (NH4)2C2O4 (ammonium oxalate) and 1.0-M K2CrO4 (potassium chromate), respectively.
  4. After you have finished discovering which Group II cations form insoluble and soluble salts, you are ready to analyze an aqueous sample of a Group II nitrate whose identity is unknown to you.
    • First devise a plan for analyzing your unknown using just two reagents. Obviously, you could identify your unknown by simply following the same steps as with your knowns, but your goal here is to understand the solubility properties of the cations sufficiently well that you can actually use them for an efficient logical analysis.
    • Have your instructor check both your results with the knowns and your scheme for analyzing your unknown.
    • After getting your instructor’s approval, obtain an unknown solution from your instructor and determine the identity of the Group II cation it contains.

_________________________________________________________________________________________________________

 

PERIODIC PROPERTIES: CATION ANALYSIS

_________________________________________________________________________________________________________ Group II

Nitrate SO42− CO32− C2O42− CrO42−

Salt

_________________________________________________________________________________________________________

Mg2+

 

 

 

 

 

Ca2+

 

 

 

 

 

Sr2+

 

 

 

 

 

Ba2+

 

 

 

 

 

 

Unknown#

 

 

 

 

Unknown #

    

_________________________________________________________________________________________________________

Write balanced net ionic equations for each solubility test you performed that gave an insoluble salt as the product of a reaction.

Conclusions

State which Group II cation forms the fewest insoluble salts and which cation forms the most insoluble salts. Describe how the relative solubilities vary according to the cation’s position in the Periodic Table.

Do you expect RaCrO4 to be soluble or insoluble in water? Why?

 

Experiment 18 PERIODIC PROPERTIES: ANION ANALYSIS

 

 

 

 

word image 1012 word image 1013 word image 1014 word image 1015 word image 1016

A common chemical reaction is the oxidation-reduction (or redox) reaction. An oxidation takes place when an atom or molecule loses one or more electrons; a reduction occurs when an atom or molecule gains one or more electrons. Therefore, a redox reaction involves one species gaining electron(s) while another species loses electron(s). The number of electrons gained must equal the number of electrons lost. In redox processes, an oxidation reaction and a reduction reaction must always occur together.

The species which is oxidized gives its electron(s) to another substance which is then reduced. In essence, the oxidized species acts as a reducing agent. Likewise, the species which gains electron(s), and is thereby reduced, acts as an oxidizing agent.

word image 1017

Oxidized species

Reduced species

 loses electrons

 gains electrons

 increases its oxidation state

 decreases its oxidation state

 is a reducing agent

 is an oxidizing agent

 

An important chemical property of the Group VII halogens is that they are all relatively easily reduced (good oxidizing agents) compared to non-halogens. The reaction for the reduction of a halogen diatomic molecule to two halide anions can be written:

X X + 2e 2

X

(1)

The diatomic halogen, X2, has 14 valence electrons and the two halide anions have a total of 16 valence electrons, thus the halogen atoms have been reduced by gaining electrons. Species that are oxidized in one reaction may be reduced in another reaction. There are some powerful oxidizing agents which can oxidize the halide anions to form a diatomic halogen molecule by removing halide electrons:

2 + 2e (2)

X

X

X

As you look at equations (1) and (2) above, notice that they are the reverse of each other. Halogens are reduced to halide ions and halide ions are oxidized to halogens. (It is understood that there are spectator cations present which balance the halide anion charge.)

What might we expect to happen when two different halogen/halide pairs are mixed together? Suppose Cl and Br are chosen as the Group VII pair. There are two combinations for mixing these: chlorine/bromide and bromine/chloride. (Note that the diatomic halogen elements have the –ine suffix and the halide anions have the –ide suffix.) Although we can write balanced chemical equations for each oxidation/reduction combination, only one of these reactions will actually occur.

 

or

 

Cl2 + 2 Br−

2 Cl− + Br2

2 Cl− + Br2

Cl2 + 2 Br−

(3)

(4)

If reaction (3) takes place, it means that chlorine (which is reduced to chloride in the reaction) is a stronger oxidizing agent than bromine, because chlorine oxidizes bromide to form bromine. If reaction (4) takes place, it means that bromine is a stronger oxidizing agent than chlorine, because bromine oxidizes chloride to form chlorine. If equation (3) is read from right to left, it is identical to equation (4). Likewise, if equation (4) is read from right to left it is identical to equation (3). Only one of these reactions can take place as written.

THE EXPERIMENT

In this experiment, you will finish your two-week study of group chemical properties by investigating the relative oxidizing strengths of the halogens chlorine, bromine, and iodine. To determine the relative oxidizing abilities of the halogens, mixtures of all combinations of halogens and halides must be made and any changes observed. Afterwards, you will be given a solution containing a halide ion and asked to identify it based on the principles you have learned.

We will take advantage of the different colors and the different solubility characteristics of the halogens and halides in order to judge whether or not a redox reaction has occurred. The halide ions are all colorless in solution in contrast to the halogens, each of which is distinctively colored: Cl2 is a very pale yellow, Br2 is red-brown, and I2 is purple.

Halogens (X2) and halides (X−) have different solubilities in water and in certain organic solvents. In polar water, the ionic halides are more soluble than are the halogens; in the relatively non-polar organic solvent, the non-polar halogens are more soluble than are the halides. And finally, polar water and the non-polar organic solvent are immiscible, that is, they are insoluble in each other, and they will separate into two distinct liquid layers upon being mixed. Depending on the density of the organic solvent used, it may be either lighter or heavier than water, which means that it may be either the upper or the lower liquid layer. Your instructor will tell you whether the organic solvent is more or less dense than water.

Even though you cannot distinguish among any of the colorless halides, you can easily do so with the corresponding colored halogens. Any halogen which is present will appear as a colored solution in the organic solvent layer in the test tube.

Using our previous Cl/Br example, suppose you first mix pale-yellow Cl2 with colorless Br− and then a red-brown color develops in the organic layer. This means that Br2 (red-brown) has formed by the redox equation (3), and thus Cl2 has oxidized Br−. As a check for this result, you could then mix colorless Cl− with red-brown Br2 (equation 4). No change of the red-brown organic layer to pale yellow (Cl2) should take place if Cl2 is indeed the stronger oxidizing agent.

PROCEDURE

NOTE: Although the halogens are more soluble in the organic solvent than they are in water, it is easier to prepare solutions of the halogens in water. When you add the organic solvent to the halogen/water, the halogen will dissolve preferentially in the organic solvent phase although some halogen will also be present in the aqueous phase.

 

General Instructions for 1, 2, and 4. below

Add 1 mL of halogen solution (X2) to a clean small test tube (it need not be dry).

Then add 2 mL of the organic solvent. Shake well (cork!) and then allow the aqueous and organic layers to separate. Note the color of the organic layer. This is the color of the halogen X2.

Now, add 1 mL of the aqueous halide solution (X−) drop-by-drop; shake gently during the addition, then cork and shake.

Record the color of the organic layer only. From the color of the organic layer, state which halogen,

X2, is present. Also state whether or not a redox reaction has taken place between the mixed X2/X− pair.

Save the contents of all your test tubes to aid in identifying your unknown.

Record your observations in your notebook by drawing the outline of a test tube and noting the position and color of the organic layer. Ignore the color of the aqueous layer. It will be easier to analyze your unknown and write equations if your observations are recorded this way.

  1. First you must prepare a series of control solutions in order to observe the color of each halogen in the organic solvent. Following the General Instructions above, choose the same X for each of the three X2/X− combinations.
  2. Following the General Instructions above, choose different X2/X− pairs until all remaining possible halogen/halide combinations have been investigated.
  3. From your data, work out a method for analyzing your unknown halide anion. There is more than one route to do this; however, some are more efficient than others. See if you can find the most efficient route. Obviously, you could identify your unknown by simply following the same steps as with your knowns, but your goal here is to understand the oxidizing properties of the halogens sufficiently well that you can actually use them for a logical analysis.

Have your instructor check your results with the knowns and your analytical scheme.

  1. After approval, obtain an unknown solution from your instructor and determine the identity of the Group VII halide it contains.

When you are finished, dispose of your experimental solutions in the HOOD.

_________________________________________________________________________________________________________

PERIODIC PROPERTIES: ANION ANALYSIS

_________________________________________________________________________________________________________

 

Halogen

 

Halide

Cl−

 

Br−

 

I−

Cl2

 

 

 

 

 

 

Br2

 

 

 

 

 

 

I2

 

 

 

 

 

 

 

Unknown X− #

______________________________________________________________________________

Write balanced chemical equations for each test you performed which underwent a redox reaction.

Conclusions

  1. Formulate a statement about the relative oxidizing powers of the halogens (X2) and their positions in the periodic table. Explain your reasoning.
  2. Do you think that At2 (astatine) could oxidize I–? Could Cl− reduce F2 (fluorine)? Write appropriate equations first, consider your statement in (1.), then briefly explain.

 

CHEM 214, CHEM 272 EXPERIMENTS

 

 

word image 1018

 

Experiment 19 SYNTHESIS OF A COORDINATION COMPOUND

 

 

 

 

word image 1019 word image 1020 word image 1021 word image 1022 word image 1023

POTASSIUM TRIOXALATOFERRATE (III) TRIHYDRATE

K3[Fe(C2O4)3].3H2O = KTOF3

Many chemical combination reactions take place by the interaction of an electron-rich species (Lewis base) with an electron-deficient species (Lewis acid). In the resulting coordination compounds, the Lewis base has formally donated two electrons to form a coordinate covalent bond. A simple example of such a reaction is that between boron trifluoride (a Lewis acid) and ammonia (a Lewis base).

F H

word image 1024 + F3B NH3

B

F

N

H

F H

 

Numerous coordination compounds are known and in fact nearly all compounds of the transition elements are coordination compounds wherein the metal or metal ion is a Lewis acid. The atoms or molecules bonded to the metal are Lewis bases called ligands. Ligands can be negatively charged or neutral and they can have one or more sites (electron pairs) to bond to the central atom. Some common monodentate ligands (ligands which use only one electron pair to bond) include water, ammonia, halide ions, and cyanide ion. Bidentate ligands bond to the metal at two positions using two electron pairs. Examples include ethylenediamine and the oxalate ion.

(See Figure 1.)

A charged complex ion is a metal ion coordinated with one or more ligands. The net ionic charge on the complex ion is counterbalanced with oppositely charged ions which are not part of the coordination complex. (See Figure 2.) The complexes assume various molecular geometries depending upon the metal and ligand involved. These molecular geometries often may be deduced using the postulates of the VSEPR Theory.

For molecules with a coordination number of six, the symmetrical octahedral geometry allows the ligands to be as far apart as possible. As you recall from VSEPR Theory, if the ligands are not identical, there exists the possibility of isomerism. Because the apexes of a regular octahedron are all equivalent positions, the three bidentate oxalate ions, (C2O4)2- in the [Fe(C2O4)3]3- complex ion are in identical environments (Figure 2). Although the chloride atoms in the two isomeric compounds cis– and trans– [Co(NH3)4Cl2]+ are in equivalent environments within each complex ion,

(Figure 2) the two complex ions are not equivalent to one another. The two chloride atoms are 90° apart in the cis isomer and 180° apart in the trans isomer. These two isomers are called geometric isomers; although they have identical molecular formulas, the geometrical arrangement of their ligands in space is different. Consequently, they have somewhat different chemical and physical properties. In this example, the cis isomer is violet colored, while the trans isomer is green.

 

Figure 1

word image 3234

 

Figure 2

word image 3235

[Co(NH3)4Cl2]+ [Co(NH3)4Cl2]+

THE EXPERIMENT

The synthesis and characterization of compounds are primary tasks of a research chemist. In this experiment, you are going to prepare the coordination compound potassium trioxalatoferrate(III) trihydrate from potassium oxalate hydrate and ferric chloride, and determine the percent yield. The equation for the preparation is

3 K2C2O4•H2O + 1 FeCl3 1 K3[Fe(C2O4)3] •3H2O + 3 KCl

In the complex ion, the iron atom is at the center of an octahedron and coordinated with three bidentate oxalate ligands (Figure 2). The water molecules are not acting as ligands coordinated to iron; instead they are part of the three-dimensional crystal lattice. Notice that they are not included inside the brackets which enclose the complex ion [Fe(C2O4)3]3−. The potassium ions, also part of the crystal, act as counter-ions.

A series of experiments will be conducted during the next couple of weeks whereby you will quantitatively analyze this coordination compound for its oxalate, iron, and water content to prove its composition and purity. Because you will be calculating a percent yield as part of this experiment, an example of a percent yield calculation, using a different synthesis, is provided below.

 

Calculation of Percent Yield

In the preparation of trans-K[Cr(C2O4)2(H2O)2]•2H2O, 4.15 g of oxalic acid dihydrate (mol.wt. =

126.07 g/mol) reacted with 1.21 g of potassium dichromate (mol.wt.= 294.18 g/mol) and 1.52 g of the product coordination compound (mol.wt. = 339.19 g/mol) was isolated. What is the percent yield for this synthesis?

K2Cr2O7 + 7H2C2O4 •2H2O 2K[Cr(C2O4)2(H2O)2] •2H2O + 6CO2 + 13H2O

From the balanced equation, 1 mole of K2Cr2O7 reacts with 7 moles of oxalic acid dihydrate to produce 2 moles of the coordination compound product.

Determination of the limiting reagent:

Calculate the moles of reactants actually used for the synthesis:

Moles K2Cr2O7 = 1.21 g / 294.18 g/mol = 0.00411 mol

Moles H2C2O4•2H2O = 4.15 g / 126.07 g/mol = 0.0329 mol

The mole ratio of reactants used is 0.0329/0.00411, which is about 8:1. From the balanced equation, the reaction requires only a 7:1 molar ratio of oxalic acid to K2Cr2O7. Therefore, oxalic acid dihydrate is in excess and potassium dichromate is the limiting reagent.

Determination of theoretical yield:

Basing the calculation on the amount of potassium dichromate used, the number of moles of product formed should be twice the number of moles of K2Cr2O7 reacted because the product/reactant molar ratio from the balanced equation is 2:1.

  1. mol product

Theoretical yield  0.00411 mol reactant  0.00822 mol product 1 mol reactant

The theoretical yield in grams of product is:

(0.00822 mol)(33.079 g/mol  2.79 g product

The percentage yield of K[Cr(C2O4)2(H2O)2] •2H2O in grams is then

1.52 g

% yield  100  54.5%

2.79 g

 

 

 

PROCEDURE

Note: The quantities of reagents given here are approximate. Record the exact amount used in your notebook so that you can calculate an accurate percent yield later. Your instructor will tell you how much ferric chloride (FeCl3) solution to add in step 2; the amount to use depends on how it is prepared in the stockroom.

Preparation

  1. Weigh a 150-mL beaker and then weigh into it approximately 4.0 grams of K2C2O4•H2O

(potassium oxalate monohydrate). Add about 10 mL distilled water and then place the beaker on a hot plate and warm gently. Stir until the solution is hot and the solid is completely dissolved. Carefully remove the beaker from the heat.

  1. Add 8 mL of the ferric chloride solution to the beaker containing the hot solution. (Remember to record the molarity of the FeCl3 in your notebook.) Stir well and then cool the beaker containing the solution to 0oC in an ice-water bath and keep it at this temperature until crystallization is complete.
  2. Decant the solution, reserving all solid material in the beaker. (Also keep the mother liquor. If crystallization was not complete, more crystals will precipitate from the mother liquor.)

These crystals are your crude product and probably contain many impurities. The impurities may be removed by performing a recrystallization.

Recrystallization

  1. Add 20 mL warm distilled water to the solid crude product. Stir to dissolve. Warm gently if necessary. Do not add an excess of water. Allow the product to precipitate out of solution by cooling at 0oC. Water-soluble impurities should remain in solution in this step.
  2. Collect the product by vacuum filtration, using filter paper in a Buchner funnel. (See the experiment “Synthesis of Aspirin” for instructions on vacuum filtration.)
  3. Interrupt the suction, pour a few milliliters of ice cold distilled water on the product, stir slightly with a stirring rod to wash the crystals (be careful not to tear the wet filter paper), and then apply suction again.
  4. Wash the crystals, in the same manner, with cold methanol which will remove the solvent water. Continue to pull air through the crystals to speed drying.
  5. Spread the crystals in a pre-weighed porcelain evaporating dish (P.E.D.) and allow to dry for at least 20 minutes. Weigh the P.E.D. with the product in it and calculate the mass of the product by difference.

The molecular weight of K2C2O4•H2O is 184.23 and that of K3[Fe(C2O4)3] •3H2O is 491.25 g/mol.

Use the mass data above to calculate the actual yield and the percent yield in the synthesis of your coordination compound product. _________________________________________________________________________________________________________

SYNTHESIS OF POTASSIUM TRIOXALATOFERRATE(III) TRIHYDRATE

_________________________________________________________________________________________________________

Balanced Chemical equation for synthesis of the coordination compound: DO NOT USE ABBREVIATIONS; FORMULAS ARE REQUIRED FOR REACTANTS AND PRODUCTS.

Mass of beaker

__________ g

Mass of beaker + K2C2O4•H2O

__________ g

Mass K2C2O4•H2O

__________ g

Moles K2C2O4•H2O

__________ mol

Volume of FeCl3(aq)

__________ mL Conc. of FeCl3(aq) = _______ M

Moles of FeCl3

__________ mol

Determination of limiting reagent: (show all calculations in notebook)

Limiting reagent: ____________________

Theoretical yield of product: mole

(show all calculations in notebook)

Color and general appearance of complex:

g

Experimental yield of product: mole

g

(show all calculations in notebook)

Percent yield of product (show all calculations in notebook): ____________

 

DATA ANALYSIS SUPPLEMENT

Data: After you have recorded all of the masses, volumes and molarities, enter your own data into the appropriate Web-based data entry form, using a computer in the lab. In your measurements and calculations, remember to pay attention to uncertainties and significant figures.

Results: Outside of class, retrieve your data in spreadsheet form (along with that of the entire class) from: http://chem.gmu.edu/results/. Then use Excel to calculate the moles of K2C2O4•H2O and FeCl3, the Theoretical yield, Experimental yield and Percent yield of K3[Fe(C2O4)3] 3H2O

Your lab report should include the following in addition to the pages from your lab notebook:

  • An Excel sheet “Raw Data” showing the raw data for the entire class (properly formatted).
  • A 2nd Excel sheet “Class Results” showing the calculated results for each student in the class – each in a separate row. Make sure to show an example calculation for each column. This can be done by copying the equation to a blank cell below the elements in the table. An apostrophe in front of it makes it text instead of an equation. Each table should be formatted as in previous classes with lines separating the table headings from the contents. A line below the contents should also be drawn. The columns should include:

o Moles of K2C2O4•H2O o Moles of FeCl3 o Theoretical yield of K3[Fe(C2O4)3] 3H2O o Experimental yield of K3[Fe(C2O4)3] 3H2O o Percent yield of K3[Fe(C2O4)3] 3H2O

  • An Excel chart showing a histogram of the distribution of percent yield results. Use equally spaced “bins” chosen to display the “spread” in the class results.

 

Printing Tables:

  • Large tables should be formatted to fit on one page. This can be done by (a) printing in landscape mode and (b) using the File|Page Setup menu to select “Fit to 1 Page”
  • You must Include column and row headings on the printout. This can be done by going to

File|Page Setup and using the Sheet tab with Select “Row and Column Headings”.

 

Discussion: In your discussion section address the following questions:

  • What percentage of the students in your section obtained percent yields that were higher than the maximum calculated in the previous question?
  • What experimental error(s) would cause the percent yield to come out too high?

Sample data—unformatted

 

word image 3236

 

 

Sample histogram–

 

word image 3237

Data tables and figures are to be printed in landscape orientation.

 

Experiment 20 DETERMINATION OF WATER OF HYDRATION IN KTOF3

 

 

 

 

word image 1025 word image 1026 word image 1027 word image 1028 word image 1029

Many inorganic compounds exist as hydrates in the solid state. When these substances crystallize from an aqueous solution, water molecules are incorporated in the crystal lattice. Copper sulfate, for example, crystallizes as the pentahydrate, CuSO4•5H2O, in which there are five molecules of water in the solid lattice for every one formula unit of CuSO4. Sodium sulfate crystallizes as the decahydrate, Na2SO4•10H2O.

word image 1030 The water molecules are usually weakly bonded to the metal ion, though they may also be attached to the anion. In either case, the water molecules are not strongly held and are lost from the lattice upon heating, leaving the anhydrous (water-free) compound. When the compound is dehydrated, its crystal structure and its color may change. Thus, anhydrous compounds often have a very different appearance from their hydrated counterpart. CuSO4•5H2O is a bright blue crystalline substance whereas CuSO4 is a white powder.

THE EXPERIMENT

In the first week of this series of experiments you prepared the coordination compound, potassium trioxalatoferrate (III) trihydrate, K3[Fe(C2O4)3] •3H2O. Today, you will analyze the hydration number for your product and compare it to the known hydration number. You will also set up the experiment to gravimetrically determine the iron content of the coordination compound.

The water molecules, which are a part of the crystal lattice in the compound you prepared, cannot escape under normal air drying conditions. They can, however, be forced out of the crystalline lattice and leave behind the anhydrous salt, K3[Fe(C2O4)3] (potassium trioxalatoferrate).

To accomplish this dehydration, you will heat the hydrated salt on a hotplate. The difference in the weights of the hydrated salt and the anhydrous salt will allow you to calculate the hydration number and the weight percent water present in the originally prepared coordination compound.

K3[Fe(C2O4)3]•3H2O K3[Fe(C2O4)3] + 3H2O

(hydrated) (anhydrous)

Calculation of Percent Composition by Weight

The weights of elements or combinations of elements making up a mole of compound may be calculated from the molecular formula of the compound and the atomic weights of its constituent atoms. For example, the percent by weight of carbon in octane (C8H18) is calculated from the molecular weight of octane (114.23 g/mol) and the atomic weight of carbon (12.01 g/mol). In one mole of octane there are 8 moles of carbon. Therefore, the weight of carbon in one mole of octane is 8 (12.01 g) = 96.08 g. The weight percent carbon is then:

word image 1031

Likewise, the weight percent of water in a hydrate salt can be calculated.

Weight of water = (5 moles H2O ) x (18.02 g/mol) = 90.10 g H2O in the compound

Weight of 1 mole of CuSO4•5H2O = 249.69 g

word image 1032

Problems (show work):

  1. Calculate the wt.% potassium (39.10 g/mol) present in the synthesized hydrated KTOF3 compound— K3[Fe(C2O4)3]•3H2O (mol.wt. = 491.25 g/mol) .
  2. Calculate the wt% H2O that should be present in zinc chloride hexahydrate, ZnCl2•6H2O (mol.wt.= 244.38 g/mol).

PROCEDURE

  1. Precisely weigh approximately 1.0 g of your compound hydrate in an evaporating dish. Label the dish with your name and place it on the tray to be put in the oven.
  2. After the sample in the evaporating dish has been removed from the oven and has cooled to room temperature, weigh the anhydrous sample in the evaporating dish.
  3. If the drying oven is not available——Place the evaporating dish containing the sample on a hotplate. Heat at 1/3 power (Setting 3) for 45 minutes; cool to room temperature for 10-15 minutes, then weigh the anhydrous sample in the evaporating dish.
Calculations
  1. Calculate the moles of anhydrous salt formed after the dehydration. The molecular weight of the anhydrous salt is 437.20 g/mol.
  2. From the weight of water lost during the dehydration, calculate the moles of water lost. The molecular weight of water is 18.02 g/mol.
  3. Calculate the mole ratio of water to hydrated potassium trioxalatoferrate in your sample. Ideally, there should be an integral number of water molecules for each salt molecule and the ratio should be the same as for the balanced dehydration equation. If you do not obtain such a whole number ratio in your calculation, round off your answer.
  4. Calculate the theoretical weight percent water in the hydrated salt. The molecular weight of the hydrated salt is 491.25 g/mol.
  5. Calculate the experimental weight percent water in your sample of hydrated salt from the weight of the sample before dehydration and the weight of water lost.
  6. Calculate the percent error in the determination of the weight percent of water in your potassium trioxalatoferrate trihydrate sample.

 

 

_________________________________________________________________________________________________________

DETERMINATION OF WATER OF HYDRATION

_________________________________________________________________________________________________________

Show all calculations.

Balanced chemical equation for dehydration:

Weight of container

__________ g

Weight of container and hydrated salt

__________ g

Weight of hydrated salt

__________ g

Moles of hydrated salt

(Show calculation in lab notebook.)

__________ mol

Weight of container and anhydrous salt

__________ g

Weight of anhydrous salt

__________ g

Moles of anhydrous salt

(Show calculation in lab notebook.)

__________ mol

Weight of H2O lost

__________ g

Moles of H2O lost

(Show calculation in lab notebook.)

__________ mol

Moles of water released per mole hydrated salt

(Show calculation in lab notebook.)

__________

Theoretical weight percent of water in hydrated salt

(Show calculation in lab notebook.)

__________

Actual weight percent water in hydrated salt

(Show calculation in lab notebook.)

__________

% Error

_________

(Show calculation in lab notebook.)

DATA ANALYSIS SUPPLEMENT

Data: After you have noted down all the masses, enter your own data into the appropriate Webbased data entry form, using a computer in the lab. In your measurements and calculations, remember to pay attention to uncertainties and significant figures.

Results: Outside of class, retrieve your data in spreadsheet form (along with that of the entire class) from: http://chem.gmu.edu/results/. Then use Excel to calculate the Weight of hydrated salt, Moles of hydrated salt, Weight of anhydrous salt, Moles of anhydrous salt, Weight of H2O lost, Moles of H2O lost, Moles of water released per mole anhydrous salt, Theoretical weight percent of water in hydrated salt, Actual weight percent water in hydrated salt, and % Error.

Your lab report should include the following in addition to the pages from your lab notebook:

 An Excel sheet “Raw Data” showing the raw data for the entire class (properly formatted).

A typical example of the unformatted raw data sheet is shown below:

DATA ANALYSIS SUPPLEMENT

RAW DATA

 

A

B

1

g_K3FeOx3_3H2O

g_K3FeOx3

2

12.123

10.345

3

11.867

9.453

 

Explanation:

g_K3FeOx3_3H2O = Weight of hydrated salt g_K3FeOx3= Weight of anhydrous salt

  • A 2nd Excel worksheet “Results” showing the calculated results for each student in the class – each in a separate row. Make sure to show an example calculation for each column. This can be done by copying the equation to a blank cell below the elements in the table. An apostrophe in front of it makes it text instead of an equation. Each table should be formatted as in previous classes with lines separating the table headings from the contents. A line below the contents should also be drawn. The columns should include:
    • Mass of H2O lost from the compound after heating.
    • Moles of water lost from the hydrated sample
    • Moles of the hydrated salt used o Mole of water divided by the mole of hydrated compound o Experimental Mass % H2O in KTOF3 o % Error of Experimental Mass % H2O in KTOF3

An Excel chart showing a histogram of the distribution of Actual weight percent results. Use equally spaced “bins” chosen to display the “spread” in the class results. Refer to the histogram tutorial on the website.

Printing Tables:

  • Large tables should be printed in the landscape mode and sometimes it might be necessary to go to “page setup” and select “Fit to 1 page” to make the page fit on one page.
  • Include column and row headings on the printout so that the instructor can more easily understand equations used. This can be done by going to File → Page Setup → Sheet tab → Select “Row and Column Headings”.

Discussion: In your discussion section address the following questions: o What is a Hydration number?

  • How are water molecules bonded to the metal ion or anion? What happens to them on heating?
  • What may happen to the crystal structure and color of a compound when it is dehydrated?
  • What are the possible sources of error in this experiment? Refer in your answer to the class histogram as well as to your own results.
  • Discuss your % Error in terms of errors in measurements. What is the major source of error?
  • Using the histogram, discuss how your mole ratio compares with the mole ratio of the rest of the class.
  • Discuss the shape of the histogram. What does it tell you about the nature of the error in the class’ measurements?

 

 

Experiment 21 REDOX TITRATION OF OXALATE IN KTOF3

 

 

 

word image 1033 word image 1034 word image 1035 word image 1036 word image 1037

 

Potassium permanganate, KMnO4, a reagent that is readily available and inexpensive, is widely used as an oxidizing agent and titrant. Permanganate titrations do not require an indicator because the endpoint is easily observed; permanganate ion is a vivid dark purple (absorption maximum = 526 nm) while the reduced manganese(II) ion is virtually colorless. The first slight excess of permanganate causes the solution to turn a pink color when all oxidizable material has been consumed.

In an acidic solution, the permanganate ion oxidizes oxalate (or oxalic acid) to produce carbon dioxide and water while the permanganate ion is concurrently reduced to manganese (II) as follows:

word image 1038 5 C2O42− + 2 MnO4 + 16 H+ 2 Mn2+ + 10 CO2 + 8 H2O (l)

(The spectator ions are not shown in this net ionic equation.) When oxalate ion is oxidized in acid solution, the oxidation half-reaction is

C2O42− 2 CO2 + 2 e− (2)

The oxidation states of the two carbon atoms in the oxalate anion change from a total of +6 to a total of +8, corresponding to a loss of 2 electrons. When permanganate ion is reduced in an acid solution, the reduction half-reaction is

MnO4 + 8 H+ + 5 e Mn2+ + 4 H2O (3)

The oxidation state of manganese changes from +7 to +2, corresponding to a gain of 5 electrons. Balancing with respect to electron loss and gain and then adding the two half reactions gives the net oxidation-reduction reaction occurring:

5 ( C2O42− 2 CO2 + 2 e )

(4)

2 ( MnO4 + 8 H+ + 5 e Mn2+ + 4 H2O )

_______________________________________________________________

(5)

5 C2O42− + 2 MnO4 + 16 H+ 2 Mn2+ + 10 CO2 + 8 H2O

(1)

 

which indicates that 2 moles of KMnO4 are reduced (10 electrons gained) for every 5 moles of C2O42− oxidized (10 electrons lost) during the titration.

THE EXPERIMENT

This week you will analyze your potassium trioxalatoferrate coordination compound K3[Fe(C2O4)3•3H2O] for its oxalate content by performing an oxidation-reduction titration with

KMnO4. Oxidation-reduction reactions, like acid-base reactions, are widely used for the titrimetric analytical determination of substances. You may find it helpful to reread the material on the volumetric analysis you performed previously– “Titration of Vinegar”, and the explanation of oxidation/reduction in “Chemical Properties of the Halogens.”

Potassium permanganate solutions are not entirely stable over long periods of time so you must protect the solution from heat and light as much as possible. The concentration of the permanganate titrant must be standardized prior to using it to analyze your compound. Sodium oxalate, Na2C2O4, which reacts with permanganate according to Equation (1), is often used as the primary standard. There will not be enough time for you to standardize your own permanganate solutions; this will already have been done by the prep room. Use the molarity shown on the reagent bottle to perform your calculations.

The logic of the analysis is as follows:

  1. The moles of potassium permanganate consumed in the titration are calculated from the titration data; you know both the molarity and the volume of permanganate solution used.
  2. The moles of oxalate which were oxidized in the titration reaction are obtained from the moles of permanganate used in the titration and the known redox ratio of the two reacting species (See Equation 1.)
  3. The mass of oxalate which was oxidized is calculated from the moles and the molecular weight of oxalate, 88.02 g/mole.
  4. If all of the above calculations are combined into one statement:

liter KMnO4 mol KMnO4 5 mol (C2O4 )2 88.02 g (C2O4 )2

g C2O42− titrated =  liter sol’n 2 mol KMnO4 mol (C2O4 2

)

or Vol. (KMnO4 ) • M (KMnO4)• reacting ratio • M.W. (oxalate)

 

 

Example

Suppose 13.25-mL of 0.2203-M KMnO4 were required to titrate a sample of Fe(C2O4)3. How many grams of oxalate were titrated?

Moles KMnO4 reduced = 0.01325 liter  0.2203 molliter  2.91910-3 mol

5 mol C O

Moles oxalate oxidized = 2.919103 mol KMnO42 4  7.29710-3 mol

2 mol KMnO4

Grams oxalate = 7.297103 mol  88.02 gmol  0.6423 g

Problems (show work)

  1. If 17.60 mL of 0.122-M KMnO4 were required to titrate a sample of K3[Fe(C2O4)3], how many moles of C2O42- are in the sample? What is the weight of oxalate in the sample?
  2. Calculate the weight percent C2O42− (M.W.=88.02) in each of the following: H2C2O4 (M.W.=90.04), CaC2O4 (M.W.=128.10), and K2[Cu(C2O4)2].2H2O (M.W.=417.36).
  3. Titration of an oxalate sample gave the following weight percentages: 15.92%, 15.87%, and 16.03%. Calculate the average and the standard deviation. Should a fourth titration be performed?
  4. In a standardization procedure, 0.251 g of sodium oxalate, Na2C2O4 (M.W.=134.00), required 29.25 mL of a KMnO4 solution to reach the end point. How many moles of oxalate were oxidized? How many moles of permanganate were reduced? What is the molarity of the KMnO4 solution?

Answers to Selected Problems

1). 5.37 x 10−3 moles; 0.473g

2). 97.8%; 68.7%; 42.2%

4). 1.87 x 10−3 moles; 7.49 x 10−4 moles; 0.0256 M

 

 

PROCEDURE

Caution: Both potassium permanganate and sulfuric acid are caustic. Do not allow either to come in contact with skin or clothes. In case of accident, wash the affected area with copious amounts of water and notify the instructor.

  1. Into each of three separate 125-mL Erlenmeyer flasks, weigh out to the nearest milligram about

0.1 g of the coordination compound. Add about 20 mL of 1.0-M sulfuric acid.

  1. Rinse a buret with two 5 mL portions of the KMnO4 solution and then fill the buret about half-full with KMnO4. Make sure there are no air bubbles in the tip of the buret. Because the titrant is so darkly colored, it is necessary to read the level of the liquid from the top of the meniscus.
  2. Perform each titration as follows:
      1. Carefully stir (with a stirring rod) the solution while warming it on a hot plate until the oxalate coordination compound dissolves and the temperature of the solution is 80-90°C. Do not boil the sample and do not use the thermometer as a stirring rod.
      2. Titrate the oxalate solution with the permanganate in the buret, stirring constantly, and keeping the solution above 60°C at all times. Do not use a thermometer as a stirring rod.
      3. As the endpoint is approached, add the permanganate dropwise allowing each drop to decolorize before adding the next. The end point is when the faintest shade of pink remains even after the solution has been allowed to stand for 15 sec.

*** (Ordinarily, the endpoint in a permanganate titration is not permanent. Due to the action of dissolved reducing species in the water, the permanganate is slowly reduced and the color fades. If the pink color persists up to 15 sec. while you are stirring, consider it the true endpoint.

***If there is any doubt as to whether or not an end point has been reached, record the buret reading and then add another drop of the permanganate. The development of an intense pink color indicates that the previous reading did correspond to the true end point.

4. Calculations

From the calculation of the moles of KMnO4 used in the titrations, calculate the moles of oxalate in your samples. Then calculate the weight in grams of oxalate in your sample and the weight percent of C2O42−. The three results should agree within 3 parts per thousand. If the results do not agree, titrate another sample.

It is necessary to convert the weights of oxalate titrated in the three samples to a common basis — the weight percent of oxalate in each individual sample (see Problem). The average of these three weight percents is then used for the final calculation which compares the amount of oxalate in your coordination compound to that which is actually present in a pure sample of coordination compound. _________________________________________________________________________________________________________

REDOX TITRATION OF OXALATE

_________________________________________________________________________________________________________

K3[Fe(C2O4)3] .3H2O Sample l

Sample 2

Sample 3

Mass of sample _______ g

Titration

KMnO4 Molarity _________ M

Buret Readings

_______ g

_______ g

Initial volume _______mL

_______ mL

_______ mL

Final volume _______ mL

_______ mL

_______ mL

Volume of KMnO4 _______ mL

Calculations (Show all calculations in lab notebook.)

_______ mL

_______ mL

Moles of KMnO4 reduced _______

_______

_______

Moles of C2O42− oxidized _______

_______

_______

Weight of oxidized C2O42− _______ g

_______ g

_______ g

Experimental weight percent of oxalate _______ in the sample

Average percent weight oxalate _______

Standard deviation _______

_______

_______

Theoretical weight percent oxalate in the coordination compound

_______

% Error ________

DATA ANLAYSIS SUPPLEMENT

Data: After you have noted down all the masses, volumes and molarities, enter your own data into the appropriate Web-based data entry form, using a computer in the lab. In your measurements and calculations, remember to pay attention to uncertainties and significant figures.

Results: Outside of class, retrieve your data in spreadsheet form (along with that of the entire class) from: http://chem.gmu.edu/results/. Then use Excel to calculate the Mass of samples, Volume of KMnO4, Moles of KMnO4, Moles of C2O4 2- and Weight of oxidized C2O4 2-, Experimental weight percent of oxalate, Average percent weight oxalate, Standard deviation, Theoretical weight percent oxalate, and Percent Error.

Your lab report should include the following in addition to the pages from your lab notebook:

  • An Excel sheet “Raw Data” showing the raw data for the entire class (as downloaded).
  • A 2nd Excel sheet “Class Results” showing the calculated results for each student in the class – each in a separate row. Make sure to show an example calculation for each column. This can be done by copying the equation to a blank cell below the elements in the table. An apostrophe in front of it makes it text instead of an equation. Each table should be formatted as in previous classes with lines separating the table headings from the contents. A line below the contents should also be drawn. The columns should include:

o Mass of samples o Volume of KMnO4 o Moles of KMnO4 o Moles of C2O4 2− o Weight of oxidized C2O4 2− o Experimental weight percent of oxalate for each titration o Average weight percent oxalate for each student o Standard deviation for each student o Theoretical weight percent oxalate o Percent Error

  • An Excel chart showing a histogram of the distribution of experimental weight percent results.

Use equally spaced “bins” chosen to display the “spread” in the class results.

Discussion: In your discussion section address the following questions:

  • Why is an indicator not required in permanganate titrations?
  • Why is the potassium permanganate solution standardized before you analyze your compound?
  • What are the possible sources of error in this experiment? Refer in your answer to the class histogram as well as to your own results.

 

Sample of Raw Data—unformatted

word image 3238

Sample of Processed Data—-unformatted

 

word image 3239

 

word image 3240

Experiment 22 DETERMINATION OF MOLECULAR MASS BY FREEZING POINT DEPRESSION

 

 

 

 

 

word image 1039 word image 1040 word image 1041 word image 1042 word image 1043 word image 1044

Solutions of nonvolatile solutes have higher boiling points and lower freezing points than the pure solvents used to prepare the solutions. Both of these phenomena result from the fact that the vapor pressure of a solution (of a nonvolatile solute) is lower than the vapor pressure of the pure solvent. The vapor pressure of a liquid is determined by the ability of particles at the liquid’s surface to escape into the vapor phase. In a pure liquid, all the particles at the surface are identical; in a solution, the surface is occupied by both the solvent and the solute particles. In a solution, fewer solvent particles are in a position to enter the vapor phase and the solution’s vapor pressure is lower than that of the pure solvent at all temperatures. This means that the solution will have to be heated to a higher temperature before its vapor pressure reaches that of the atmosphere, that is, the solution will have a higher boiling point than the pure solvent.

word image 1045 Although it may not be readily apparent, the same phenomenon of vapor pressure lowering that causes the boiling-point elevation also causes freezing-point depression. The freezing point of a solution is the temperature at which the solvent in a solution and the pure solid solvent have the same vapor pressure. Vapor-pressure lowering results in freezing-point depression provided that the solution does not freeze as a solid solution. The solid that forms must be pure solvent.

Both boiling-point elevation and freezing-point depression are examples of colligative properties of solutions where the magnitude of the effect depends on the number of solute molecules in a given mass of solvent. The more solute dissolved in a given solvent, the more the freezing point is lowered and the more the boiling point is elevated. Because of the importance of these colligative properties, a special concentration unit, molality (m), is employed.

moles solute

m  kilogram solvent

It has been experimentally determined that when one mole (6.02×1023 molecules) of a nonvolatile, non-electrolyte solute is dissolved in one kilogram of water, the freezing point of water is lowered from 0oC to –1.86oC. If two moles are dissolved, the freezing point is –3.72oC (21.86) and so on. Thus, the freezing point change is directly proportional to the amount of solute (not its identity) dissolved in the water solvent. The constant of proportionality, which for water is 1.86, is called the molal freezing point depression constant (kf). The molal freezing point depression constant is unique for each different solvent. For camphor, kf = 39.7. The relationship between kf, m, and the freezing point depression (tf) is:

Tf = kfm

From this equation, you can calculate any quantity if the other two quantities are known. What are the units for the molal freezing point depression constant?

Example

  1. Determine the molality of a solution of methanol solute in an aqueous solution that contains 100 g water. The experimentally determined freezing point was –2.43oC.

The freezing point depression for the aqueous solution is

Tf = 0o – (–2.43o) = 2.43o

and thus the molality is:

m = Tf / kf = 2.43o/1.86o m−1= 1.31 m

  1. How many grams of methanol are present in this solution? The molecular weight of methanol is

32.04.

From the definition of molality:

moles solute = m kg solvent = 1.31m 0.100 kg = 0.131 mol

Finally,

(32.0 g/mol) 0.131 mol = 4.19 g methanol

Problem:

A solution is prepared by mixing 2.17 g of an unknown non-electrolyte with 225.0 g of chloroform. The freezing point of the resulting solution is –64.2oC. The freezing point of pure chloroform is –

63.5oC and its kf = 4.68oC m-1. What is the molecular mass of the unknown? (Answer: 69.3 g/mol)

THE EXPERIMENT

Determining the molecular mass of an unknown substance by freezing-point depression is a classical, but still useful, technique. In this experiment you will determine the molecular mass of isopropyl alcohol (rubbing alcohol) by measuring the freezing point of an aqueous solution of the alcohol. Even though isopropyl alcohol, like many alcohols, is more volatile than water, the vapor pressure of isopropyl alcohol is low enough at 0oC to be considered negligible.

The cold bath which is used to freeze the solution is itself an example of freezing point lowering. It consists of a mixture of water, ice and rock salt.

 

 

THE PROCEDURE

The experimental freezing point apparatus consists of a test tube fitted with a rubber stopper into which is inserted a thermometer and stirrer. The cold bath is a large beaker containing the mixture of ice, water and rock salt.

The density of isopropyl alcohol is 0.785 g/mL. The density of water at room temperature may be found from a table in a previous experiment.

Before you begin, check your thermometer to make sure there are no breaks in the liquid column.

  1. Prepare an ice/water/rock salt bath by placing equal amounts of ice and rock salt and a small amount of water in the beaker. Stir the mixture. The temperature of the bath must be at least – 10oC. If it is not, pour off some of the water, add more ice and rock salt, and stir the mixture again. Continue this process until the desired temperature is achieved.
  2. Rinse the freezing point apparatus well with distilled water. Shake out as much water from the test tube as possible. Measure 15.0 mL of distilled water with a graduated cylinder and transfer it completely into the test tube. Stopper the test tube with the stopper holding the thermometer and stirrer. The thermometer scale marks below 0oC should be visible. Clamp the test tube in position in the cold bath so that the contents are below the level of the liquid in the bath.
  3. Move the stirrer up and down constantly to agitate the water, pausing regularly at 30-second intervals to record the time and the temperature. Continue until the temperature readings are constant for five consecutive readings or until the water has frozen sufficiently that stirring has become difficult. Allow the frozen water to thaw and return to room temperature. Repeat this freeze/thaw procedure.
  4. Using a graduated cylinder, measure out 3 to 4 mL of isopropyl alcohol, then pour into the freezing point apparatus which has thawed for the second time.
  5. Stopper the freezing point apparatus and place it in the cold bath (remember to monitor the bath temperature – it must be below –10oC). Repeat the process of agitation, recording the temperature every 30 seconds. Continue until the temperature readings have been constant for five consecutive readings or until the mixture has frozen sufficiently that stirring has become difficult. Allow the frozen mixture to thaw and return to room temperature. Then repeat this freeze/thaw procedure.

If the alcohol/water mixture does not freeze at a constant temperature, it may be that the mixture is supercooled, that is, the liquid does not solidify until it is cooled slightly below its freezing point. The freezing point can be determined graphically, as follows. Plot the temperature vs. time data which should resemble the graph shown below. The freezing point is the intersection of the two straight-line parts of the plot.

 

Temperature

 

Time

 

 

After concluding the procedure, pour off the water from the bath and if any solid rock salt remains, place the ice/rock salt mixture in the designated place for reuse by other classes. The mixture of isopropyl alcohol and water may be poured down the drain. _________________________________________________________________________________________________________

DETERMINATION OF MOLECULAR MASS BY FREEZING POINT DEPRESSION

_________________________________________________________________________________________________________

Show all calculations on a separate page. INCLUDE ALL OF THE PROPER AND NECESSARY UNITS IN YOUR DATA TABLE AND IN YOUR SAMPLE CALCULATIONS.

Volume of distilled water solvent ___________

Mass of distilled water solvent ___________

Time/Temperature readings for pure water

Record data in notebook. You will need to leave space for approximately 15 to 20 readings.

Freezing point of pure water

________

________

Average

________

 

Isopropyl alcohol:

Initial buret reading

___________

 

Final buret reading

___________

 

Volume of isopropyl alcohol

___________

 

Mass of isopropyl alcohol

___________

 

Time/Temperature readings for mixture of isopropyl alcohol and water Record data in notebook. SAME AS ABOVE.

Freezing point of mixture

________

________

Average

________

 

Tf

___________

 

mexp’t

___________

 

Moles isopropyl alcohol in mixture

___________

 

Molecular mass of isopropyl alcohol

___________

 

Percent error

___________

 
DATA ANALYSIS SUPPLEMENT

Data: After you have recorded the time/temperature readings for pure water, initial and final buret readings for isopropyl alcohol, and time/temperature readings for mixture of isopropyl alcohol and water, enter your own data into the appropriate Web-based data entry form, using a computer in the lab. In your measurements and calculations, remember to pay attention to uncertainties and significant figures.

isopropyl alcohol, freezing point depression, molality, moles of isopropyl alcohol in mixture, molecular mass of isopropyl alcohol and percent error.

Your lab report should include the following in addition to the pages from your lab notebook:

An Excel sheet “Raw Data” showing the raw data for the entire class (as downloaded). Make sure that you format it so that units are included.

  • A 2nd Excel worksheet “Results” showing the calculated results for each student in the class – each in a separate row. Make sure to show an example calculation for each column. Each table should be formatted according to the instructions given in the Excel Tutorial on the retrieval page. The columns should include:
  • m_H2O = Mass of distilled water solvent
  • m_Iso = Mass of Isopropyl alcohol
  • dt_f = Freezing point depression
  • m_expt = Molality
  • Mol (Isopropyl) = Moles of isopropyl alcohol in mixture
  • MW (Isopropyl) = Molecular mass of isopropyl alcohol
  • % Error = Percent error
  • An Excel chart ‘Freezing Point Determination’ with the variables time on the x-axis and temperature on the y-axis to determine the freezing point.
  • A 2nd Excel chart ‘ Histogram’ showing a histogram of the distribution of percent error in the experimental molecular masses of isopropyl alcohol. Use equally spaced “bins” chosen to display the “spread” in the class results. Refer to the histogram tutorial on the website.

Discussion: In your discussion section address the following questions:

  • Define the freezing point of a solution?
  • What may be a reason why an alcohol/water mixture does not freeze at a constant temperature?
  • What are the possible sources of error in this experiment? Refer in your answer to the class histogram as well as to your own results.
  • Discuss the shape of the histogram. What does it tell you about the nature of the error in the class’ measurements?

 

 

 

Experiment 23 CHEMICAL KINETICS

 

 

 

 

 

word image 1046 word image 1047 word image 1048 word image 1049

Crystal Violet Dye Reaction Order

  1. Graphical Determination of Reaction Order for Crystal Violet Dye
  2. Determination of the Complete Rate Law
Purpose

For a simple reaction system consisting of an organic dye and aqueous base, determine the order of the reaction with respect to the organic dye. This general method involves monitoring a reaction by visible spectroscopy, and obtaining data sets of percent transmittance versus time. Later processing of these data sets will lead to the unambiguous assignment of reactant order by a graphical method.

Introduction

Kinetics is a large and diverse field of study within chemistry. It is primarily concerned with discovering, and if possible controlling, those factors that govern the rates of reaction. Starting with the relatively simple reaction system A→ B, immediately two ways to determine rate become apparent:

  • Determine the rate at which the reactant, A, is consumed, or
  • Determine the rate at which the product, B, is formed

Such a simple analysis leads to the construction of a general rate expression for this system, mathematically expressing the two general methods outlined above:

word image 1050

This general rate expression can be read in terms of either A or B. For the reactant A: “Rate is equal to minus the change in the concentration of A over the change in time.” For the product B: “Rate is equal to the change in the concentration of B over the change in time.” This treatment allows the reaction rates to have positive values for [A] and [B], and for the two rates to be expressed in terms of one another.

What this general expression really says mathematically, is something that is intuitively obvious physically: that the rate at which product B forms cannot be any faster than the rate at which reactant A is being consumed to produce it. General rate expressions become more complicated, but keep these same features, when a reaction system under study becomes more complex.

Below is a more complicated general reaction equation, and a corresponding general rate expression for that reaction:

𝑎𝐴 + 𝑏𝐵 → 𝑐𝐶 + 𝑑𝐷

word image 1051

The introduction of reaction coefficients brings one additional term to each reactant and product in the general rate expression.

Relevant Background Information:

The more complicated general rate expression above can still be read the same way as the simple one, and has the same utility. Being able to measure the change in concentration as a function of time for any single reactant or product allows the measurement of the rate of change over time for any of the other reactants or products. In this experiment, the disappearance of the color of one of the reactants can be used, if one so desires, to characterize the rates for all chemical species in the reaction system.

Thus the design of this experiment is to monitor, by visible spectroscopy, this disappearance of color, and convert that information into a general expression for how a particular reactant behaves. One can assume the disappearance of color to be related directly to the consumption of the reactant and an appearance of product. You will also determine the order of reaction with respect to this consumed reactant, which becomes part of what is called the rate law for the reaction.

You will be concerned with how the change in concentration versus time can be turned into a reactant order for a given reactant. Usually reaction orders have integer values, or zero. They can also be fractional, and can even be negative. The reaction under study is one of the type A(color) → B(colorless), and the rate can be conveniently expressed in terms of disappearance of color:

word image 1052

 

The reaction order, x, can be found in one of two ways. In the first general method, the rate can be found by varying the concentration of A. For example, a ‘baseline’ concentration of 1 for A, yields a corresponding ‘baseline’ rate, also of 1. If you were to then double the concentration of A to 2, assuming all other factors remain the same, the rate of reaction could again be measured. The dependence of changing the concentration of A on rate could be measured directly, and the order solved mathematically by:

𝑅𝑎𝑡𝑒 = [2]𝑥

If the rate were observed to increase by a factor of two, this expression simplifies to:

2 = [2]𝑥

The reaction order, x, is therefore equal to 1. The reaction can then be said to be, “first-order with respect to A.” Convince yourself that if doubling the concentration of A led to no change in the observed rate, that the reaction order x would be 0. Also, if doubling the concentration of A led to a factor of 4 increase in rate, that the reaction order x would be 2.

In the second general method, the use of integrated rate law expressions is used to determine reaction order. The results for zero-, first-, and second-order reactions are presented in the following table.

 

Zero-Order First-Order Second-Order

Rate Law

k[A]0 k[A]1 k[A]2

Integrated Rate Expression

[A]t = kt + [A]0 ln[A]t = kt + ln[A]0 1/[A]t = kt + 1/[A]0

Linear graph

[A] versus time ln[A] versus time 1/[A] versus time

Slope

k k k

 

Relevant Chemistry and Experimental Details

I. Determination of Reaction Order with respect to Crystal Violet Dye

The reaction of the organic dye crystal violet with aqueous base is a relatively simple one, whose progress can be tracked over time by the disappearance of a purple color:

crystV(aq) + OH(aq) ——–>crystVOH(aq)

(purple) (colorless)

A general rate expression for this reaction can be written:

word image 1053

The complete rate law for this system is:

𝑅𝑎𝑡𝑒 = 𝑘[𝑐𝑟𝑦𝑠𝑡𝑉+]𝑥[𝑂𝐻]𝑦

The disappearance of the purple color, indicating the consumption of the crystal violet dye, can be monitored over time with a Spectronic 20 visible spectrometer. This instrument is able to determine how much of a given incident beam of light is being allowed through a sample at a given wavelength. The more color, the less light passing through. As the reaction proceeds, and the color fades, more of the incident light is allowed through. The experimental procedure asks that readings from the spectrometer be taken in percent transmittance, which should be seen to increase as a function of time. This would be indicative of an initially “rich” color fading over time, thus allowing more of the incident light to be ‘transmitted’ through it. Percent transmittance readings should be converted to absorbance with this equation:

Absorbance = log(%T/100)

Absorbance can be related to concentration (although the exact concentration of the crystal violet dye is neither found nor needed, as is shown shortly) using the Beer-Lambert Law.

Use of the Beer-Lambert Law — Exploiting the Absorbance/Concentration Relationship

The Beer-Lambert Law is a simple relationship between the amount of incident light a given analyte absorbs as a function of the concentration of that analyte:

A = 0 b c

The terms 0 (extinction coefficient) and b (path length) are constant under the conditions of this experiment, allowing the further simplification

Absorbance (is directly proportional to) concentration

Without this simplification, it would be necessary to determine the actual concentrations of crystal violet dye at every time measurement. Since the absorbance is directly proportional to the concentration of crystal violet dye, the data sets collected will actually be of absorbance versus time information. You will assume this proportionality is true, and that your measured absorbances will be treated as if they were actually concentrations of crystal violet dye. It does not change any of the attributes of the integrated rate laws to use absorbances in place of concentrations under the conditions of this experiment.

Additional Treatments of Experimental Data: Complete Rate Law and Activation Energy

II. Determination of the Complete Rate Law

Once the order of reaction with respect to crystal violet dye has been determined, it is possible to find the complete rate law by also determining the order of reaction with respect to aqueous hydroxide, OH. The method used takes advantage of two facts:

  • the rate constant, k, is unchanged if the temperature is unchanged
  • the concentration of crystal violet dye is constant under the experimental conditions The rate law can accordingly be simplified to a form such as this:

original rate law: Rate = k [crystV]x [OH]y modified rate law: Rate = k’ [crystV]x where k’ = k [OH]y

The new symbol k’ assumes the crystal violet dye concentration, and the rate constant, are in fact constant. It should now be possible, by varying the hydroxide ion concentration, to determine the exponent y, the order of reaction with respect to hydroxide. It is a consequence of using the modified rate law that now, the actual rate constant, k, is dependent on hydroxide ion concentration. The slope of the graph prepared that reveals the reaction order with respect to crystal violet dye is in fact this term k’, and not the actual rate constant, k. Therefore it is necessary to determine the reaction order with respect to crystal violet dye before attempting to determine the order with respect to hydroxide ion concentration. This simplification cannot be made and used without knowing the exponent x.

If, therefore, a series of experiments were run in which the hydroxide concentrations were varied, a dependence of this variance on the rate would be obtained directly, and would reveal itself as hydroxide ion-dependent values for k’. If two such experiments with varying hydroxide ion concentrations were run (denoted by “[OH]1y” and “[OH]2y“) the order with respect to hydroxide ion could be found with the following relationships.

word image 1054

which can be simplified to (k1′ / k2′ ) = ([OH]1 / [OH]2)y. This method has the additional benefit if being able to provide a value for the actual rate constant, k, with just a little extra math. It is not necessary to know the rate constant to construct the complete rate law, only the exponents.

Since this method uses actual experimental numbers, and does not rely on a fitted straight line, it is important to point out that while most reaction orders are whole numbers, it is unlikely actual data will lead directly to a whole number result. When considering which order to assign a given reactant by this method, select the nearest whole number for the reaction order. An answer of

1.124, for example, is best rounded back to 1, and so on.

 

 

THE EXPERIMENT

Collection of %T versus Time Data, Constant Temperature

Data sets are to be collected individually, and for two different concentrations, of aqueous sodium hydroxide solution with crystal violet dye. Standard concentrations might include 0.025M NaOH, and 0.0125M NaOH. Be sure to check the reagent bottles in your laboratory section for the exact concentrations available.

Begin by turning on a Spectronic 20 and letting it warm up for at least 15 minutes. Be sure the wavelength indicator is set to 540 nm, and do not adjust it again for the remainder of the experiment. When the instrument has warmed up, use a clean small test tube from your drawer, filled 2/3 with the first NaOH solution you plan to use, as a blank to set 0% and 100% transmittance.

You will be taking a %T reading every 30 seconds for at least 7 minutes, so it may be convenient to wait until the room clock’s second hand approaches 12 before beginning the following procedure. Add one single drop of the crystal violet dye solution, invert to mix the contents of the test tube, and carefully place the tube into the instrument. When the drop hits the surface is your “time zero”, and 30 seconds later you will begin taking readings of %T. Leave the sample tube in the instrument throughout the data collection process. Record %T values to one decimal place, as this allows you to convert to absorbance and retain three significant figures.

When you have taken your 15th and final reading, remove the sample tube from the instrument and prepare to re-zero the instrument with a new NaOH blank in another test tube. Then repeat the entire 7-minute data collection a second time.

Remember, everyone will be required to measure two trials of a full 7-minute data on two different NaOH samples of differing concentration. The real work of this experiment comes in the data analysis, only some of which you will be required to do in the laboratory.

When you and your partner have finished collecting data, turn off the instrument, discard your samples down the drain, and carefully clean the test tubes you used. Enter whatever data is required into the 212 database before you leave for the day.

Results

After consulting the table of reaction orders and their attributes, the order of the crystal violet dye should make itself apparent if, after collecting the absorbance and time data, three graphical plots are prepared for each trial: • One graph of absorbance vs. time

• One graph of ln(absorbance) vs. time

• One graph of 1/absorbance vs. time

Only one type of these plots should be a straight line. This is therefore a direct, graphical determination of the order of reaction with respect to crystal violet dye.

Discussion:

  1. What is the order of the reactant crystal violet?
  2. Compare the values for slope from the linear graphs you obtained. Does there appear to be any trend you can spot when comparing these two slopes and the change in concentration of hydroxide ion? Rates of reaction should be proportional to observed values for k.
  3. In general, which of the two concentrations reached the 90% mark more quickly? State any trend that stands out.
  4. The starting concentration (time zero) of crystal violet is 1 x 10-4 M. Using the proper integrated rate law and the calculated value for k, calculate the concentration of crystal violet after 240 seconds.
  5. Use the y-intercept from your straight line to obtain a value for e0, the extinction coefficient, for crystal violet at time zero. Use the Beer-Lambert Law to do this.
Data Tables

_________________________________________________________________________________________________________

CHEMICAL KINETICS: CRYSTAL VIOLET DYE

_________________________________________________________________________________________________________

Concentration of NaOH Solution __________

Concentration of Crystal Violet Dye Solution __________

Room Temperature __________

(assumed to be reaction temperature)

Trial #1 Trial #2

Time(sec) %T Time(min) %T

30 _______ 30 ________

60 _______ 2 ________

90 __________ 90 __________

120 __________ 120 __________

150 __________ 150 __________

180 __________ 180 __________

210 __________ 210 __________

240 __________ 240 __________

270 __________ 270 __________

300 __________ 300 __________

330 __________ 330 __________

360 __________ 360 __________

390 __________ 390 __________

420 __________ 420 __________ 450 __________ 450 __________

Data Tables continued

Concentration of NaOH Solution __________

Concentration of Crystal Violet Dye Solution __________

Room Temperature __________

(assumed to be reaction temperature)

Trial #1 Trial #2

Time(sec) %T Time(sec) %T

30 __________ 30 __________

60 __________ 60 __________ 90 __________ 90 __________

120 __________ 120 __________

150 __________ 150 __________

180 __________ 180 __________

210 __________ 210 __________

240 __________ 240 __________

270 __________ 270 __________

300 __________ 300 __________

330 __________ 330 __________

360 __________ 360 __________

390 __________ 390 __________

420 __________ 420 __________

450 __________ 450 __________

Experiment 24 MEASUREMENT OF AN EQUILIBRIUM CONSTANT

 

 

 

 

 

word image 1055 word image 1056 word image 1057 word image 1058 word image 1059

The true thermodynamic equilibrium constant, Keq, is formulated in terms of the activities of the chemical species present at equilibrium. Because the activities of some species may be difficult to determine, we define an experimentally accessible equilibrium constant, Kc. In a hypothetical solution behaving as at infinite dilution, the activities are numerically equal to the concentrations, and under these conditions Keq equals Kc. For a balanced chemical equation of the form

pP + qQ rR + sS

word image 1060 the expression for the equilibrium constant is

[R]r [S]s Kc = [P]p [Q]q

where the brackets indicate the molar concentrations of the relevant species.

The magnitude of Kc indicates the position of equilibrium of a chemical system at a given temperature, that is, it indicates the relative concentration of reactants and products when equilibrium is established. A numerical value of the equilibrium constant larger than one indicates that the equilibrium lies to the right and thus favors the products, whereas a numerical value of Kc less than one indicates that the equilibrium lies to the left and thus favors the reactants.

THE EXPERIMENT

In this experiment, the value of an equilibrium constant, Kc, will be determined spectrophotometrically for the reaction between hydrated ferric ion (Fe3+) and thiocyanate ion ((NCS)−1) to form the ferric thiocyanate complex ion, [Fe(NCS)]2+. The overall balanced chemical equation is

[Fe(H2O)6](NO3)3 + Na(NCS) [Fe(H2O)5(NCS)](NO3)2 + H2O + NaNO3 (1)

for which we can write the net equation

[ Fe(H2O)6]3+ + (NCS) −1 [Fe(H2O)5 (NCS)]2+ + H2O

colorless colorless red

which can be further simplified by omitting to write the water ligands

Fe3+ + (NCS)−1 [Fe(NCS)]2+ + H2O

(1a)

(1b)

Because the activity of water in dilute aqueous solutions is equal to 1, we can simplify and write the equilibrium constant as

[Fe(NCS)2+ ]

Kc = [Fe3+ ][(NCS)1-] (2)

When solutions containing Fe3+ and (NCS)−1 are mixed, reaction (1) occurs rapidly to form the deep blood-red colored complex ion, [Fe(NCS)]2+, which absorbs light strongly at a wavelength of 447 nm. Since neither Fe3+ nor (NCS)−1 absorb light at this wavelength, the measured absorbance of the solution can be used to determine the concentration of [Fe(NCS)]2+ in the equilibrium mixture. The concentrations of Fe3+ and (NCS)−1 at equilibrium with [Fe(NCS)]2+ can then be calculated from the known initial concentrations of Fe3+ and (NCS)−1. The value of Kc is obtained from the equilibrium expression (2).

The experiment is carried out in two parts. In Part A, you will prepare a set of standard solutions containing known concentrations of [Fe(NCS)]2+. You will determine the absorbance of each of these solutions using the Spectronic 20 spectrophotometer and establish a concentration calibration by plotting Absorbance versus concentration of [Fe(NCS)]2+.

The standard solutions of known concentration of [Fe(NCS)]2+ are prepared by mixing solutions of (NCS) −1 at low concentrations with solutions containing a very much larger concentration of Fe3+. Under these conditions, the equilibrium in equation (1) is shifted far to the right, converting almost all of the (NCS)−1 into [Fe(NCS)]2+. We can therefore conclude that the concentration of the [Fe(NCS)]2+ is approximately equal to the initial (NCS)−1 concentration.

***You should review the discussion in the “Absorption Spectroscopy” experiment on establishing a plot which relates the absorbance and concentration of an absorbing species.

In Part B you will use this calibration plot to determine the concentration of [Fe(NCS)]2+ which is in equilibrium with Fe3+ and (NCS)−1 in several solutions. In these solutions, neither of the reactants is in large excess and all the concentrations can be calculated. After calculating the equilibrium concentrations of all three ions you will determine the equilibrium constant, Kc.

CALCULATIONS AND PRE-LABORATORY ASSIGNMENT

As you do the calculations here and the ones for the experiment, keep the following points in mind.

  1. You will be making and using solutions of known concentration (M) and known volume (mL). The concentrations of the solutions you make are calculated using a concept already familiar to you from previous experiments:

V1  M1 = V2  M2

  1. At equilibrium, the Fe is present in two forms – as part of the [Fe(NCS)]2+ complex and as unreacted Fe3+. The amount of Fe present at equilibrium must be equal to the amount of Fe present at the beginning;.

At equilibrium, the NCS is present in two forms – as part of the [Fe(NCS)]2+ complex and as unreacted NCS-1. The amount of NCS present at equilibrium must be equal to the amount of NCS present at the beginning;

  1. If one of the reactants is present in great excess, essentially all of the other reactant (the limiting reagent) reacts to form product.

 

Fe

3+

 

 

+

(NCS

)

1

 

 

 

)]

[

Fe(NCS

2+

 

 

at beginning

 

Calculate from

 

V

1

 

M

1

 

V

=

2

 

M

2

 

 

 

 

0

 

amount reacted

 

Reacts 1:1 to form 1 product

 

 

at equilibrium

 

Beginning

 

amount reacted

 

 

D

etermined from graph

 

 

Example:

5.0 mL of a 0.0200 M Fe(NO3)3 solution is mixed with 5.0 mL of 0.00200-M (NCS)−1. Assume the volumes are additive. The blood-red [Fe(NCS)]2+ ion that forms has an equilibrium molar concentration of 7.00 x 10-4 mol/liter as determined from a calibration plot of its absorbance vs. concentration. Calculate the Kc for [Fe(NCS)]2+ formation.

Fe(NO3)3 + Na(NCS) Fe(NCS)(NO3)2 + Na(NO3)

  1. Fe3+ concentration at beginning

M2 = (V1 M1 )/ V2 = (5.0 mL 0.0200 M) / 10.0 mL = 0.0100 M

  1. (NCS)−1 concentration at beginning

M2 = (V1 M1 )/ V2 = (5.0 mL 0.00200 M) / 10.0 mL = 0.00100 M

  1. [Fe(NCS)]2+ concentration at equilibrium

From a calibration plot = 7.00 x 10-4 M

  1. concentration of Fe3+ reacted

Because the Fe3+ which has reacted is present as [Fe(NCS)]2+, their concentrations are identical = 7.00 x 10-4 M.

  1. concentration of (NCS)−1 reacted

Because the (NCS)−1 which has reacted is present as [Fe(NCS)]2+, their concentrations are identical = 7.00 x 10-4 M.

  1. concentration of Fe3+ remaining unreacted at equilibrium

The concentration unreacted = concentration at beginning – concentration reacted = (1.00 x 10−2) − (7.00 x 10−4) = 9.3 x 10−3 M

  1. concentration of (NCS)−1 remaining unreacted at equilibrium

The concentration unreacted = concentration at beginning – concentration reacted =

(1.00 x 10−3) − (7.00 x 10−4) = 3.0 x 10−4 M

j. Kc = [Fe(NCS)2+] / [Fe3+][NCS−1]

Kc = (c.) / (f.)(g.) = 251

 

 

 

 

 

 

The calculations below are to be completed before the laboratory period. Show your calculations at the bottom of the page.

7.0 mL of a 0.0026-M Fe(NO3)3 solution is mixed with 5.0 mL of 0.0066-M (NCS)−1. The blood-red

[Fe(NCS)]2+ ion that forms has an equilibrium molar concentration of 5.3 x 10-4 mol/liter as determined from a calibration plot of its absorbance vs. concentration. Calculate the Kc for

[Fe(NCS)]2+ formation.

 

 

a. Fe3+ concentration at beginning

 

M

b. NCS−1, concentration at beginning

 

M

c. [Fe(NCS)2+]concentration at equilibrium

 

M

d. concentration of Fe3+ reacted

 

M

e. concentration of NCS−1 reacted

 

M

f. concentration of Fe3+ remaining unreacted at equilibrium

 

M

g. concentration of NCS−1 remaining unreacted at equilibrium

 

M

j. Kc

__________

 

PROCEDURE

Part A Preparation of Calibration Plot—Standard Solutions Prepared by Preps Room

  1. Clean a buret and rinse it with two 10-mL portions of 0.0020 M Na(NCS).
  2. Fill the buret with fresh solution and deliver 0, 3, 5, 7, and 10 mL of 0.0020-M Na(NCS) into 5 separate, labeled 100-mL volumetric flasks. If you do not deliver the volume specified, then record exactly the volume you delivered (e.g. 4.05 mL).
  3. Pipet 25.0 mL of 0.200-M Fe(NO3)3 into each volumetric flask, swirl to mix, and dilute carefully to the 100-mL mark with 0.1-M HNO3.
  4. Place the cap on the flask and invert it several times to mix each solution thoroughly. These standard solutions are used to establish a calibration curve. (Do not use a graduated cylinder to measure the volume.)
  5. After the Spectronic 20 instrument has been turned on for 20 minutes, set the absorption wavelength at 450 nm and, with no cuvette in the instrument, set the % Transmittance to 0% (Absorbance = ) using the left-hand knob.
  6. Rinse a cuvette with several portions of Solution 1 and then partially fill it with Solution 1. Carefully dry the outside of the cuvette with a clean tissue to remove moisture and fingerprints. Handle the cuvette only by its top edges. Solution 1, the reference solution, is used to calibrate the spectrophotometer.
  7. Calibrate the spectrometer by adjusting the right-hand knob to 100% Transmittance (Absorption = 0) with the reference solution in the light path. (Keep this reference solution cuvette for Part B.)
  8. Rinse another cuvette with several portions of Solution 2. Partially fill it with fresh solution. Dry the outside of the cuvette with a clean tissue. Record the solution’s absorbance at 450 nm.
  9. Repeat with solutions 3, 4, and 5.
  10. Record the concentration (M) of [Fe(NCS)2+] in each of the standard solutions (1-5), remembering that the Na(NCS) is the limiting reagent. Plot the Absorbance (y-axis) versus concentration of [Fe(NCS)2+] (x-axis) for the five solutions. Draw the best straight line through your points.

 

Composition of Solutions for Calibration Curve

Solution Volume of Volume of

0.200 M Fe(NO3)3 0.00200 M Na(NCS) (in 0.1 M HNO3) (in 0.1 M HNO3)

_______________________________________________________________________________________________________

  1. 25.0 mL 0 mL
  2. 25.0 mL ~3 mL
  3. 25.0 mL ~5 mL
  4. 25.0 mL ~7 mL
  5. 25.0 mL ~10 mL
Part B Determination of Equilibrium Constant
  1. Prepare the solutions below in 10-mL volumetric flasks. Use a graduated cylinder to measure out the indicated volumes of Na(NCS) and Fe(NO3)3 solutions. . (Note the difference in concentration of the ferric nitrate from Part A. Do not mix up the two different solutions.) Dilute to the mark with HNO3 and mix the contents of the flask.

 

 

Test Solutions

Fe(NO3)3

NaNCS

HNO3

Total

Volume

(mL)

0.002 M

0.002 M

0.100 M

(mL)

(mL)

(mL)

1

1.0

1.0

8.0

10.0

2

1.0

1.5

7.5

10.0

3

1.0

2.0

7.0

10.0

4

1.0

2.5

6.5

10.0

5

1.0

3.0

6.0

10.0

6

2.0

1.0

7.0

10.0

7

2.0

1.5

6.5

10.0

8

2.0

2.0

6.0

10.0

9

2.0

2.5

5.5

10.0

10

2.0

3.0

5.0

10.0

 

  1. Use the reference solution from Part A to calibrate the spectrophotometer.
  2. Mix each test solution until equilibrium is reached (approximately 1 minute). Rinse a cuvette with several portions of the test solution and partially fill it. Clean and dry the outside of the cuvette. Record absorptions as in Part A.
  3. Using the calibration plot from Part A, determine the [Fe(NCS)]2+ concentration present in each equilibrium mixture by drawing a dotted horizontal line from the Absorbance axis to its intersection with the best straight line. Drop a vertical dotted line to the concentration axis — this value is the concentration of the solution.
  4. From the equilibrium [Fe(NCS)]2+ concentration and the initial Fe3+ and (NCS)−1 concentrations used in each reaction solution, calculate the equilibrium Fe3+ and (NCS)−1 concentrations. Hence determine Kc for each solution using equation (2). _________________________________________________________________________________________________________

MEASUREMENT OF AN EQUILIBRIUM CONSTANT

_________________________________________________________________________________________________________

A. Calibration Plot

Concn. of Fe(NO3)3 M Concn. of Na(NCS) M

_________________________________________________________________________________________________________

Solution 1 2 3 4 5

_________________________________________________________________________________________________________

Concn _______ M _______ M _______ M _______ M _______ M

Absorbance at 450 nm ________ ________ ________ ________ ________

 

B. Equilibrium Solutions

Conc. of Fe(NO3)3 ________ M Conc. of NaNCS ________ M

____________________________________________________________________________________________________________ Solution: 1 2 3

____________________________________________________________________________________________________________

 

Initial Vol. Na(NCS)

 

 

_______ mL

_______ mL

_______ mL

Initial Conc. (NCS)−1 in 10-mL flask

 

 

_______ M

_______ M

_______ M

Initial Vol. Fe(NO3)3

 

 

_______ mL

_______ mL

_______ mL

Initial Conc. Fe3+ in 10-mL flask

 

 

 

_______ M

_______ M

_______ M

Absorbance at 450 nm

 

 

_______

_______

_______

Equil. Conc. [Fe(NCS)2+]

 

 

_______ M

_______ M

_______ M

Equil. Conc. Fe3+

 

 

_______ M

_______ M

_______ M

Equil. Conc. NCS−1 (M)

 

 

_______ M

_______ M

_______ M

Kc

 

 

_______

_______

_______

Average Value of Kc

 

 

_______

 

Standard Deviation:

 

 

_______

 

Show a typical calculation for each of the above sets of calculations.

ATTACH CALIBRATION PLOT

 

DATA ANALYSIS SUPPLEMENT

Data: After you have noted down all volumes, absorbances and molarities, enter your own data into the appropriate Web-based data entry form, using a computer in the lab. In your measurements and calculations, remember to pay attention to uncertainties and significant figures.

Results: Outside of class, retrieve your data in spreadsheet form (along with that of the entire class) from: http://chem.gmu.edu/results/.

Your lab report should include a printout worksheets labeled: RawData, ResultsA, MyResults ResultsB1, ResultsB2, MyGraph in addition to the pages from your lab notebook. Each table should be properly formatted with lines by including lines above and below the raw to separate the data from the rest of the worksheet as well as a title for the worksheet and titles for each of the columns. Look at the downloadable sample worksheets at the above URL for examples of the correct formatting. Be sure to include titles for each column of results. Include a sample of each calculation on the spreadsheet at the bottom left hand side of each worksheet.

  • RawData: This worksheet should show the raw data for the entire class (as downloaded).
  • ResultsA: Include the results from part A in this worksheet. Include columns for o Each of the five Fe(NCS)2+ concentrations.
    • Linear least squares calculation of the slope of a plot of this concentration and the absorbances (AA1..AA5) from the RawData table. Use the “=slope(y1:y5, x1:x5)” function.
    • Linear least squares calculation of the intercept of a plot of this concentration and the absorbances (AA1..AA5) from the RawData table. Use the “=intercept(y1:y5, x1:x5)” function.
  • MyResults: List your Fe(NCS)2+ concentrations and absorbances for part A in columns. This data will be used to make your graph.
  • MyGraph: Make a graph in this worksheet of the data in MyResults. Remember to follow standard rules for making graphs: the graph should be in landscape orientation on a page by itself….covering at least 80% of the page (borders included); use a clear title for the graph; label axes properly and include units. Do not include a legend, since you have only one set of data.
  • ResultsB1: Using volumes of iron and NCS as well as the concentrations of the stock solutions from the RawData page, determine initial concentrations (before reaction occurs) of each of these for the solutions prepared. Include columns for
    • Fe3+ concentration for each of the three solutions.
    • NCS concentration for each of the three solutions.
  • ResultsB2: Using results from linear least squares analysis in ResultsA and the absorbances of the solutions for part B of the experiment (RawData columns labeled AB1. AB3), determine the equilibrium concentrations of the complex in each solution. From this and the concentrations on ResultsB1, determine the equilibrium concentrations of the uncomplexed ions (Fe3+ and NCS). From these concentrations determine the equilibrium constant for each solution. Columns to include on this worksheet:
    • Equilibrium concentration of complex for each solution.
    • Equilibrium concentration of Fe3+ in each solution. o Equilibrium concentration of NCS in each solution.o Equilibrium constant for each solution.o Average of the equilibrium constant.o Standard deviation of equilibrium constant.o Class average of the equilibrium constant.
    • Class standard deviation of the equilibrium constant.

Discussion: In your discussion section address the following questions in addition to points discussed in the book:

  • How does your equilibrium compare with the rest of the class?
  • Is there much class variation in the results?
  • What does the calculated equilibrium constant tell you about the equilibrium position of the reaction you studied? Does the reaction favor reactants or products?
  • Compared with the class is your mean and standard deviation similar? Why or why not?

 

 

Experiment 25 ACIDITY CONSTANT FROM A TITRATION CURVE

 

 

 

 

word image 1061 word image 1062 word image 1063 word image 1064 word image 1065

In a previous experiment, you learned how to calculate the Ka of a weak acid solution in two ways, both of which involved determining the concentrations of the species present and substituting into the equation for the acid dissociation constant

Ka word image 1066 (1)

There is yet another way to calculate the Ka of a weak acid which depends upon a useful relationship, the Henderson-Hasselbalch equation, which can be derived as follows. Taking the logarithm of both sides of equation (1) gives

word image 1067 logKa  log( word image 1068 ) (2)

Remembering that log(xy) = log(x) + log(y)

logK a  log[H]log word image 1069 (3)

By definition, −log [Ka] = pKa and −log [H+] = pH. Multiplying all terms by −1 and substituting gives the Henderson-Hasselbalch equation

[A]

pKa  pH  log

[HA]

(4)

When half of the weak acid has been neutralized by a base, the concentration of HA equals the concentration of A, i.e., [HA] = [A]. In this case, since [A]/[HA] = 1, substituting into equation (4) leads to log([A]/[HA]) = 0 and pH = pKa.

The Ka of a weak acid may be calculated by first titrating the weak acid with a base and then by constructing a titration curve in which the continuously changing pH of the titrated solution is plotted against the volume of titrant (the base). The pH is monitored with a pH meter. See the Figure on the next page.

The equivalence point is the pH at which the base has completely neutralized the acid and the moles of acid initially present equal the moles of base added (for monoprotic acids). The equivalence point on the titration curve is easily seen and so a color-changing indicator such as phenolphthalein is not necessary.

The half-neutralization point is the pH at which half the volume of base required for complete neutralization has been added. At half-neutralization, as discussed for the HendersonHasselbalch equation above, the concentration of HA equals the concentration of A, and the pH equals the pKa. The pKa of the weak acid can therefore be read directly from the pH titration curve.

pH

 

Volume of base added (mL)

 

word image 1070

equivalence point

 

pK

a

 

 

 

THE EXPERIMENT

 

An acid-base titration curve shows the changes in pH that occur during the course of an acid-base titration. In this experiment you will construct titration curves for the titration of a strong acid (HCl) with a strong base (NaOH) and for the titration of a weak acid (CH3CO2H, acetic acid) with a strong base (NaOH).

You will use a pH meter to monitor the pH of the solution as a function of the volume of base added to it. The titration curve is obtained by plotting the pH on the vertical axis (y) against the volume of base added on the horizontal (x) axis. You should read the sections in your lecture text concerning buffers and pH titration curves.

PROCEDURE
Part A Titration of a Strong Acid with a Strong Base
  1. Rinse a buret with aqueous NaOH. Using a graduated cylinder, measure out 15 mL of NaOH and pour into the buret.
  2. Calibrate the pH meter. Follow the directions provided and copy them into your lab notebook for future use.
  3. Deliver about 10 mL of the HCl solution (measured exactly) into a 125-mL Erlenmeyer flask and add a few drops of phenolphthalein indicator. Measure the pH of the acid, making sure that the electrodes are immersed in the solution.
  4. Titrate the acid in the beaker by adding the NaOH from the second buret in 1.0 mL increments, reading the pH after each addition. Remember to mix well after adding NaOH base to the solution in the beaker. When about 20 mL of the base have been added, add smaller increments of 0.50 mL, measuring the pH after each increment until the pH reaches ~10, and then continue 2.0 mL additions until pH 12. Take note of the pH and volume of NaOH added when the phenolphthalein indicator turns from colorless to pink.
Part B Titration of a Weak Acid with a Strong Base
  1. Empty the buret containing the HCl, rinse it well with water and then rinse it with the weak acid, acetic acid (~0.1 M). Fill the buret with the acetic acid.
  2. Titrate the acetic acid with the NaOH using the same procedure as in Part A.
Calculation and Results
  1. For each of the acids, HCl and CH3CO2H, calculate the volume of NaOH which should be added to your acid solution to exactly neutralize it. Use the relationship

VNaOHMNaOH = VacidMacid

  1. On separate graphs, plot the pH of the solution against the added volume of NaOH for Parts A and B. Because you have so much data, it is sufficient to plot the results for about every 6 mLs of NaOH added, except near the equivalence point where the pH changes drastically.
  2. On each titration curve, mark the equivalence point and circle it. Draw dotted perpendicular lines to the x and y axes. What is the pH and the volume of NaOH added (VmL) at the equivalence point?
  3. For the titrations of HCl and acetic acid separately, compare the volume of NaOH required to reach the equivalence point from: the calculation in (1.) above, the graph, and the observed phenolphthalein endpoint.
  4. For titration of acetic acid only, calculate the volume corresponding to half-neutralization (VmL/2). On the graph, draw a dotted line perpendicular to VmL/2 until it intersects the titration curve. Mark and circle the point and then draw a horizontal dotted line which intersects the pH axis. This is the pKa of acetic acid.
  5. Compare the pKa you obtained from the titration curve with the pKa’s you obtained in the previous experiment.
  6. Explain why the weak acid-strong base equivalence point is not at pH = 7.0.

Your instructor may ask you to turn in graphs printed from a computer spreadsheet application such as Excel. If so, here are some guidelines for submitting these along with your data and calculations in your notebook.

Enter all of the volume and pH data for the experiment in worksheet columns which are clearly labeled. The two graphs of the data should be created as separate charts, not as embedded charts in the worksheet. Choose an “x-y scatter” chart and choose a chart option which will draw a curve through your data points. Mark the graphs as specified in the Procedure, using drawing objects and text boxes. Include titles and axis labels on the graphs. Print the graphs in a Landscape orientation.. Attach the printouts to your lab report. _________________________________________________________________________________________________________

ACIDITY CONSTANT FROM A TITRATION CURVE

_________________________________________________________________________________________________________ Part A

Volume HCl in Erlenmeyer Flask ____________mL Conc. of HCl ______________

Volume of NaOH added ___________mL pH of solution ________________

 

*Leave space for multiple recordings

Part B

Volume Acetic Acid in Erlenmeyer Flask ____________mL Conc. of Acetic Acid ______________

Volume of NaOH added _____________mL pH of solution _______________

 

 

*Leave space for multiple recordings

CALCULATIONS (show all calculations)

From the phenolphthalein endpoint:

HCl CH3COOH

VmL at equivalence point ________ VmL at equivalence point ________ From the calculation:

VmL at equivalence point ________ VmL at equivalence point ________ From the titration curve:

VmL at equivalence point ________ VmL at equivalence point ________ pH at equivalence point ________ pH at equivalence point ________

VmL/2 for acetic acid ________ pKa of acetic acid ________

VmL/2 for acetic acid ________ pKa of acetic acid ________ DATA ANALYSIS SUPPLEMENT

Data: Record the molarities of standard solutions. During titration note down the volumes and pHs in a tabular form in the laboratory notebook. In your measurements and calculations, remember to pay attention to uncertainties and significant figures.

Results: From the phenolphthalein endpoint, from calculation and from the titration curve, estimate the pH and Veq at equivalence points for HCl and CH3COOH. Also, calculate Veq /2 (half equivalence point) and pKa of acetic acid. (The lab text refers to Veq as V mL)

*There is no download for this experiment.

Your lab report should include the following in addition to the pages from your lab notebook:

  • An Excel sheet “Raw Data” showing the volumes of NaOH vs HCl and NaOH vs CH3COOH.

Enter the numbers in the first spreadsheet from your lab notebook.

  • A 2nd Excel sheet “Graph-HCl” showing the titration curve of HCl titration (solution pH vs volume of NaOH). On the titration curve, mark the equivalence point and circle it. Draw dotted perpendicular lines to the x and y axes. Estimate the pH and volume of NaOH added at the equivalence point.
  • A 3rd Excel sheet “Graph – CH3COOH” showing the titration curve of CH3COOH titration (solution pH vs volume of NaOH). On the titration curve, mark the equivalence point and circle it. Draw dotted perpendicular lines to the x and y axes. Estimate the pH and volume of NaOH added at the equivalence point. Also calculate the volume corresponding to halfneutralization (Veq /2). On the graph, draw a dotted line perpendicular to Veq /2 until it intersects the titration curve. Mark and circle the point and then draw a horizontal dotted line which intersects the pH axis. This is the pKa of acetic acid.

Discussion: In your discussion section address the following questions:

  1. Why is it not necessary to use phenolphthalein indicator to detect the endpoint in this titration?
  2. Can the pKa of acetic acid be calculated from the pH at a volume other than the half equivalence point? Explain how.
  3. What is the pH and the volume of NaOH added (VmL) at the equivalence point? Calculate the pH at equivalence point using your experiment Ka, concentrations and volumes, and compare it with your experiment pH. Find the percent error. What are some sources of error that would account for this difference? BE QUANTITATIVE.
  4. Explain why the weak acid-strong base titration equivalence point is not at pH= 7.0.
  5. What are the possible sources of error in this experiment?

 

Experiment 26 DISSOCIATION CONSTANTS OF ACIDS AND BASES

 

 

 

 

word image 1071 word image 1072 word image 1073 word image 1074 word image 1075

Acid-Base Equilibria

There are three useful ways of thinking of acids and bases, each of which attempts to explain their properties: the Arrhenius, the Bronsted-Lowry, and the Lewis theories. We will consider here only the Bronsted-Lowry Theory which defines acids and bases in terms of proton transfer reactions. (A proton is the hydrogen ion, H+.) According to Bronsted and Lowry:

an acid is any substance (neutral molecule, anion or cation) which donates a proton in an acid-base reaction.

a base is any substance (a neutral molecule, anion or cation) which accepts a proton in an acid-base reaction.

word image 1076 A Bronsted-Lowry acid might be represented by HA and a Bronsted-Lowry base by :B. When these two species react with each other, the acid donates its H+ to :B, which by accepting it, acts as the base.

HA

Acid

+

B

Base

HB+

Conjugate

Acid

+

A-

Conjugate

Base

 

The species which remains after an acid has transferred its proton is called the conjugate base of that acid. Likewise, the species which is produced after a base has accepted a proton is called the conjugate acid of that base. In the example above, HA and :A are a conjugate acid/base pair, as are :B and HB+.

Because aqueous solutions are both common and important, we will consider what happens when an acid or a base is added to water. In the following examples, water is the solvent, as well as acting either as an acid or a base.

HA

+

H2O

H3O+

+

A-

 

B-

+

H2O

HB

+

OH-

 

Substances which can be either acids or bases, such as water, are called amphiprotic.

The specific examples below show HCl, H2PO4, and NH4+ behaving as Bronsted-Lowry acids by donating protons to the solvent water to produce the hydronium ion, H3O+, and their conjugate

bases Cl, HPO42−, and NH3, respectively.

    

Acid

 

Base

 

 

Acid

 

Base

 

 

 

 

 

 

 

 

HCl

 

H2PO4

 

NH4+

+

 

+

 

+

H2O

 

H2O

 

H2O

 

 

 

 

 

H3O+

 

H3O+

 

H3O+

+

 

+

 

+

Cl

 

HPO42−

 

NH3

Similarly ammonia, :NH3, and the phosphate ion, PO43-, are bases in water, according to Bronsted-Lowry, because they accept protons from solvent water molecules (functioning as a Bronsted-Lowry acid in this case) to produce the ammonium ion, NH4+, and the hydrogen phosphate ion, HPO42-, their conjugate acids.

Base

 

Acid

 

 

Acid

 

Base

 

 

 

 

 

 

 

 

:NH3

 

PO43−

+

 

+

H2O

 

H2O

 

 

 

NH4+

 

HPO42−

+

 

+

OH

 

OH

Acid-Base Dissociation Constants

Both acid-base conjugate pairs (as in the examples above) may exist in an aqueous solution in a state of dynamic equilibrium. The quantity of each depends upon the relative strength of the acids and bases present. Acids vary in their ability to donate protons in water solution. The same is true for bases. Acids or bases are classified as strong or weak depending on the degree to which they transfer protons. In terms of equilibria, strength is a measure of the relative extent of the forward and reverse reactions; the stronger the acid or base, the larger is the Keq for the acid-base

equilibrium.

[: A][H O] For an acid in water: Keq  3

[HA][H2O]

[HB][OH]

For a base in water: Keq 

[: B ][H2O]

In a dilute aqueous solution, the concentration of water is essentially constant because it is present in such great excess. Also, because H3O+ and H+ are merely different symbols for a proton in aqueous solution, [H3O+] = [H+]. For a monoprotic acid, the acid dissociation constant, Ka, can be defined for the simplified acid dissociation equation. Likewise, the base dissociation constant Kb can be defined for the base dissociation equation.

word image 1077 [HB][OH]

Ka  Kb  [: B] (1)

The greater the degree of dissociation, the stronger the acid or base and the larger the value of Ka or Kb.

Water as an Equilibrium System

Pure water ionizes to a small extent: H2O(l) H+(aq) + OH(aq). Although only one water molecule is shown in the simplified equation here, each ion is surrounded by polar water molecules and is said to be hydrated or solvated. This is often emphasized by writing the hydrogen ion as H3O+, the hydronium ion. In pure water at 25°C, a state of equilibrium exists such that

[H+] = [OH] = 1.0 x 10−7 M

and the equilibrium constant expression for water, Kw = [H+] [OH], equals 1.0 x 10-14M2.

If H+ ions are added to water by the proton dissociation of an acid, the position of equilibrium must change to maintain the equilibrium constant of water. The number of water molecules will increase as the concentration of OH ions decreases by their reaction with the excess protons. The overall result is an increase in [H+] (>10−7 M) and a decrease in [OH] (<10−7 M). But because Kw must remain constant, [H+] [OH] = 1 x 10−14 M2 although no longer does the equality [H+] = [OH] hold.

pH and Hydrogen Ion Concentration

There are two broad categories of acids and bases which refer to their relative strengths, that is, the extent to which they are dissociated in solution. Strong acids and bases exist completely as their conjugate species in dilute aqueous solutions. In contrast, solutions of weak acids and bases contain measurable concentrations of their undissociated molecular forms.

Acidity, or acid strength, is a function of [H+]. Because the range of concentration of H+ in aqueous solutions is enormous, and because numbers involving powers of ten are clumsy to handle, a logarithmic system (pH) is used to describe H+ and OH concentrations.

The logarithm of the hydrogen ion concentration gives a convenient linear scale of acidity.

pH = -log[H+] or log (1/ [H+])

If pH < 7, the solution is acidic. If pH = 7, the solution is neutral. If pH > 7, the solution is basic (alkaline). Commonly encountered values of pH in aqueous solution range from 0 to 14. In pure water the hydrogen ion concentration is 10−7 (as is the hydroxide ion concentration), and the pH is

7.

The pH can be experimentally determined with a pH meter. Using this instrument, the [H+] in a given solution is compared with the [H+] in a standard solution from which it is separated by a glass membrane. Hydrogen ions in solution, like other ionic species, conduct an electric current. A voltage difference develops across the membrane separating the two solutions of different hydrogen ion concentrations. The voltage is transformed and displayed on the meter as a number between 0 and 14, the pH.

Equilibrium Constants and Problem Solving

If data are available from which the concentrations of the species in the equilibrium system can be determined, then the acidity or basicity constant (dissociation constant) can be calculated. Once the dissociation constant for a given equilibrium system is known, the constant may be used in calculations involving different concentrations. The percent dissociation can also be calculated.

For example, consider a generalized aqueous system containing an acid, HA, and let the initial concentration (concentration before any dissociation reaction takes place) of HA = Ci. Then the initial concentration of A = 0 (and we neglect the dissociation of water) and [H+] also = 0. After reaching equilibrium, an unknown amount, x, of HA will have dissociated. From the 1:1 stoichiometry of the reaction equation, if x amount of HA dissociates, then x amount of both H+ and A will form.

 

 

initial concentrations

 

equilibrium concentrations

 

HA

 

Ci

 

Cix

 

 

 

 

 

 

 

H+

 

0

 

x

 

+

 

 

 

 

 

A

 

0

 

x

 

 

Substituting the equilibrium concentrations of each species into the dissociation constant equation:

[x][x] x2

Ka   (2)

[Ci x] (Ci x)

If Ka and Ci are known, the quadratic equation can be solved for x precisely or approximately. Approximate calculations are justified if x is small compared to Ci. (See your textbook for examples.) This equation is also useful for calculating Ka when Ci and x are known, as in the experiment to follow. Measuring the pH allows a direct determination of x because

x = [H+] = antilog (−pH)= 10−pH (3)

Salt Hydrolysis

 

Many salts are strong electrolytes and exist as ions in aqueous solution. Although most salts contain neither H+ nor OH, they do contain other cations and anions that may produce these species in aqueous solution, a process termed hydrolysis. Thus, the salt solution is acidic or basic depending on whether the ions donate or accept a proton from water.

For the polyprotic phosphoric acid, H3PO4, dissociation occurs in three steps, with

decreasingly smaller dissociation constants.

    

 

    

H3PO4

 

H2PO4−

 

HPO42−

+

 

+

 

+

H2O

 

H2O

 

H2O

 

 

 

 

 

H2PO4−

 

HPO42−

 

PO43−

+

 

+

 

+

H3O+

 

H3O+

 

H3O+

Ka1 = 7.5 x 10−3

 

Ka2 = 6.2 x 10−8

 

Ka3 = 4.8 x 10−13

 

When the phosphate salts Na(H2PO4) and Na2(HPO4) are dissolved in water, they generate the H2PO4 and HPO42- ions in solution which have the ability to function as both proton donors and proton acceptors through the hydrolysis reactions shown below.

H2PO4−

 

H2PO4−

 

HPO42−

 

HPO42−

+

 

+

 

+

 

+

H2O

 

H2O

 

H2O

 

H2O

 

 

 

 

 

 

 

HPO42−

 

H3PO4

 

PO43−

 

H2PO4−

+

 

+

 

+

 

+

H3O+

 

OH

 

H3O+

 

OH

Ka2 = 6.2 x 10−8

 

Kb = 1.3 x 10−12

 

Ka3 = 4.8 x 10−13

 

Kb = 1.6 x 10−7

In each case, the stronger of the two possible dissociations will dominate and determine the overall pH of the solution. For the hydrolysis of a weak acid salt, [H+] can be calculated from equation (2). Provided the dissociation is very weak and thus Ci >> [H+]:

[H]  Ci  Ka

For the hydrolysis of a weak base, B, [OH] can be calculated from equation (1B) and a known Kb and initial base concentration, again provided that the dissociation is very weak:

[OH]  Ci  Kb

THE EXPERIMENT

The preceding paragraphs are the highlights of acid-base chemistry which you need to know for today’s experiment. For a complete introduction, supplement these with appropriate readings and examples from your textbook.

In Part A of this experiment you will make two separate determinations of the acid dissociation constant, Ka, of acetic acid, a weak acid. First, you will measure the pH of an aqueous solution of acetic acid of known concentration CHA, to calculate Ka for the equilibrium dissociation reaction

CH3CO2H H+ + CH3CO2

using equations (2) and (3).

Second, the acetic acid will be partially neutralized by addition of a known quantity of strong base, NaOH, to yield a mixture of unreacted acetic acid, and its conjugate base, acetate ion, formed in the neutralization reaction

CH3CO2H +OH H2O + CH3CO2

Because NaOH is such a strong base, the reaction proceeds essentially completely as written and thus the concentrations of acetic acid and acetate ion can be calculated from the known concentration of NaOH. If y moles of NaOH react completely with excess acetic acid, zy moles of acetic acid will remain and y moles of acetate ion will be produced.

After determining the pH of the solution after mixing, Ka can be determined by equations (2) and (3). In Part B of the experiment, you will both calculate and measure the pH of the phosphate salts NaH2PO4, Na2HPO4 and Na3PO4 derived from the polyprotic phosphoric acid, H3PO4 .

 

 

PROCEDURE

Note: You will need a calculator with a log function.

The pH Meter–– Calibrate the pH meter according the instructions given.

Part A

  1. Using a graduated cylinder, measure approximately 20 mL of the acetic acid into a clean, dry, 125-mL Erlenmeye flask. Determine the pH of the solution at least twice. Do not remove or rinse the electrode before proceeding to the next step. Calculate the Ka of the acetic acid solution.
  2. With the pH meter on standby, carefully deliver about 8 mL of the 0.1 M NaOH into the flask. Gently swirl the solution in the beaker for about one minute, to completely mix the solutions. Again measure the pH of the solution at least twice. Calculate the concentrations of acetic acid and acetate ion present after partial neutralization with NaOH. Calculate the Ka of the solution.

Part B

  1. Review the section on Salt Hydrolysis. By comparing the magnitudes of the acid and base dissociation constants for H2PO4 and HPO42−, decide which is the dominant behavior, acidity or basicity, for each of the salts. Calculate the pH expected for a 0.10 M solution of NaH2PO4 and Na2HPO4 respectively.
  2. Dissolve about 0.6 g of the salts NaH2PO4 and Na2HPO4.7H2O in 30 mL of distilled water in two small beakers. Standardize the pH meter and then measure the pH of each. Compare the pH values with those calculated above.
  3. Measure the pH of a solution of Na3PO4.12H2O (0.2g in 10mL of distilled water).

For the hydrolysis of phosphate ion, Kb is fairly large,

PO43− + H2O HPO42− + OH Kb = 2.1 x 10−2

and the solution is expected to be strongly basic (pH = 12). Compare this with the value measured for Na3PO4. _________________________________________________________________________________________________________

DISSOCIATION CONSTANTS OF ACIDS AND BASES

_________________________________________________________________________________________________________ Show all calculations.

Part A

Initial volume

__________ mL

  

Final volume

__________ mL

  

Volume acetic acid

__________ mL

Acetic acid solution

__________ M

pH __________

__________

Mean pH

__________

[H+] _______

[CH3CO2] _______

[CH3CO2H] _______

Ka __________

 

Initial volume __________ mL

Final volume __________ mL

 

Volume NaOH __________ mL

NaOH solution __________ M

Total volume of solution after mixing

__________ mL

pH __________ __________

Mean pH __________

NaOH reacted __________ mol

acetate ion formed __________ mol

acetic acid present initially __________ mol

acetic acid unreacted __________ mol

[H+] _______ [CH3CO2] _______

[CH3CO2H] _______ Ka __________

 

_________________________________________________________________________________________________________

DISSOCIATION CONSTANTS OF ACIDS AND BASES

_________________________________________________________________________________________________________ Show all calculations.

Part B

 

pH (calculated)

pH (measured)

H2PO4−

_______

_______

HPO42−

_______

_______

PO43−

12

_______

Write equations for the reactions taking place with the salts above.

Discussion:

  1. Based on your data and knowledge of dissociation constants, sketch (in lab notebook) the theoretical titration curves for a titration of Acetic Acid with NaOH and a titration of Phosphoric Acid with NaOH. Label all Equivalence and Half-Equivalence Points, as well as all pKas. State any assumptions made.
  2. Based on the today’s experiment describe a process to identify an unknown, weak acid. How would you determine the number of acidic groups present? What calculations should be done?
  3. List possible sources of errors in both Parts A and B. State how each error would effect the calculated Ka (part A) and pH (part B) values.

 

Experiment 27 SOLUBILITY PRODUCT CONSTANT (KSP) AND COMMON-ION EFFECT

 

 

 

 

 

Purpose

 

word image 1078 word image 1079 word image 1080 word image 1081 word image 1082 word image 1083

Determine the solubility product constant (Ksp) for a sparingly soluble salt. Study the effect on the aqueous solubility equilibrium of this salt in the presence of an additional amount of one of the ions that it contains (in this case, the effect of added Ca2+).

Introduction

word image 1084 This experiment deals with saturated aqueous solutions of salts with limited solubility in water. Under ordinary conditions, it might be fair to call the material you’ll use today as insoluble, or at best slightly soluble. At such conditions of limited aqueous solubility, the solid salt and its ions in solution are in a state of dynamic equilibrium between the solid phase and the solution phase. Concentrations of ions in solution are exceedingly low, and care must be taken to obtain good results, when performing the experimental procedure outlined shortly.

As an example of a slightly soluble salt, consider a saturated solution of silver sulfate, Ag2SO4.

In this example, the equilibrium lies far to the left, owing to the fact that silver sulfate is only slightly soluble in water:

Ag2SO4(s)  2Ag+(aq) + SO42(aq)

The corresponding equilibrium constant expression for this system is:

word image 1085

Since the solid component has a near-constant concentration, it is not used in the equilibrium expression, simplifying it to the solubility product:

K = [Ag+]2 [SO42−]

The new symbol Ksp is called the solubility product constant. Most solubility product constants have very small values, as is the case for silver sulfate, where Ksp = 1.1 x 10-5.

1.1 x 10-5= [Ag+]2 [SO42−]

In 1-liter of a saturated, aqueous solution of silver sulfate, the molar concentration of silver ion can be found by letting x equal the sulfate ion concentration, 2x the silver ion concentration, and substituting into the solubility product equation the known value for Ksp:

1.1 x 105 = [Ag+]2 [SO42−]

1.1 x 105 = [2x]2 [x]

1.1 x 105 = 4 x3 2.75 x 106 = x3 x = [SO42] = 1.4 x 10-2 M

+ -2 M

2x = [Ag ] = 2.8 x 10

Background

The equilibrium process in this experiment is a saturated aqueous solution of calcium iodate, Ca(IO3)2. The relevant solubility equation and solubility product expression, are both shown below.

Ca(IO3)2(s)  Ca2+(aq) + 2IO3 (aq)

Ksp = [Ca2+] [IO3]2

For a saturated solution of calcium iodate, if you can determine either the molar concentration of calcium ion, or the molar concentration iodate ion, the solubility product constant can be found using the reverse of the process shown above. There was found the silver ion concentration, in a saturated aqueous solution, from a known value for Ksp. In other words, if the calcium ion concentration in today’s experiment was found to be 0.1 M, you could immediately say the concentration of iodate ion must be half that value, or 0.05 M, according to the stoichiometry of the solubility equation given above. The solubility product constant could then be found with simple arithmetic. In this experiment, the iodate ion concentration of a saturated calcium iodate solution will be found via a redox titration with sodium thiosulfate, Na2S2O3.

Also, the effect of adding a source of additional calcium ion on the solubility of calcium iodate will be determined. This is accomplished via redox titration of a second calcium iodate solution containing an additional source of aqueous calcium ion (e.g., added CaCl2). It is possible to predict the result qualitatively upon the calcium iodate equilibrium, before actually performing the experiment.

According to LeChâtelier’s principle, adding an additional source of calcium ion (besides that of calcium iodate itself) should cause the solubility equilibrium to shift to the left, favoring more solid calcium iodate. This is an example of the common-ion effect. The common-ion effect, in this experiment, should lead to a reduced solubility of calcium iodate, and a corresponding change in the solubility product constant.

Details of the Redox Titration of Iodate Ion

The concentration of iodate ion (IO3) will be determined by titration with a standardized sodium thiosulfate (Na2S2O3) solution in the presence of potassium iodide (KI). Starch will be used as an indicator, and a sharp blue-to-clear transition will mark the equivalence point. The relevant reaction equations are summarized as follows.

IO3 (aq) + 5I (aq) + 6H3O+(aq) → 3I2(aq) + 9H2O(l)

This step, which occurs after adding both solid KI, and aqueous acid, to aliquots of saturated iodate solutions, has the net effect of converting iodate ions to aqueous iodine. Thiosulfate ion then reacts with aqueous iodine according to:

2- – 2-

I2(aq) + 2S2O3 (aq) → 2I (aq) + S4O6 (aq)

The net titration reaction can be obtained by combining the two reactions above, then balancing for mass and charge:

IO3 (aq) + 5I (aq) + 6H3O+(aq) → 3I2(aq) + 9H2O(l)

3I2(aq) + 6S2O32(aq) → 6I (aq) + 3S4O62 (aq)

IO3 (aq) + 6S2O32 (aq) + 6H3O+(aq) → I (aq) + 3S4O62 (aq) + 9H2O(l)

It is important to notice the molar ratio of iodate ion to thiosulfate ion in the above complete, balanced titration reaction. Ask your instructor if the thiosulfate solution available for you to use has been standardized, or whether you will have to standardize the thiosulfate solution as part of the experimental procedure. Aliquots from prepared, saturated solutions of calcium iodate in water, with and without additional calcium ion, will be titrated according to the reaction chemistry of the equations outlined above.

Experimental Notes

  • It is important to handle the saturated iodate solutions with great care, so as not to allow any of the solid calcium iodate into you titrations. This will obviously lead to a higher value for concentration of iodate than is actually present in the saturated solutions.
  • It is important for the Erlenmeyer flasks and other glassware you’ll use to be as clean as possible when performing these redox titrations. Please take a moment to at least give them a thorough rinsing–they need not be completely dry before proceeding.
  • Saturated solutions of calcium iodate in water, and calcium iodate with added calcium ion, are available in the hoods. Please take some care to not disturb the solid calcium iodate present in the bottles when obtained your portions of solutions to use.
  • It is recommended that you perform 4 titrations of unknown calcium iodate after standardizing the thiosulfate solution. The first of these should be a ‘scout’, where 1 mL portions of thiosulfate solution are added quickly to determine the approximate volume needed to reach equivalence.
  • You may find it convenient to set up all four calcium iodate unknown determinations at once. This is fine to do, as the titration solutions, once prepared, should remain stable until you are ready for them.
  • Titrated solutions, and any left over thiosulfate or iodate solutions, are safe to dilute and rinse down the sinks.

THE EXPERIMENT

  1. Rinse a buret with sodium thiosulfate solution that has already been standardized by the Preps Room. Record its concentration should be printed on the label. Using a graduated cylinder, measure out 50 mL of the sodium thiosulfate solution and add it to the buret.
  2. Add 5 mL of a calcium iodate saturated solution into a clean 125-mL Erlenmeyer flask. Add 20 mL of distilled water to this saturated iodate solution, swirl to mix. Dissolve ~1 g of solid KI into the iodate/water solution, then add 20 drops of 2-M HCl from a dropper bottle. Swirl to mix the contents, obtaining dark red homogeneous solutions. To complete the sample preparation, add approximately 30-40 drops of a 0.2% starch solution. The titration solutions should be a dark blue-black color. Make 3 of the samples.
  3. Repeat Step 2 using the solution of calcium iodate with added calcium ions. You will have a total of 6 samples.

Performing the Titrations

The first titration is a “scout” using an “Added Calcium Ion” sample, intended to make the remaining 4 titrations go faster and with greater accuracy. Perform the scout titration by adding 1 mL of the thiosulfate solution at a time, swirling and noting any changes in color. Remember that the starch/iodine endpoint is a sharp blue-black to clear transition, and it is unlikely you will miss it. Continue adding 1 mL portions of thiosulfate solution, until 1 mL added causes the solution to become clear and colorless. Assuming that all the titration solutions to follow have been prepared in the same manner as the scout, this gives an idea of what the actual equivalence volume is to ± 1 mL.

For the remaining titrations, it is permissible to add enough thiosulfate solution so that you are within 2 mL of the equivalence volume. Be sure and record initial volumes before titrating, and final volumes at equivalence, from the barrel of the buret to 2 decimal places. NOTE: The equivalence point will be different for the solution with added calcium ions, so you must do a scout for this titration as well.

Example:

Using the above scout procedure, imagine the equivalence volume was determined to be after adding 12 mL of thiosulfate solution; but before adding 13 mL. This means that 12 1 mL additions failed to produce a color change, but the 13th turned the solution from blue to clear.

For the next titration, an accurate initial volume of 0.06 mL was read, the titration was performed by adding 10 mL of the thiosulfate solution all at once to begin. If the samples were prepared identically, this should be before enough thiosulfate has been added to reach the equivalence volume. From this point, the titration continues on a drop by drop basis. Eventually, one single drop was added that caused the color of the titration solution to change from blue to colorless. The final volume reading was taken at that time to be 12.74 mL.

The volume added to reach equivalence is therefore carefully determined, and found to be (12.74 mL − 0.06 mL) = 12.68 mL. This volume, along with the molarity of the thiosulfate solution, allows the most accurate possible determination of moles thiosulfate.

It is a few steps from there to convert moles thiosulfate, to moles iodate ion, to concentration of iodate ion, finally to solubility product constant (Ksp) for calcium iodate.

Data Tables — Solubility Product Constant/Common-Ion Effect

Part A: Calcium Iodate, No Added Calcium Ion

 

Trial #1

Trial #2

Concentration of Sodium Thiosulfate Solution (M)

__________

__________

Volume of Calcium Iodate Solution Added (L)

__________

__________

Final Volume, Thiosulfate Solution (mL)

__________

__________

Initial Volume, Thiosulfate Solution (mL)

__________

__________

Total Volume, Thiosulfate Solution (mL)

__________

__________

Total Volume, Thiosulfate Solution (L)

__________

__________

Moles Sodium Thiosulfate Used

__________

__________

Moles Iodate

__________

__________

Equilibrium Concentration of Iodate Ion (M)

__________

__________

Equilibrium Concentration of Calcium Ion (M)

__________

__________

Molar Solubility of Calcium Iodate (M)

__________

__________

Ksp, Calcium Iodate

__________

__________

Average For Ksp, Calcium Iodate

__________

Data Tables — Solubility Product Constant/Common-Ion Effect

Part B: Calcium Iodate, Added Calcium Ion

 

Trial #1

Trial #2

Concentration of Sodium Thiosulfate Solution (M)

__________

__________

Volume of Calcium Iodate Solution Added (L)

__________

__________

Final Volume, Thiosulfate Solution (mL)

__________

__________

Initial Volume, Thiosulfate Solution (mL)

__________

__________

Total Volume, Thiosulfate Solution (mL)

__________

__________

Total Volume, Thiosulfate Solution (L)

__________

__________

Moles Sodium Thiosulfate

__________

__________

Moles Iodate

__________

__________

Equilibrium Concentration of Iodate Ion (M)

__________

__________

Equilibrium Concentration of Calcium Ion (M)

__________

__________

Molar Solubility of Calcium Iodate

(added Ca2 ion source)

__________

__________

Average Molar Solubility of Calcium Iodate

(added Ca2 ion source)

“Scout Titration”

__________

Volume of Calcium Iodate Solution Added (L)

__________

Final Volume, Thiosulfate Solution (nearest mL)

__________

Initial Volume, Thiosulfate Solution (nearest mL)

__________

Total Volume, Thiosulfate Solution (nearest mL)

__________

 

DATA ANALYSIS SUPPLEMENT

Results:

Your lab report should include the following in addition to the pages from your lab notebook:

  • An Excel sheet “Raw Data” showing the volumes of sodium thiosulfate, calcium iodate, and calcium iodate with added calcium ion solutions for each trial.
  • A 2nd Excel sheet “Results Part A” which should include the columns:
    • Volume Thiosulfate o Moles Thiosulfate o Moles Iodate o Equilibirum concnetraion of Iodate Ions o Equilibirum concnetraion of Calcium Ions o Molar Solubity o Ksp o Average Ksp o % Error

 

  • A 3rd Excel sheet “Results Part B” which should include the columns:
    • Volume Thiosulfate o Moles Thiosulfate o Moles Iodate o Equilibirum concnetraion of Iodate Ions o Equilibirum concnetraion of Calcium Ions o Molar Solubity o Average Molar Solubility
  • An Excel sheet ‘Histogram’ depicting the spread of the Experiement Ksp values.

 

Discussion:

  1. Discuss the calculated Ksp, and compare that to the known Ksp value.
  2. Discuss the results of the common ion (added Ca) trial. Were they as expected? Why or why not?
  3. Discuss sources of error and how these errors could have led to the results. Relate your answer to the class histograms.

 

 

Experiment 28 THERMODYNAMICS OF THE DISSOLUTION OF BORAX

 

 

 

word image 1086 word image 1087 word image 1088 word image 1089 word image 1090

SIMULTANEOUS DETERMINATION OF SEVERAL THERMODYNAMIC

QUANTITIES: K, G°, H°, AND S°

Purpose

Study a system of a sparingly soluble salt in water. From the solubility information at various temperatures, a variety of other thermodynamic quantities can be determined for the system. Among these thermodynamic quantities are the changes in standard enthalpy, standard entropy, and standard free energy.

Introduction

word image 1091 The system you will be studying involves a relatively simple solubility equilibrium of borax (Na2B4O7•10H2O) in water:

Na2B4O7 • 10H2O(s)  2Na+(aq) + B4O5(OH)42(aq) + 8H2O(l)

Note that the reaction equation, as it is written above, has a few features you should notice:

1). It is written as an equilibrium process

2). Eight (8) water molecules from the hydrated salt are lost to the reaction medium

When defining an equilibrium constant expression for this reaction, the “waters of hydration”, written as a reaction product above, must now be represented in any expression for the equilibrium constant for this process. Thus, the proper equilibrium expression for the reaction above has this form:

word image 1092

The experimental procedure will specifically state that you are to make sure there is always some solid borax remaining in the sample mixture before you remove some of it to analyze. That being the case, the ‘concentration’ of solid borax can be assumed constant, and the equilibrium expression can be simplified:

K = [Na+]2 [B4O5(OH)42] [H2O]8

The water term in the expression (which was originally part of borax’s crystalline matrix) is lost to the sample mixture. This added amount of water does not significantly change the concentration of water already present. The water term can also be ‘simplified out’ of the equilibrium expression:

K = [Na+]2 [B4O5(OH)42]

What remains is now a solubility product constant (Ksp) expression for borax in water. To determine a value for the solubility product, a method must be found to assay either the amount of sodium ion, or borate ion, in the sample mixture. The original equilibrium expression, and balanced solubility equilibrium reaction, leads to a convenient way to express either ion in terms of the other, so that it is possible to define Ksp in terms of either the concentration of sodium ion, or concentration of borate ion. A series of substitutions, based on the original balanced solubility equilibrium equation, gives the desired Ksp, expression defined in terms of the borate ion only:

[Na+] = 2 [B4O5(OH)42]

K = [ (2 [B4O5(OH)42] ) ]2 [B4O5(OH)42 ]

Ksp = 4 [B4O5(OH)42]3

Finding the concentration of borate ion, in any sample at any given temperature, leads directly to a value for solubility product (Ksp) at that temperature.

Design of the Experiment and Relevant Background

For a series of borax-containing samples taken at various temperatures, the concentration of borate ion will be determined. The borate ion reacts with monoprotic acid (such as HCl) in a 1:2 fashion:

B4O5(OH)42 (aq) + 2HCl(aq) + 3H2O(l) → 4 B(OH)3(aq) + 2 Cl (aq)

The amount of HCl consumed (in Liters) multiplied by the molarity of the acid gives moles HCl. It is a trivial exercise to determine moles borate ion from there, and dividing by the volume of borate ion-containing sample in liters gives concentration of borate ion. This general procedure is repeated for borate ion-containing samples of constant volume obtained at various temperatures.

Calculation of these temperature-dependent values for Ksp of borate ion is but the first step in a greater sequence of obtaining complete thermodynamic parameters for the dissolution of borax in water.

word image 1093 There are several other important thermodynamic parameters that can be found: free energy change, enthalpy, and entropy, upon further treatment of the equilibrium constant/temperature base data set. The next goal will be to estimate the free energy change for this solubility equilibrium. The following equation immediately shows the relationship between free energy change ( ) and equilibrium constant (K):

G = RT ln Ksp (A)

A value for the equilibrium constant at a given temperature gives (after a sense) the value for free energy change directly. The free energy change at a given temperature is itself related to both the change in enthalpy, and the change in entropy, by the following equation:

G = H  TS (B)

Combining the two equations relates K, H, and S in a single statement:

RT ln Ksp = = H  TS

The preceding equation can be rearranged into a form that is linear. When re-arranged into a linear form, 1/T and ln Ksp can be used as (x,y) points on a graph. The slope of such a graph is related to the change in enthalpy and the change in entropy is related to the y-intercept:

A summary of the post-data collection activities reveals the true beauty of this experiment. Once the solubility product constants have been determined for 5 different temperatures:

ln Ksp = -H + TS

RT RT

ln Ksp = -H+ S

1

T

RR

y = m x + b

word image 1094

compiled. This is most conveniently done using equation (A), with values for Ksp and the temperatures for which they are valid.

  • A graph of ln Ksp vs. 1/T should be made, the slope of which is related to change in enthalpy.
  • The same graph has a relationship between change in entropy and its y-intercept

The experiment requires that the solubility of borax be found at various temperature values. Samples of saturated borax solution are collected at no less than 5 different temperatures, four above room temperature, and one close to or at room temperature. These samples are then warmed (if necessary) to re-dissolve any precipitated borax, and titrated to the yellow bromocresol green endpoint with standardized aqueous hydrochloric acid.

Experimental Procedure (groups of 2 students)

Note: To complete this procedure, it will be necessary to carefully divide the tasks between partners. Take a moment before beginning to make a plan between the two of you to do this.

  1. Add 2 scoops of borax and 75 mL of distilled water to a 400-mL beaker. Heat the solution until it is 5-8°C above the assigned temperature. Heat the solution slowly and ensure that it does not go over 50°C. Once the solution is slightly above the assigned temperature, take the beaker off the hotplate. The solution must be saturated, therefore, if all the borax has dissolved add more to the beaker.
  2. At the same time, heat 150 mL of distilled water to your assigned temperature. This will be used for your dilutions.
  3. Pour 5 mL of the solution (no solid) into a graduated cylinder once the solution is at the desired temperature. Quantitatively transfer this solution to a 125-mL Erlenmeyer flask. In addition, add 50 mL of the distilled water that has been heated to the assigned temperature and 5 drops of the bromoscresol green indiciator to the Erlenmeyer flask. Swirl to mix.
  4. There will be a total of three samples made this way.
Titration of the Borax Samples
  1. Rinse a buret with HCl. Using a graduate cylinder, measure out approximately 20 mL of HCl and added it to the buret.
  2. Carry out a scout titration, adding the HCl in 1 mL, incriments until the endpoint (yellow color) is reached. For the remaining titrations, it is permissible to add enough HCl so that you are within 2 mL of the equivalence volume and then continue to add the HCl dropwise until the endpoint is reached. Be sure and record initial volumes before titrating, and final volumes at equivalence, from the barrel of the buret to 2 decimal places.

 

 

Data Tables — Simultaneous Determination of Several Thermodynamic Quantities

Molarity of Standardized HCl

__________

Room Temperature, °C _______

Room Temperature, K _______

Solution Temperature

Trial #1

 

Trial #2

Volume of Borax Solution

__________

 

__________

Final Volume, HCl

__________

 

__________

Initial Volume, HCl

__________

 

__________

Total Volume, HCl

__________

 

__________

Moles HCl

__________

 

__________

Moles Borax

__________

 

__________

Concentration of Borax in sample

__________

 

__________

Solubility Product, Ksp __________ __________

Average Ksp __________

Free Energy Change, G __________ __________

Graphically Determined Values

Change in Enthalpy, H __________

Change in Enthalpy, S __________

 

DATA ANALYSIS SUPPLEMENT

Results: Outside of class, retrieve your data in spreadsheet form (along with that of the entire class) from: http://chem.gmu.edu/results/. Then use Excel to calculate the moles of HCl titrated, moles of Borax, molarity of Borax solution, Ksp, then ΔH and ΔS from a graph of ln(Ksp) vs 1/T. Finally, you will determine ΔG using two different equations and compare the two.

Your lab report should include the following in addition to the pages from your lab notebook:

  • An Excel sheet “Raw Data Sheet” showing the raw data for the entire class.
  • A 2nd Excel worksheet “Calculation of Ksp”. This should include the following columns:
    • Volume HCl titrated in L for trial 1 o Volume HCl titrated in L for trial 2 o Moles of Borax for trial 1 o Moles of Borax for trial 2 o Concentration of borax for trial 1 o Concentration of borax for trial 2 o Ksp calculated from concentration of borax for trial 1 o Ksp calculated from concentration of borax for trial 2 o Average Ksp
  • A 3rd Excel worksheet “Calculation of Thermodynamic Values” showing the calculated results for each student in the class – each in a separate row. This should include the following columns:
    • Natural log of average Ksp o Temperature (converted to Kelvin) o Reciprocal Temperature (1/T) o Slope of ln(Ksp) vs. 1/T graph o Intercept of ln(Ksp) vs. 1/T graph o Gibbs Free Energy calculated using equation A: G = RT ln Ksp o Average of ΔG o Enthalpy (found from slope) o Entropy (found from intercept)
    • Gibbs Free Energy calculated using equation B: G = H – TS o Average of ΔG
  • A 4th, 5th, and 6th Excel sheet “Graphs” showing a graph of ln(Ksp) vs 1/T, ΔG vs T (using

Equation A to calculate ΔG), and ΔG vs T (using Equation B to calculate ΔG)

 

 

Discussion:

  1. Compare your two calculated G values; do they agree? Based on your values, is the dissolution of borax spontaneous?
  2. Why is it necessary to ensure that there was always some solid borax in the solution?
  3. What was the Ksp value for your temperature? What does this say about the solubility of borax at this temperature? Compare this to the Ksp values at other temperatures and make a statement regarding the effect of temperature on solubility. If your results do not match what you expected, analyze sources of error.
  4. Literature values for enthalpy and entropy of the dissolution of borax in water are 110 kJ/mol and 380 J/mol*K respectively. Do your values agree? Calculate your percent error and discuss sources of experimental error. You assumed that H and S are constant over the temperature range studied. Are they?

 

 

 

 

 

 

 

 

 

 

 

 

Experiment 29 .ELECTROCHEMICAL MEASUREMENTS

 

 

 

 

word image 1095 word image 1096 word image 1097 word image 1098 word image 1099

ELECTROCHEMICAL CELLS: DETERMINATION OF REDUCTION POTENTIALS FOR A SERIES OF

METAL/METAL ION SYSTEMS, VERIFICATION OF NERNST EQUATION, AND DETERMINATION OF

FORMATION CONSTANT OF Cu(NH3)42+AQUA COMPLEX

Purpose

word image 1100 There are three parts in today’s lab experiments. In part A, the purpose is to construct simple electrochemical cells, and measure cell potentials. Through these measurements, it will be possible to calculate the standard cell potential and confirm the observed “activity series” for some of the common metals. In part B, you will verify the Nernst equation by measuring the cell potential as a function of ion concentration in one of the half-cells. In part C, the measured cell potential, after addition of NH3 as the complexing agent, in the Cu/Cu2+ half-cell will be used to calculate the formation constant of a Cu(NH3)42+ aqua complex.

Introduction

You will need to rely heavily on the explanations of voltaic cells, concentration cells, and the Nernst equation that are provided in your lecture text. There is simply not enough space to cover all the necessary information here in the lab manual. This text is designed to help you to build the cells, take the experimental measurements, and suggest ways to interpret and quantify the information you will obtain while doing the experiment. For your information, an abbreviated table of standard reduction potentials is included as part of this text.

Design of the Experiment: Standard and Non-Standard Reduction Potentials For Metal/Metal Ion Half-Cells

Typical electrochemical cells are comprised of two half-cells, linked by a salt bridge (allowing transport of ions in both directions), with a voltmeter completing the circuit to measure the difference in voltage between the two halves. One half-cell has a reduction taking place (the cathode) while the other half-cell has an oxidation taking place (the anode). The cell potential, Ecell, as measured by a voltmeter is a measure of the free energy change, G = nFEcell, that occurs in the overall oxidation/reduction reaction.

Electrons will move in response to this potential difference, with the flow going toward the cathode from the anode. Thus, the sign of the potential difference (as measured by a voltmeter) is sufficient for identifying which half-cell has the oxidation reaction and which has the reduction reaction.

 

Standard half-cell reaction potentials are measured using the standard hydrogen electrode (SHE) as a reference point. The standard hydrogen electrode is assigned a value for potential of 0.00 V, and the half-cell reaction for the standard hydrogen electrode is the reduction of 1.00 M aqueous acid to gaseous hydrogen at 1 atm pressure on a platinum electrode:

+  H2 (g, 1 atm)

2H (aq, 1 M) + 2e

All other half-cell reactions can be measured in tandem with the standard hydrogen electrode, which gives rise to the table of standard reduction potentials, an abbreviated version of which appears as part of this text.

Measurements of standard half-cell potentials are supposed to be carried out at standard conditions i.e., at solution concentrations of 1.0 M and at 25°C. It is not always practical (or possible) to carry out electrochemical reactions at these conditions, as is the case in today’s experiment. For all reversible reactions the cell potentials measured at non-standard conditions (Ecell) can be related to potentials at standard conditions (E°cell) by the Nernst equation: Ecell = Eocell − (0.0591/ n) log Q (Q = equilibrium quotient)where, Eocell = Eooxidation + Eoreduction

Also, it is not practical to set up a standard hydrogen electrode for each of you to do your measurements. Instead, you will use the potential of the Ag+/Ag reduction half-reaction as your standard (+0.80 V). Using this value, and assuming it to be the cathode half-reaction, will allow you to determine half-cell potentials for the other systems you will study (half-cells of Cu, Sn, Pb, and Zn) whose general reduction half-reactions are of the form:

M[3]+ + 2e  M

As an example of today’s measurements, let us look at the half-cell reactions of Al/Al[4][5]+ and Sn/Sn[6]+ at standard conditions. The reduction half-reactions and potentials are as follows (from table of standard reduction potentials):

Overall Reaction: 2Al + 3Sn2+  2Al3+ + 3Sn Ecell° = +1.52 V

In today’s experiment, solution concentrations are 0.1 M (temperature effects will not be considered) and not the 1.0 M of “standard conditions”. Thus the Nernst equation that would be needed to “correct” the cell potential of the above example to non-standard conditions is:

Ecell = Eocell − (0.0591 / n) log ([0.1]2 / [0.1]3) where, n = 6 (the overall number of electrons required to balance the reactions)

Ecell = 1.52 − (0.0099)

= 1.51 V  This is what you should observe on the voltmeter

The above principle can be used to measure the standard potential of an unknown redox couple if the cell potential is measured and if one half of the couple has known standard reduction potential (in this case aluminum). Continuing the above example, if you know the “standard” reduction potential for the aluminum half cell (−1.66 V from Table), you then measure the overall cell potential with the voltmeter, and solve for the half-cell potential for tin.

Standard reference half-cell potential: Eoox = +1.66 V (for the Al3+/Al half-cell as anode)

Measured overall cell potential: Ecell = +1.50 V

Solution concentrations: [Sn2+] = 0.01 M and [Al3+] =0.01 M

+1.50 = Eocell – (0.0591 / 6) log ([0.01]2 / [0.01]3)

Eocell = + 1.52 V = Eoox + Eored o =  0.14 V (Non standard reduction potential for Sn2+ + 2e  Sn)

E red

So…the actual objective is to determine half-cell potentials for Sn, Pb, Cu and Zn at non-

+ standard conditions, using the Ag /Ag half-reaction as your standard reference and examining the all-possible combinations of half-cells.

 

Notes on Experimental Procedure and Precautions

  • There are 15 sets of pre-assembled and labeled weighing boat setup’s located on the various benches. Obtain one of these for use by your group of two students.
  • There are also several sets of solutions placed around the room on the benches. These are used by those students occupying a given bench space, so please use what you need and leave the sets where you found them so that others may also get what they need. No one should have to cross the lab to find a reagent they need.
  • Treat the pH meters (being used here as voltmeters) with care, and return them to the cabinets when you are finished.
  • Waste bottles for silver and lead salts generated during this experiment are in the hoods. Please carefully follow the waste disposal directions.
  • You must not spill silver solution on your skin, clothes or on the bench; it may produce a semipermanent mark where spilled. Promptly clean and wash the spilled area with plenty of water.

 

Experimental Procedure

A. Measurement of Non-Standard Reduction Potentials
  1. Assemble the cell as depicted in below. Do not let the salt bridges dry out. The red wire goes with the silver wire and into the clear silver solution and the black wire goes on the other metals (Cu 0.01M and Zn strip).

 

 

 

 

 

 

 

 

 

 

 

word image 3241

 

  1. With one lead still on the silver electrode, move the other lead to the other metals to take your reading.

To construct “salt bridges”, dip one end of a piece of filter paper into the boat containing KNO3 solution, and the other into a boat containing a metal ion solution. A total of 5 such pieces of filter paper must start from the KNO3 boat, and end up in each of the 5 different solutions. Do not let the salt bridges dry. Keep a medicine dropper handy to put a few drops of KNO3 on the filter paper whenever necessary.

Obtain pieces of Cu, Ag, and Zn metals to use as electrodes. Place each metal strip into its corresponding metal ion solution (copper metal into the copper nitrate solution, etc), bending each strip as needed to allow a portion to stick above the solution surface for potential measurements. It is important to handle these strips as little as possible, and never with bare fingers.

Attach electrical leads to the pH meter by carefully inserting the connector and giving a slight twist. Making sure the meter is set to “mV”, turn on the meter and connect the two leads (alligator clips) to one another. A voltage reading of zero should result. If not, use the calibration knob to set the voltage to zero. Digital meters usually autozero by itself.

Attach one lead to the end of the strip of zinc metal, the other lead to the strip of silver metal and make sure that the metal strips remain dipped in solution. If the voltage is negative, reverse the leads to obtain a positive voltage. Measure the instantaneous cell voltage and record two more readings immediately after that. Take the average of the three as the measured cell potential.

At this point, the silver electrode is acting as the cathode (i.e., reduction is occurring at the silver metal). Leaving the connection to the silver electrode intact, move the other lead to the other metals in any sequence you wish and take reading. Obtain voltage values for all possible combinations of two metals.

Convert your readings to volts (if you have been taking them in millivolts) and use the Nernst equation to find half-cell potentials for the Zn2+/Zn and Cu2+/Cu half-cells. Remember, Ag+/Ag half-cell as the reference, assuming it to be the cathode, with a potential of Eored = + 0.80 V (+0.7993 V to be exact).

B. Verification of Nernst equation for Cu/Cu2+ half-cell by changing [Cu2+]

Measure cell potentials for all Cu/Cu2+ (0.1, 0.01, and 0.001 M) // Ag+ (0.10 M)/Ag cells. You should have three measurements for three different copper concentrations (0.1, 0.01, and 0.001M). For the overall cell reaction: 2Ag+ + Cu (s)  Cu2+ + 2Ag (s), according to Nernst equation:

Ecell = Eocell  (0.0591/2) log ( [Cu2+]/ [Ag+]2)

a plot of Ecell vs. log ( [Cu2+]/ [Ag+]2) should be a straight line with a slope = −0.0295 V (per tenfold change in concentration) and an intercept, Eocell. By knowing standard reduction Eo for Ag/Ag+ and the Eocell, you can calculate the standard reduction potential for Cu/Cu2+. Compare the calculated value with that of the literature.

 

 

Waste Disposal and Clean Up

Place two drops of concentrated HCl (or H3PO4) into the Pb and Ag solutions, and stir briefly. Allow the solid formed to settle, and decant the solution into the drain. Place the solid slurries of PbCl2 (or Pb3(PO4)2) and AgCl (or Ag3PO4 ) into their appropriate containers in the hood. All other solutions may be rinsed down the drain. Return the metal strips to the containers near the balances.

_____________________________________________________________________________

ELECTROCHEMICAL MEASUREMENTS

_____________________________________________________________________________ Part A

Cell #

Electrode

(Cathode)

Reduction

1

Electrode (Anode)

Oxidation

2

Ecell (V)

Measured

Eocell, (V)

Nernst

Eocell(V)

Standard

1

Ag

 

Zn

 

 

 

 

2

Ag

 

Cu

 

 

 

 

5

Zn

 

Cu

 

 

 

 

 

Explanations
  • First, identify the anode (Ox) and cathode (Red) and place them in proper columns. Placement shown in table may not be correct.
  • EoCell Nernst is obtained from the Nernst equation, (EoCell Nernst = Ecell, Measured + (0.0591/n) log Q). This is the non standard cell potential. You must write a balanced redox reaction to find n.
  • EoCell Standard is the standard cell potential = Eoox + Eored. Standard potentials are listed in Table 2 as obtained from table of standard reduction potentials in your textbook.

 

Part B

Half Cell Reaction

Eored (V)

Non standard

Eored (V)

Table value

Percent error

Ag+ + e  Ag

 

+0.7993

N/A

 Cu2+ + 2e  Cu

 

+0.339

 

 Zn2+ + 2e  Zn

 

−0.762

 

Note: Knowing the Eocell, Nernst from Table 1 and given the EoAg/Ag+ one can calculate the nonstandard reduction potential.

Discussion
  1. List the major sources of experimental errors. Which parameter should be measured most accurately to obtain the best result? (Hint: Check Nernst equation)
  2. Is it possible to obtain the number of electrons required for the overall reaction from the slope of the graph? How?
  3. What is the purpose of using the Nernst equation in this experiment?

 

 

Experiment 30 IDENTIFICATION OF AN ORGANIC ACID

 

 

 

 

 

word image 1101 word image 1102 word image 1103 word image 1104 word image 1105

Introduction

In this experiment, you will measure the melting point, molecular weight, and observe the solubility of an unknown organic acid. Using this data, you will be able to identify the acid from a given list of possible unknowns. The purpose of this experiment is to learn how to determine the melting point using a melting point apparatus, to gain experience with various solubility tests, and to practice titration to determine molar concentration of an unknown.

Theory of Melting Point Determination

word image 1106 The melting point of a compound can be used as an indicator of the compound’s purity and as a means of identification. Pure compounds melt within a 1-3 degree range with a higher melting point, and impurities in the sample will result in a broader melting point range and a lowering of the melting point relative to the pure compound.

You will be measuring the range for the melting point of your unknown acid; the first temperature in your range will be recorded when the first drop of liquid appears, and the second value is the temperature at which the entire sample turns to a clear liquid. Sometimes, NONMELTING behavior can be seen such as discoloration, decomposition, shrinkage, etc. Don’t worry! There’s nothing wrong. Just record what you observe. It is alright if you obtain slightly different results for the same compound (within 1-2 degrees).

THE EXPERIMENT

Part 1: Determination of Melting Point
  1. Obtain an unknown organic acid sample from your instructor and record the unknown number in your laboratory notebook. Obtain a Mel-temp and record its number in your lab notebook
  2. Prepare the capillary tube using the following steps:
    • Crush a small sample of your acid with a spatula or watch glass o Lightly press the open end of the tube into the sample so that there is about 1-2 mm of your sample in the tube
    • Drop the tube through a length of glass tubing to pack the sample into the bottom of the capillary tube. Do NOT tap the tube against the lab bench!
  3. For your samples, do one fast determination to get a rough approximation of your melting point range. Record this range in your notebook.
  4. Next, after allowing your Mel-temp to cool, heat rapidly to within 10-15 degrees of the anticipated melting point, then lower the heat setting to slow the heating, approximately 12 degrees/min. You do NOT want to be heating rapidly when you hit your anticipated melting point as this will give you an inaccurate range.
  5. Record temperature at which first drop appears and then record temperature at which all solid is melted
  6. Compare the experimental range with literature value(s).

Notes on Melting Point Procedure: Take only a small amount of your samples. Remember to record the number on the vial containing the unknown. Record the number of your Mel-Temp. Do not attempt to redo the melting point on a previously melted sample. Put used tubes in the broken glassware container.

Part 2: Solubility Tests

Because the titration requires that the acid be in solution, it is necessary to determine the solubility of the unknown acid in suitable titrations – water, ethanol, or a mixture of the two.

  1. To test the solubility of your compound in water, add ~40 mg of the unknown acid (small amount on the end of a spatula, do not weigh) to 2-3 mL of water in a medium sized test tube.
  2. Shake carefully for approximately 1 minute.
  3. If water does not completely dissolve the sample of the acid, repeat the procedure in a clean test tube using ethanol instead of pure water.
  4. If the sample will not dissolve in ethanol, then confirm that the sample dissolves in a 50:50 water-ethanol mixture
  5. Check with instructor before proceeding
Part 3: Titration to Determine Molecular Weight of Acid

Note: Your first titration will be a “scout” titration to determine, roughly, the volume of NaOH required to reach the endpoint. Your second and third titrations will be your quantitative trials.

  1. Accurately weigh about 0.2 g of the unknown acid.
  2. Quantitatively transfer the acid to a 125-mL Erlenmeyer flask. This means that you should carefully rinse out your weigh boat with a small aliquot of solvent (see next step) to ensure that ALL of the solid acid has been transferred to the Erlenmeyer flask.
  3. Add 15 mL of the solvent which will dissolve the acid sample as determined previously to the Erlenmeyer flask.
  4. Add 2-4 drops of indicator

o Water or 50:50 ethanol mixture → use phenolphthalein o 95% ethanol → use bromothymol blue

  1. Rinse the buret; first with water and then with NaOH.
  2. Add approximately 20 mL of the NaOH to the buret. Record initial volume of NaOH in the buret. You can refill the buret if need be.
  3. Add the NaOH from the buret to the Erlenmeyer flask containing the acid sample, somewhat rapidly at first, then more slowly. Notice the color change as you add base, this will help you determine when you are nearing your endpoint.
  4. Swirl the flask continuously as you add the NaOH.
  5. Endpoint is achieved when the solution in the flask changes color o Phenolphthalein: colorless→ pink (red) o Bromothymol blue: yellow → blue

After you have performed these tests, review your results and determine the identity of your unknown acid from the list given below. Check with your instructor.

 

Table 1. List of Possible Organic Acids

Acid

Molecular

Formula

Melting Point

(C)

Acidic Groups

Mol. Wt. (g/mol)

meta-toluic

C8H8O2

108-110

1

136

azelaic

C9H16O4

109-11

2

188

dl-mandelic

C8H8O3

120-122

1

152

benzoic

C7H6O2

122-123

1

122

dl-malic

C4H6O5

131-133

2

134

trans-cinnamic

C9H8O2

133-134

1

148

maleic

C4H4O4

134-136

2

116

ortho-chlorobenzoic

C7H5ClO2

138-140

1

148

acetylsalicylic

C9H8O4

140-142

1

180

benzilic

C14H12O3

150-153

1

228

adipic

C6H10O4

152-154

2

146

citric (anydrous)

C6H8O7

152-154

3

192

salicylic

C7H6O3

158-160

1

138

1-naphthoic

C11H8O2

160-162

1

172

para-tert-butylbenzoic

 

L-tartaric

C11H14O2

C4H6O6

165-167

170-172

1

2

178

150

para-toluic

C8H8O2

180-182

1

136

succinic

C4H6O4

185-188

2

118

para-hydroxybenzoic

C7H6O3

214-216

1

138

Discussion:

1. What was the identity of your unknown organic acid? Discuss the tests you performed and how each test helped you identify and characterize your unknown. Some ideas of topics you should discuss:

  • How does your experimental m.p. compare to the theoretical m.p.? What does your experimental m.p. range tell you about the purity of your sample?
  • Compare your experimental molecular weight to the theoretical molecular weight.
  • How do the structural features of the molecule factor into the melting point? Include a drawing of the structure for your organic acid.
  • If there were results that disagreed, offer analysis.
  • Discuss possible sources of error.

 

 

 

APPENDIX

What follows are instructions on using Vernier data acquisition equipment to take data from certain probes that attach to it. The main data acquisition (LabQuest2) is shown Figure A- 1 below.

word image 3242 word image 1107

Figure

A

 

3

 

Temperature probe

 

word image 3243 word image 1108

Figure A

 

1

 

Vernier Data Acquisition Unit (LabQuest 2)

 

This unit control collects and analyze data from a sensor. Several sensors are used in the General Chemistry classes. These probes are listed below along with instructions on how to acquire data. They plug into the LabQuest 2

pH probe is used to measure the pH LabQuest 2 via a USB port. This will be used when you need to measure the pH in an experiment such as in Experiment 24Experiment 25.

word image 3244 word image 1109

Figure A

 

2

 

pH probe

 

 

 

 

 

 

Temperature probe. As the temperature varies the probe develops a small voltage that is a measure of the temperature.

Thus, with proper settings in the LabQuest 2 one can know the temperature of the material in contact with this probe. This will be useful in experiments (such as Experiment 9) where heat is absorbed or evolved and we need to know how this affects the temperature which might lead to something like the determination of the heat of formation.

Voltage probe In certain experiments such as when we do electrochemistry (e.g. Experiment 29 where we measure the voltage of several cells we make) a voltage is developed between wires in different half-cells. This probe allows the voltage to be measured and can help us figure out the half-cell potential of one or more half-cells.

word image 3245 word image 1110

Figure A

 

4

 

Voltage probe

 

 

 

 

Spectrophotometer. The amount of light absorbed at different wavelengths can be measured using a spectrophotometer. The Go Direct SpectroVis Plus

word image 3246 word image 1111

Figure A

 

5

 

Go Direct SpectroVis Plus

Spectrophotometer

 

Spectrophotometer uses an LED and Tungsten light source to produce radiation which goes through a small monochromator (diffraction grating). The reduction in the

amount of light making it to the detector is measured and is a function of the wavelength. This can allow us to determine wavelength of maximum absorption as well as the concentration of the compound absorbing radiation.

Experiment 10 andExperiment 24 are examples where this instrument might be used.

Instructions on how to use these probes follows.

 

A 1. INSTRUCTIONS FOR THE PH PROBE

1. Set up

  1. Plug in the pH sensor to one of the CH# ports (left side).
    1. The Labquest App should show the red box showing the sensor

1. “CH #: pH” and the current reading

  1. Tap the “Mode/Rate/Duration” box. Set:
    1. Mode = “Time Based”
    2. Rate = 2 samples/s
    3. Duration = how long your measurements will take
  2. Calibrate: Tap the red box showing the current pH
    1. word image 3247
    2. Tap ‘Calibrate’. Tap ‘Calibrate Now’
    3. word image 3248 iv. Submerge the pH probe in the first buffer solution.

Enter the pH of the current buffer solution.

word image 3249

  1. When the “Live voltage” is stable, tap ‘Keep’
  2. Rinse and dry the pH probe, then submerge in the second buffer solution and repeat

word image 3250 vii. After tapping ‘Keep’, you return to the live pH reading. If the probe is still in the buffer solution, it should read within +/- 0.01 of the buffer pH.

viii. word image 3251

  1. Measurement
    1. To start a measurement run, click the play button on screen or the hardware button on the right of the Labquest
    2. The Labquest will switch to the graphing tab and graph the data as it collects
    3. word image 3252
    4. To stop the run early, push the stop button on the screen, or press the hardware button again

i. The graph will autoscale after stopping a run.

    1. To measure another data run, temporary store the current data by clicking the file cabinet button

i. The current data graph will disappear, ready for a new data set

  1. Processing
    1. To view graphs of multiple runs, tap “Run #” and select “All Runs”
    2. To view individual data points, tap the location on the graph you would like to read
        1. The measurement at that particular time will show on the right for each data run
        2. You can also click the Table tab, which shows all data entries for all current runs
  2. Save data
      1. *The file cabinet does NOT save the data. It only temporarily stores it to compare to other runs*
      2. To save all open data to the device for easy opening later, click File -> Save…
          1. Click on “untitled” to set a filename for your dataset
          2. Click “Save”

c. To save all open data to a flash drive, insert a USB flash drive into the Labquest

          1. Click File -> Export…
          2. Verify that a USB icon is in the upper left corner. iii. Click on “untitled” to set a filename for your dataset

iv. Click “Save”

  1. Shutdown
      1. Unplug the pH sensor and submerge the probe in a storage solution.
      2. Click File -> Quit
      3. Push the power button on top of the Labquest

 

A 2. TEMPERATURE PROBE

 

  1. Set up

a. Plug in the temperature sensor to one of the CH# ports (left side).

i. The Labquest App should show the red box showing the sensor 1. “CH #: Temperature” and the current reading

  1. word image 3253

b. Tap the “Mode/Rate/Duration” box. Set:

i. word image 3254 ii. Mode = “Time Based”

  1. Rate = 2 samples/s
  2. Duration = how long your measurements will take
  3. Measurement
    1. To start a measurement run, click the play button on screen or the hardward button on the right of the Labquest
    2. The Labquest will switch to the graphing tab and graph the data as it collects
    3. To stop the run early, push the stop button on the screen, or press the hardware button again
      1. The graph will autoscale after stopping a run.
    4. To measure another data run, temporary store the current data by clicking the file cabinet button
      1. word image 3255
      2. The current data graph will disappear, ready for a new data set
  4. Processing
    1. To view graphs of multiple runs, tap “Run #” and select “All Runs”
    2. word image 3256
    3. To view individual data points, tap the location on the graph you would like to read
      1. The measurement at that particular time will show on the right for each data run
      2. You can also click the Table tab, which shows all data entries for all current runs
  5. Save data
    1. *The file cabinet does NOT save the data. It only temporarily stores it to compare to other runs*
    2. To save all open data to the device for easy opening later, click File -> Save…
      1. Click on “untitled” to set a filename for your dataset
      2. Click “Save”
    3. To save all open data to a flash drive, insert a USB flash drive into the Labquest
      1. Click File -> Export…
      2. Verify that a USB icon is in the upper left corner. iii. Click on “untitled” to set a filename for your dataset

iv. Click “Save”

  1. Shutdown
    1. Unplug the Temperature sensor
    2. Click File -> Quit
    3. Push the power button on top of the Labquest

A 3. SPECTRO VIS INSTRUCTIONS

  1. Set up
    1. Plug in the SpectroVis USB cable to the LabQuest2 USB port (left side)
      1. In the LabQuest App, the red Sensor box will appear “USB: Abs”
      2. word image 3257
    2. Tap anywhere inside the red USB: Abs box. Click “Calibrate…” on the menu that appears
      1. word image 3258
      2. The SpectroVis will take a dark sample
      3. The SpectroVis will wait 90 seconds for the lamp to warm up

1. word image 3259

      1. Insert a cuvette for a blank spectrum. Wipe the cuvette sides before inserting. Click “Finish Calibration”. Click “OK”
          1. The smooth sides should be facing the white light (Abs. light source, right) and the white triangle (detector, left) icons
          2. The cuvette could be empty or contain your solvent, such as water

a. The instrument will subtract this blank spectrum from subsequent measurements

3.

 

word image 3260 word image 3261

4.

    1. Labquest will return to the screen with the main box.
  1. Measurements
    1. Full Spectrum
      1. To take a full spectrum, (380nm – 950 nm), press the play button on the screen, or the play hardware button.
        1. The Labquest will show a live spectrum. Click the Stop button, or the play hardware button again to freeze the spectrum
        2. word image 3262
        3. The Labquest will now show the latest spectrum, autoscaled.
        4. word image 3263
        5. To temporary store the run, click the file cabinet button.
          1. The spectrum disappears, and it is ready to take another measurement.
          2. word image 3264
    2. To view several spectra at the same time, click “Run #” next to the file cabinet button, then click “All Runs”
    3. To view the absorbance for a particular wavelength, click on the spectrum. The absorbance will be displayed on the right for each Run.

 

word image 3265

1. You can step to the previous or next wavelength by tapping the buttons at the bottom of the spectrum.

  1. Individual wavelength
    1. Take a full spectrum as above
    2. Return to the Sensor screen. iii. Click “Mode: Full Spectrum” iv. word image 3266
    3. Select Mode: “Time Based” Click “OK”.
    4. word image 3267
    5. Labquest will ask if you want to save previous data (your previous full spectra), click save and enter a filename
    6. The sensor screen will now show a live absorbance @ 380.5 nm
    7. word image 3268
    8. You can change the wavelength by click the red box and “Select

Wavelength…”. Enter your desired wavelength

  1. Individual wavelength kinetics.
    1. Set up the individual wavelength measurement as above.
    2. You can change the acquisition parameters by clicking the Mode/Rate/Duration box.
    3. To begin the measurement, begin the reaction in your cuvette, insert the cuvette and press the play button.
    4. The graph tab will open and your live data will graph as it acquires.
        1. You can stop the acquisition early by pushing the stop button, or wait until it reaches the defined endtime.
        2. word image 3269

v.

        1. Save data
      1. *The file cabinet does NOT save the data. It only temporarily stores it to compare to other runs*
      2. To save all open data to the device for easy opening later, click File -> Save…
          1. Click on “untitled” to set a filename for your dataset
          2. Click “Save”

c. To save all open data to a flash drive, insert a USB flash drive into the Labquest

          1. Click File -> Export…
          2. Verify that a USB icon is in the upper left corner. iii. Click on “untitled” to set a filename for your dataset

iv. Click “Save”

4. Shutdown

      1. Unplug the SpectroVis USB cable
      2. Click File -> Quit
      3. Push the power button on top of the Labquest

 

  1. “Keeping a Laboratory Notebook”, A. Eisenberg; J. Chem. Educ., 59 (12), 1045 (1982).

  2. Adapted from: Machado, C.; Chachaldo, V.G. “An Easy and Versatile Experiment to Demonstrate Solvent Polarity Using Solvatochromic Dyes”, J. Chem. Ed. 1992, 78, 649-651.

  3. + + 2e−  Sn Eo = 0.14 V Sn

  4. + + 3e−  Al Eo = 1.66 V Al

    To construct a complete electrochemical cell with a positive overall potential, it will be necessary to consider the aluminum half-cell not as a reduction, but as an oxidation. This requires writing the half-cell reaction shown above in reverse. It also requires changing the cell potential magnitude and sign to equal, but opposite, that of the reduction half-reaction. The new half-reactions, and the overall reaction are shown below.

  5. 3+ + 3e) x 2 Eo = +1.66 V

    Oxidation (anode): (Al  Al

  6. + + 2e−  Sn) x 3 Eo = 0.14 V Reduction (cathode): (Sn

SOLUBILITY PRODUCT CONSTANT LAB RAW DATA

Raw Data for the Titration of Calcium Iodate with Sodium Thiosulfate Scout Titration

Volume of Calcium Iodate (mL) 5

Concentration of Sodium Thiosulfate (M) 0.1003

Initial Buret Reading (mL) 10

Final Buret Reading (mL) 16

Raw Data for the Titration of Calcium Iodate (w/added Ca2+ ions) with Sodium Thiosulfate Scout Titration

Volume of Calcium Iodate (mL) 5

Concentration of Sodium Thiosulfate (M) 0.1003

Initial Buret Reading (mL) 16.93 Final Buret Reading (mL) 20.1

Trial One Trial Two

5 5

0.1003 0.1003

16 23

21.95 29.2

Trial One Trial Two

5 5

0.1003 0.1003

20.1 23.5

23.5 26.7

SOLUBILITY PRODUCT CONSTANT LAB RAW DATA

Raw Data for the Titration of Calcium Iodate with Sodium Thiosulfate Scout Titration

Volume of Calcium Iodate (mL) 5

Concentration of Sodium Thiosulfate (M) 0.1003

Initial Buret Reading (mL) 10

Final Buret Reading (mL) 16

Raw Data for the Titration of Calcium Iodate (w/added Ca2+ ions) with Sodium Thiosulfate Scout Titration

Volume of Calcium Iodate (mL) 5

Concentration of Sodium Thiosulfate (M) 0.1003

Initial Buret Reading (mL) 16.93 Final Buret Reading (mL) 20.1

Trial One Trial Two

5 5

0.1003 0.1003

16 23

21.95 29.2

Trial One Trial Two

5 5

0.1003 0.1003

20.1 23.5

23.5 26.7

Order a unique copy of this paper
(550 words)

Approximate price: $22