Lab 13-Titration of Hydrogen Peroxide

Physical Sciences Department

General Chemistry II CHM 241

LABORATORY MANUAL

Second Edition

TABLE OF CONTENTS

Page

Course Schedule 4

Laboratory Safety 5

Integrity of Data Guidelines 8

Experiment 1. Heat of Fusion 9

Experiment 2. Intermolecular Forces 14

Experiment 3. Spectroscopy 23

Experiment 4. Percent Copper in Brass 31

Experiment 5. Freezing Point Depression 39

Experiment 6. Chemical Kinetics 50

Experiment 7. Le Châtelier’s Principle 63

Experiment 8. Coordination Number 75

Experiment 9. Identification of a Weak Acid 83

Experiment 10. Solubility Product 94

Experiment 11. Qualitative Analysis of Cations 101

Experiment 12. Titration of Hydrogen Peroxide 109

Experiment 13. Electrochemistry 119

Experiment 14. Bicarbonate-carbonate mixture 128

Appendix A. Common Laboratory Equipment 135

Appendix B. Volumetric Glassware 136

Appendix C. Graphing 137

Appendix D. Titration 139

Appendix E. Filtration 140

Appendix F. Periodic Table 141

General Chemistry II Laboratory

COURSE SCHEDULE

Fifteen Week Semester Twelve Week Semester

1. Lab Safety and Exp. 1 1. Lab Safety and Exp. 1

2. Exp. 2 2. Exp. 3 and Exp. 2, Parts A&B (or Part C)

3. Exp. 3 3. Exp. 4 and Exp. 2, Part C (or Parts A&B)

4. Exp. 4 4. Exp. 5

5. Exp. 5 5. Exp. 6

6. Exp. 6 6. Exp. 7 and Exp. 8 part A

7. Exp. 7 7. Exam 1 and Exp. 8 part B

8. Exam 1 and Exp. 8 part A 8. Exp. 9 and Exp. 8 part C

9. Exp. 8 parts B and C 9. Exp. 10*

10. Exp. 9 10. Exp. 11

11. Exp. 10* 11. Exp. 12

12. Exp. 11 12. Exam 2 and Exp. 13

13. Exp. 12

14. Exp. 13 or Exam 2

15. Exam 2 or Exp. 13

*The NaOH solution standardized in Experiment 9 is used again in this experiment.

SAFETY REGULATIONS FOR THE CHEMISTRY LABORATORY

1. Read these safety regulations carefully and be sure you understand them. Before each

laboratory session, your instructor will discuss any safety hazards that may be associated

with that day’s experiment. Therefore, it is imperative that you come to lab on time.

2. Due to safety concerns students who arrive after the pre-lab presentation may not be
allowed to perform that particular lab experiment.

3. It is strongly suggested that you obtain a hall locker from the Security Office. Only your
lab manual, notebook, and calculator are allowed on the lab bench.

4. Report all accidents, no matter how minor, to your instructor at once. No one in the lab is
permitted to give out bandages or medication. You must see the College Nurse.

5. Safety glasses or goggles are required and must be worn by everyone in the lab when
experiments are being conducted. Contact lenses are not recommended in the chemistry

lab. Safety glasses are provided by the college, but students may purchase their own.

6. Do not perform any unauthorized experiment.

7. Do not taste anything in the laboratory. Never eat, drink or smoke in any of the labs.

8. You must tie back long hair. Do not wear open-toed shoes, shorts, fuzzy sweaters, loose
sleeve shirt or any dangling jewelry. You must cover bare midriffs. You are advised to

wear a lab coat or old clothing to the lab.

9. Do not fill pipettes by mouth. Rubber bulbs or pipette pumps are provided. The
instructor will demonstrate how these are to be used.

10. Exercise care when noting the odor of fumes. Use ‘wafting’ if you are directed to note
an odor.

11. Do not force glass tubing or a thermometer into rubber stoppers. Lubricate with water
and introduce it gradually and gently into the stopper, or insert through a cork borer.

Protect your hands with toweling when inserting without a cork borer.

12. Never point a test tube containing a reaction mixture (especially when heating) toward
yourself or another person.

13. No ‘fooling around’ in the laboratory. A less than serious approach to lab work may
result in an accident.

14. Before connecting or disconnecting electrical equipment, make sure that the switches
are in the off position.

15. Never work in the laboratory alone.

16. Make sure all apparatus is properly supported on the workbench.

17. Read the label on every bottle twice before using it in the laboratory. Many chemical
names are very similar but are very different chemically.

18. Replace caps and stoppers on bottles immediately. Return spatulas to their correct place
immediately after use. Do not mix them up.

19. Do not remove or relocate any chemical that has been placed in the hood. Sample it in
the hood.

20. Never light a Bunsen burner with a cigarette lighter. Use the strikers that are provided.

21. Students are responsible for keeping their work area neat and orderly. All spills are to be
cleaned up immediately using the spill kits located on the instructor’s desk. Solid

chemical waste should be and placed in the appropriately labeled container. Liquid

chemical waste should be poured into the appropriately labeled container. All waste

material should be left in the hood for subsequent disposal. If there is doubt about proper

disposal, ask the instructor.

22. Wash all glassware immediately after use. Place clean glassware on drying rack or in the
designated bin on the counter.

23. Dispose of broken glassware in the labeled broken glassware boxes.

24. Wash your hands before leaving the laboratory.

25. You must notify your instructor of any chemical to which you are allergic.

26. If you are pregnant or planning to become pregnant this semester, you must notify your
physician that you are enrolled in a chemistry lab course. You and your physician must

decide whether or not it is appropriate for you to remain in the course.

Note the location of the following safety equipment so that you can get to it quickly

in an emergency.

SAFETY EQUIPMENT LOCATION

FIRE EXTINGUISHER

SAFETY SHOWER

EYEWASH

EMERGENCY PHONE

FIRE ALARM

NEAREST EXIT

True False

1. Safety glasses must be worn by everyone working in the lab. T F

2. Only major accidents in the lab need to be reported T F

3. Material Safety Data Sheets are provided in the lab T F

4. Eating and drinking are permitted in the lab T F

5. It is OK to taste a chemical as long as it smells good T F

6. Only authorized experiments are to be performed T F

7. You should wear shoes at all time in the lab T F

8. In order to save time, it is permissible to weigh hot objects T F

9. Broken glassware should be disposed of in the appropriate box T F

10. Working alone in the lab is an acceptable practice T F

A typical Chemistry Laboratory safety YouTube video link is given below: (hold Ctrl Key

and hover the mouse over the link) https://www.youtube.com/watch?v=UKovNdse5MU

Please complete sign this attached form. Remove it from the safety regulations and hand it to

your Laboratory Instructor.

I, the undersigned, have read the Divisional Safety Regulations for the Chemistry

Laboratories. I understand them and will abide by them.

Print your name: ________________________________________________________

Signature: ____________________________________________________________

Date: ____________________________________________________________

Course Name and Number:_______________________________________________

INTEGRITY OF DATA GUIDELINES

One purpose of a laboratory course is to reinforce the concepts covered in the lecture

course. A second, equally important purpose, is to experience working in a chemistry lab, and

to learn about practices and procedures that are employed in such an environment. In addition

to specific laboratory procedures that will be covered in the array of experiments, there are

two universal practices in all laboratory settings- Laboratory Safety, which was discussed in

the previous pages, and Integrity of Data Guidelines.

These guidelines are used in all laboratory settings- from the traditional research

laboratory to hospitals and the physician’s office. The purpose of the guidelines is to ensure

that data is recorded in such a way that its veracity, or authenticity, cannot be questioned.

Taken as a whole, these practices protect the integrity of the data by preventing it from being

changed or recorded in error. Students are expected to follow these integrity of data guidelines

when collecting and recording data. The guidelines are as follows:

1. Data sheets must include the date and the student’s name.

2. Data is recorded in blue or black non-erasable ink; no white-out is permitted.

3. If a mistake is made while entering data, a single line is used to cross it out

and the correct entry is made nearby. (The original entry must be legible.)

4. No transcription is permitted. (Data is recorded directly into the data sheets.)

5. Data is recorded at the time it is observed.

In most laboratories today, notebooks are electronic rather than paper. Although this

renders a different set of guidelines, their purpose is the same- to ensure the authenticity of

data. Laboratory notebook software does not permit a change to be made once data has been

entered. In instances where a change is required, there is a record of the original entry. When

a measurement is recorded on a scrap of paper, that original data is scanned and becomes a

part of the notebook. These and other practices concerning electronic lab notebooks, along

with the guidelines described above regarding paper notebooks, work together to protect the

integrity of experimental data.

General Chemistry II Laboratory

Experiment 1

Heat of Fusion

OBJECTIVE: To determine the heat of fusion of water.

BACKGROUND:

When the solid phase of a molecular substance is converted to the liquid phase, energy,

in the form of heat, must be added in order to break the attractions between the molecules.

These intermolecular forces in a solid hold the molecules locked into position. Although the

molecules vibrate in place, they do not move relative to each other, i.e. they have no

translational movement. In contrast, the molecules in the liquid phase, although close to one

another, do have translational movement. They are constantly making and breaking

intermolecular attractions as they move about in random translational motion.

As heat is added to a molecular substance in the solid phase, the kinetic energy of the

molecules increases, resulting in greater vibrational motion, and evidenced by an increase in

temperature. This process continues until the melting point is reached, when molecules begin

to have sufficient energy to break the attractive forces holding them in position, and the

substance begins to melt. At this point, added energy results in breaking attractive forces rather

than in increased movement, and the temperature remains constant throughout the melting

process. When the entire sample has become a liquid, added heat increases the kinetic energy

and the temperature increases once again.

A similar transition occurs in converting a substance from the liquid to the gas phase.

As heat is added once the boiling point has been reached, this energy is used to break

intermolecular forces between molecules in the liquid phase. Again, during the process of

vaporization, the temperature remains constant.

These relationships can be summarized in a heating curve, as illustrated in the figure

on the following page.

The amount of heat required to convert a substance from the solid to the liquid phase

is quantified as the heat (or enthalpy) of fusion, ∆Hfus. It is a physical property, and can be

reported as heat per gram of substance or per mole of substance. The latter is often referred to

as the molar heat of fusion.

Experiment 1 Heat of Fusion

General Chemistry II Laboratory

In this experiment, the heat of fusion of water, in joules/gram, will be determined using

a coffee cup calorimeter where a sample of ice has melted in tap water. The amount of heat

given up by the tap water in the calorimeter as it cools (qwater) will be absorbed as heat by a

sample of ice as it melts (qfusion) and as this melted ice warms (qmelted ice). Assuming no loss of

heat to the surroundings, the sum of these must equal zero.

qwater + qfusion + qmelted ice = 0

Rearranging,

qfusion = – qwater – qmelted ice (eq. 1)

Values for both qwater and qmelted ice are obtained from the following equations, where m

represents mass, c represents the specific heat of water (4.18 J/g ºC) and ∆T represents the

change in temperature.

q = m c ∆T (eq. 2)

∆T = Tfinal – Tinitial (eq. 3)

Once the heat of fusion is determined, the experimental error can be found as follows.

𝑃𝑒𝑟𝑐𝑒𝑛𝑡 𝐸𝑟𝑟𝑜𝑟 =
|𝑇ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 𝑉𝑎𝑙𝑢𝑒 − 𝐸𝑥𝑝𝑒𝑟𝑖𝑚𝑒𝑛𝑡𝑎𝑙 𝑉𝑎𝑙𝑢𝑒|

𝑇ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 𝑉𝑎𝑙𝑢𝑒
× 100

Three determinations will be made using different sample sizes. A graph of the heat,

in joules, absorbed in melting the ice (qfusion) as a function of the mass of melted ice, in grams,

will be constructed. The slope of the line represents the heat of fusion of water.

T
em

p
er

at
u
re

Heat added

melting

vaporization

Figure 1: Heating Curve

(eq. 4)

Experiment 1 Heat of Fusion

General Chemistry II Laboratory

REAGENTS: Ice EQUIPMENT: 150-mL beaker
Tap water 100-mL graduated cylinder

coffee cup calorimeter
400-mL beaker

thermometer, or thermocouple

PROCEDURE:

1. Measure and record the mass of a 150-mL beaker. Set it aside ready to use in step 5.

2. Tare a coffee cup calorimeter. Add 100 mL tap water using a graduated cylinder.

Measure and record the mass. Place the calorimeter in a 400-mL beaker for stability.

3. Measure and record the initial temperature of the tap water in the calorimeter.

4. Add sufficient ice to fill the volume of water, and gently stir with the thermometer.

5. When the temperature reaches between 0ºC and 5ºC, record the final temperature.

Immediately pour the water into the 150-mL beaker, leaving the unmelted ice behind.

6. Measure and record the mass of the beaker and contents.

7. Repeat steps 1 – 6 using 70 mL tap water, and again using 40 mL.

Disposal: Water may be disposed of down the drain.

CALCULATIONS:

A. Perform the following calculations for each of the three determinations.

1. Determine the mass of the contents of the beaker. This is the mass of the original tap water

in the calorimeter, plus that of the melted ice.

2. Determine the mass of the melted ice by subtracting the mass of the tap water from the

mass of the beaker contents.

3. Determine the temperature change for the tap water, ∆Twater, using eq. 3.

4. Determine the temperature change for the melted ice, ∆Tmelted ice, using eq. 3. The initial

temperature for the ice is assumed to be 0 ºC.

5. Determine qwater using the mass and temperature change of the tap water and eq. 2.

6. Determine qmelted ice using the mass and temperature change of the melted ice and eq. 2.

7. Determine qfusion using eq. 1.

B. Prepare a graph in Excel* of qfusion, in joules, as a function of the mass of melted ice, in

grams. Determine the heat of fusion for water from the graph.

*See Appendix C for directions on graphing.

 General Chemistry II Laboratory

Date: ____________________ Name: ________________________

Experiment 1: Heat of Fusion

Data: Determination 1 Determination 2 Determination 3

Initial mass of beaker ____________ ____________ ____________

Mass of tap water ____________ ____________ ____________

Initial temperature ____________ ____________ ____________

Final temperature ____________ ____________ ____________

Final mass of beaker ____________ ____________ ____________

Results:

Mass of beaker contents ____________ ____________ ____________

Mass of melted ice ____________ ____________ ____________

ΔT of tap water ____________ ____________ ____________

ΔT of melted ice ____________ ____________ ____________

Heat for water, qwater ____________ ____________ ____________

Heat for melted ice, qmelted ice ____________ ____________ ____________

Heat for fusion of ice, qfusion ____________ ____________ ____________

Heat of fusion of ice _______________________

General Chemistry II Laboratory

Date: ____________________ Name: ________________________

Experiment 1: Heat of Fusion

POST-LAB QUESTIONS:

1. The heat of fusion of water is 333 J/g. Determine the percent error using equation 4.

2. Determine the amount of heat required to raise the temperature of a 22.5-gram

sample of copper from 125 ºC to its melting point of 1084 ºC, and then melt the

copper. (The specific heat of copper is 24.4 J/mol ºC and its heat of fusion is

13.26 kJ/mol.)

3. If the ice had begun at a temperature lower than 0 ºC, would the calculated value of

the heat of fusion have been higher, lower, or unchanged? Briefly explain.

General Chemistry II Laboratory

Experiment 2

Intermolecular Forces

OBJECTIVE: To relate intermolecular forces of molecules to physical properties.

BACKGROUND:

The attractive forces between molecules and their neighbors are called intermolecular

forces. These forces are much weaker than the intramolecular forces within a substance- the

covalent (or ionic) bonds. Intermolecular forces are the attractions that need to be overcome

for a molecular solid to melt, and for a liquid to vaporize. Therefore, the strength of these

attractive forces influence a substance’s physical properties. The stronger the intermolecular

forces, the higher melting point, boiling point, heat of vaporization, and other properties. The

following table summarizes the boiling points of some molecular compounds.

Compound Formula Polarity Molecular Structure Boiling Point

Methane CH4 Nonpolar

– 161oC

Propane C3H8 Nonpolar

– 42oC

Butane C4H10 Nonpolar

10oC

Hexane C6H14 Nonpolar

70oC

Acetone C3H6O Polar

56oC

Ethanol C2H6O Polar

77oC

Water H2O Polar

100oC

Experiment 2 Intermolecular Forces

General Chemistry II Laboratory

There are three general types of intermolecular forces. All substances exhibit London

Dispersion Forces (LDF), and they are generally the weakest of the three types. These London

forces are due to the attractions between small, temporary dipoles that arise from the constant,

random movement of the electrons in a substance. As molar mass increases, the size of the

electron cloud increases as well. It becomes more easily distorted, and produces temporary

dipoles of greater magnitude. This causes the attractions to be stronger, requiring more energy

for both fusion and vaporization. For halogens, this results in increasing melting and boiling

points, shown by the fact that at room temperature F2 and Cl2 are gaseous, Br2 is liquid and I2

is solid. The extent to which the electron cloud can be distorted is called polarizability.

Dipole-dipole forces exist between molecules that are polar. Since the dipoles are

permanent, these attractions are generally stronger than London Dispersion Forces. This

means that a polar molecule with similar molar mass as a nonpolar molecule will have higher

melting points and boiling points. Not all molecules containing polar bonds are polar. The

polar bonds must be unevenly dispersed in the molecule in order to produce a polar molecule.

CO2 and CBr4, for example, have polar bonds but are not polar molecules.

The third type of intermolecular force is hydrogen bonding, a specific type of dipole-

dipole attraction that is stronger than other dipole-dipole attractions. Hydrogen bonds form

when a hydrogen atom is covalently bonded to a very electronegative atom. This causes its

electron to be drawn away from its nucleus. The positive hydrogen is then attracted to the very

electronegative atom in a neighboring molecule. In order to observe hydrogen bonding, the

hydrogen atom must be covalently bonded to fluorine, oxygen or nitrogen. A hydrogen atom

bonded to a carbon atom cannot create a hydrogen bond. It’s

important to note that, despite its name, a hydrogen bond is

an intermolecular force, not a bond. The figure to the right

illustrates H-bonding between water molecules. H-bonding

is important in biochemistry; the structure of a biopolymer is

largely determined by the formation of hydrogen bonds.

The relative strengths of the three types of intermolecular forces, and thus boiling

points, are generally as follows:

London Dispersion Forces < Dipole-Dipole Forces < H-Bonding

However, this is not always true. Since molar mass is also a factor, a large non-polar molecule

can have a higher boiling point than a compound that interacts with dipole-diploe forces, or

even a substance with H-bonding. For example, octane, a component of gasoline, has a boiling

point of 125oC- much higher than acetone (dipole-dipole) and H2O (H-bonding). This is due

to the polarizability of the large electron cloud.

To make comparisons of the intermolecular forces of a substance, evaporation rate can

be used instead of boiling point. Evaporation rate is the ratio of the change in temperature to

the change in time as a substance evaporates. A faster rate of evaporation translates to a lower

boiling point and, in turn, weaker intermolecular forces.

Experiment 2 Intermolecular Forces

General Chemistry II Laboratory

Boiling point is not the only physical property affected by the type of intermolecular

forces a substance has. Solubility is also dependent upon the polarity of a molecule. The term

“like dissolves like” suggests that polar solutes dissolve in polar solvents and nonpolar solutes

dissolve in nonpolar solvents. Therefore, polarity, and the associated intermolecular forces,

determine a substance’s solubility in water and in other solvents.

The solubility of a solid in a liquid is readily observed. When liquids mix forming a

homogeneous solution, they are said to be miscible; if they do not mix, they are immiscible.

If two liquids are miscible, there is no observable interface between the two. If the two liquids

are immiscible, two distinct layers are seen.

In this experiment, both miscibility and evaporation rates of acetone, ethanol, hexane,

and water will be determined. Salt solubility in an ethanol-water mixture will also be observed.

REAGENTS: acetone EQUIPMENT: Thermometers
ethanol filter papers

hexane rubber band

distilled water tape

sodium chloride stop watch

5 test tubes containing a 50% by volume: small test tubes

water and acetone wood block

water and hexane

hexane and acetone

hexane and ethanol

ethanol and acetone

SAFETY ALERT:

– Do not pour any materials into the sink!

– Wash hands and laboratory bench after the experiment.

– Acetone: Extremely flammable liquid and vapor. Vapor may cause flash fire. Causes eye

irritation. Breathing vapors may cause drowsiness and dizziness. Causes respiratory tract

irritation. Aspiration hazard if swallowed. Can enter lungs and cause damage. Prolonged or

repeated contact may dry the skin and cause irritation.

Hexane: Extremely flammable liquid and vapor. Vapor may cause

flash fire. Breathing vapors may cause drowsiness and dizziness.

Causes eye, skin, and respiratory tract irritation. May be harmful if

absorbed through the skin. Aspiration hazard if swallowed and enters

lungs causing damage. Possible risk of impaired fertility. Long-term

exposure may cause damage to the nervous system of the extremities.

General Chemistry II Laboratory

Date: ____________________ Name: ________________________

Experiment 2: Intermolecular Forces

PRE-LABORATORY QUESTIONS

1. Which of the substances used in this experiment must be handled in the fume hood?

2. Identify the strongest type of intermolecular forces in acetone, ethanol, water and hexane.

(Structures listed on page 15.)

3. Predict the relative strength of the intermolecular forces in the four liquids above.

______________ < _______________ < …

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