Potentiometric Acid Base Titrations Lab Report

 

CHEM25770

Analytical Chemistry 1

 

Laboratory Manual

Fall

2020

Table of Contents

 

Laboratory 1

Determination of KHP in a Sample………………………………………………………………………. 3

 

Laboratory 2

Potentiometric Acid – Base Titrations…………………………………………………………………….10

 

Laboratory 3

Gravimetric Determination of Chloride in a Soluble Sample………………………………………….25

 

 

General Instructions for Labs

 

 

You must make sure that you are on time for your lab. During the first few minutes of the lab, there will be a pre-lab quiz and then your lab instructor will give you verbal instructions for that week’s experiment and may demonstrate some procedures. If you arrive at the lab too late to do the quiz, you will not have an opportunity to make it up. If you miss the instructor’s explanation, you will not know all the details of what you need to do, and you will miss important safety information. Any student who arrives at the lab more than 5 minutes after the start of the lab period will not be admitted to the lab – NO EXCEPTIONS. There will not be an opportunity to make up a missed lab.

 

Anyone who is in the laboratory must wear an appropriate lab coat, safety glasses, and flat, closed shoes at all times (no sandals or flip-flops). You must also wear the PPE necessary for COVID-19, such as face masks and gloves. Social distancing must be followed. You will have to leave the laboratory if you do not meet these requirements.

 

CHEM25770 2 Fall 2020

Laboratory 1

Determination of KHP in a Sample

Objective

A solution of sodium hydroxide will be prepared and then standardized using the primary standard, potassium hydrogen phthalate (KHP). Then using the standardized NaOH solution, the concentration of KHP in an impure sample will be determined by titration.

Background

  • NaOH is commonly used in the preparation of standard solutions. Standardization of the solution is necessary after preparation as the NaOH used is not highly pure, it is hygroscopic in nature and also reacts with CO2 from the atmosphere. These shortcomings can be overcome by the use of 50% NaOH solution.
  • Accurate measuring apparatus such as the volumetric flask and pipette are to be used when a solution’s concentration must be known exactly such as with a stock or standard solution; approximate measuring apparatus such as graduated beakers, cylinders and reagent bottles suffice when only approximate concentrations are needed.
  • The 0.1 M NaOH solution is stored in a polyethylene bottle as the base reacts with glass, forming silicates.
  • Potassium hydrogen phthalate (KHP) is used as a primary standard to standardize the

0.1 M NaOH.

  • Reverse osmosis (RO) water must be used for all chemical procedures as tap water would cause errors in your results.

Calculations

Standardization:

 

Standardization is a way of accurately finding the concentration of a titrant solution. So, for example, we may say that a titrant solution is 0.1 M NaOH, and that is its approximate concentration. But in order to be useful as a titrant, its concentration needs to be known more precisely, usually to 4 significant figures, e.g., 0.1089 M or 0.09753 M. Standardization of a titrant is sometimes done by using that titrant to titrate a precisely known mass of a primary standard. That is how it is done in this experiment.

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CHEM25770 3 Fall 2020 At the endpoint of the titration, the number of moles (nB) of base titrant (0.1 M NaOH) that you have added is equal to the number of moles (nA) of acid (KHP) that was in the flask.

𝑛𝑛𝐴𝐴 = 𝑛𝑛𝐵𝐵 (1)

The number of moles of acid is equal to the weight of pure KHP that you weighed out divided by the molar mass of KHP.

𝑤𝑤𝑤𝑤𝑤𝑤𝑤𝑤ℎ𝑡𝑡 𝑜𝑜𝑜𝑜 𝐾𝐾𝐾𝐾𝐾𝐾

𝑛𝑛𝐴𝐴 = 𝑀𝑀𝑀𝑀 𝑜𝑜𝑜𝑜 𝐾𝐾𝐾𝐾𝐾𝐾 (2)

Since the molarity of the base = (no. of moles)/volume, the number of moles of base is equal to the molarity of the base times the volume of the base:

𝑛𝑛𝐵𝐵 = 𝑀𝑀𝐵𝐵𝑉𝑉𝐵𝐵 (3) Combining equations (1), (2), and (3), we get

𝑤𝑤𝑤𝑤𝑤𝑤𝑤𝑤ℎ𝑡𝑡 𝑜𝑜𝑜𝑜 𝐾𝐾𝐾𝐾𝐾𝐾

𝑀𝑀𝑀𝑀 𝑜𝑜𝑜𝑜 𝐾𝐾𝐾𝐾𝐾𝐾 = 𝑀𝑀𝐵𝐵𝑉𝑉𝐵𝐵 (4)

Because you know all of the values in this equation except 𝑀𝑀𝐵𝐵, the molarity of the base, you can solve the equation to find 𝑀𝑀𝐵𝐵.

You will prepare the NaOH solution in such a way that its molarity is about 0.1 M. The purpose of the standardization (when you titrate pure KHP with your NaOH solution) is to find out what the NaOH molarity is exactly. You must take all of your readings (weight of KHP, volumes of titrant) to 4 or more significant figures. This will allow you to calculate the NaOH molarity to 4 significant figures.

 

Determination of the unknown:

After you have done the calculations for the standardization, you will know the exact concentration of your NaOH solution. It will be around 0.1 M, but the purpose of the standardization is to find out what it is exactly (to 4 significant figures). In the second week of the experiment, you use your standardized NaOH to titrate your unknown sample. Your unknown contains KHP, but it is not pure KHP. Use equation (4) to find the weight of KHP in your unknown. Divide that weight of KHP by the weight of the sample (that you titrated) to find the percentage of KHP in your unknown.

Chemicals Required

Week 1:

50% w/w NaOH

KHP primary standard

Phenolphthalein indicator

Week 2:

Phenolphthalein indicator

Standardized NaOH from the previous week

KHP unknown

 

Equipment Required

Week 1:

10 mL graduated cylinder

1 L volumetric flask

1 L polyethylene bottle

50 mL burette

3-250 mL Erlenmeyer flasks

Burette stand

Burette cap

 

Procedure

Week 1

Week 2:

3-250 mL Erlenmeyer flasks

Burette stand

50 mL burette

Burette cap

Wash bottle

 

Preparation of 0.1 M NaOH solution

  1. Wash a 1 L volumetric flask thoroughly and rinse it with RO water. Add RO water to the volumetric flask until it is about half full.
  2. Using a graduated cylinder, pour 6 mL of 50% NaOH solution into the 1 L volumetric flask.
  3. Add RO water slowly to the flask up to the base of the neck and then bring it up to the calibration mark. You do not have to get it exactly to the calibration mark. Why not?
  4. Invert the flask 12 – 15 times to ensure good mixing of the solution.
  5. Condition a clean 1 L polyethylene bottle by pouring about 5 – 10 mL of the 0.1 M NaOH solution into the bottle and swirling it around, then discarding the liquid in the sink, while running the tap water. Repeat this with another 5 – 10 mL of the 0.1 M NaOH. Store the rest of the NaOH solution in the polyethylene bottle. Label it properly with contents, concentration, date of preparation, and your initials.
  6. Squeeze as much air as possible out of your polyethylene bottle after pouring the NaOH solution in and tightly cap the bottle.

Standardization of 0.1 M NaOH solution

  1. The primary standard KHP will be dried for 2 hrs at 110°C and cooled in a desiccator ready for use.
  2. Weigh accurately (to the nearest 0.1 mg) about 0.7 g of primary standard KHP into a clean and dry 250 mL Erlenmeyer flask. (If you have washed the flasks, dry them with paper towel.) Do this two more times so that you have three Erlenmeyer flasks, each containing about 0.7 g KHP, accurately weighed. (Why must the flasks be dry before you weigh the KHP into them?) Add 50 – 75 mL of RO water to each flask and swirl to dissolve.
  3. Condition a clean 50 mL burette with the 0.1 M NaOH solution that you prepared.
  4. Fill the burette with the 0.1 M NaOH solution to a level close to 0.00 mL, but do not waste time trying to get it exactly to 0.00. Record the initial volume to two decimal places, for example 0.83 mL. It is not good enough to record it as 0.8 mL. Since NaOH can absorb CO2 from the air, your burette should be capped.
  5. Take one of the flasks containing the weighed KHP and add 3 – 4 drops of phenolphthalein indicator. Titrate with NaOH until a very faint pink endpoint. The endpoint should not have a strong colour. The endpoint occurs when you can see the weakest possible pink colour and that colour persists for at least 30 – 45 s. Record the final volume to two decimal places, for example, 35.17 mL.
  6. Repeat steps 4. and 5. for the other two flasks containing primary standard KHP.
  7. Calculate the molarity of NaOH from each titration. Determine the mean molarity, the standard deviation, and the % relative standard deviation. Use the mean molarity of NaOH to determine the % KHP in your unknown sample.

Week 2

Determination of KHP in an Impure Sample

  1. Your instructor or technologist will provide you with your unknown sample (which has been dried by the lab technologist). Immediately record in your lab notebook the number of your unknown sample.
  2. Weigh out 0.8 g (to the nearest 0.1 mg) of your unknown KHP sample into each of three 250 mL Erlenmeyer flasks and dissolve in about 75 mL of RO water.
  3. Fill the burette with the 0.1 M NaOH solution to a level close to 0.00 mL, but do not waste time trying to get it exactly to 0.00. Record the initial volume to two decimal places. Since NaOH can absorb CO2 from the air, your burette should be capped.
  4. Take one of the flasks containing the weighed unknown KHP and add 3 – 4 drops of phenolphthalein indicator. Titrate with NaOH until a very faint pink colour of the indicator persists for at least 30 – 45 s. Record the final volume to two decimal places.
  5. Calculate the percent by weight of KHP in your unknown sample. Make sure you have 3 good results from your titrations.
  6. When you have finished with your sample, return the remaining unknown to its original container and place in the box near the whiteboard.

 

Lab Report Format

 

This lab report will include:

 

COVER SHEET

TITLE

PURPOSE

OBSERVATIONS

CALCULATIONS

RESULTS

CONCLUSION

POSSIBLE SOURCES OF ERROR

REFERENCES

PHOTOCOPY OF ALL DATA FROM LAB NOTEBOOK

 

COVER SHEET

The cover sheet is to be completely filled in and attached to the front of the report. If information such as sample number, mean and standard deviation are missing, the analytical results cannot be marked, resulting in a mark of 0 for that part of the lab.

 

TITLE

 

The title should be the same as that in your lab manual.

 

PURPOSE

 

The purpose is to be 1 – 2 short sentences describing the purpose of your lab.

 

OBSERVATIONS

 

This section includes both physical and numerical observations.

 

Physical observations such as colour change may be written in a short paragraph which includes the change in colour of the indicator and the rate of colour change. Any in between changes must be noted.

Numerical observations are to be put into table form. The table must have an appropriate title and sub headings.

 

Each week’s work is to be summarized in 1 table only.

 

Appropriate headings are:

 

Sample number

Sample mass

Initial volume of titrant

Final volume of titrant

Net titrant volume

 

CALCULATIONS

 

At the beginning of this section, the name, chemical formula, and molar mass of each substance involved in the titration are to be listed. All measurements in this lab are done to four significant figures, and all calculations must be done to four significant figures.

 

The balanced equation for the reaction must be present immediately before any calculations.

 

A sample calculation is to be done for one sample only.

 

Use the average molarity of the NaOH when calculating the % KHP in your unknown sample. Calculate the % KHP as the final result for each of the three determinations. Calculations for mean, standard deviation, and % relative standard deviation are included here.

 

RESULTS

 

The final results for the unknown must be presented in table format. Two results tables are required, one for the standardization of the titrant and one for the analysis of the unknown.

The first table should show the three values found for the molarity of the titrant, the mean (with units), the standard deviation (with units), and percent relative standard deviation (with units).

The second table should show the unknown sample number, the three values found for %KHP, the mean (with units), the standard deviation (with units), and percent relative standard deviation (with units).

 

CONCLUSION

The Conclusion is a sentence which states “Sample number ______ was found to contain ____ % KHP with a relative standard deviation of ____%.”

 

POSSIBLE AND ACTUAL SOURCES OF ERROR

At least 3 sources of error are to be discussed here with their effect on the % KHP in the unknown.

 

REFERENCES

List all references used in report. All web sites must be easily accessed by the instructor or they will be deemed invalid.

 

Any diagrams, charts, pictures, tables, etc, downloaded from the web or copied from a book must be specifically referred to in your text and have a reference site under the diagram.

Marks will be deducted for lack of proper references.

Some or all marks will be deducted for plagiarism. Copying from a reference is deemed plagiarism. When using references to present your answers, use your own words.

 

Laboratory 2

Potentiometric Acid – Base Titrations

Objective

To obtain titration curves by titrating a solution of phosphoric acid with a standardized strong base and using these titration curves to determine of the concentration of the acid. The curves will be obtained in two ways, manually, and with an automated titrator.

Background

This experiment will illustrate the titration curve obtained when titrating a triprotic weak acid with a strong base. Data obtained from these types of titrations can be used for standardizing a solution, analyzing an unknown acid or base sample, and for determining the dissociation constants (pKas and pKbs) of weak acids and bases. They are called potentiometric titrations.

Each pair of partners will receive a solution containing an unknown concentration of phosphoric acid. This is a two week experiment and it will be done in two parts, one part each week. In one part of the experiment, the students will titrate the unknown solution manually using a regular burette and a pH meter. In the other part of the experiment, the students will titrate the same unknown solution using an automated titrator. Titrations on the automated titrator will be run in triplicate.

The titration curve that you get in this experiment will look something like this. The horizontal axis shows volume of titrant added and the vertical axis represents pH.

word image 2479

You can see that the pH rises slowly at first, then increases more steeply as it approaches the first end point. It then increases slowly again after the first end point. There will be a second end point at a volume of titrant which is twice the volume at the first end point.

When you are doing the manual titration, you will collect data points which are (volume, pH). When you are measuring pH in a part of the curve that is rising slowly, you need to measure

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CHEM25770 10 Fall 2020 the pH only every 1 mL or so. When you get close to an endpoint, you must measure the pH at volumes that are closer together, because if you don’t, you will miss the endpoint.

Chemicals Required

Standardized 0.1 M NaOH (prepared for you) Phosphoric Acid unknown solution

RO water

Buffers (pH 4, 7, and 10)

 

Equipment Required

Manual Titration: Automated Titration:

25 mL volumetric pipette 10 mL volumetric pipette

20 mL volumetric pipette Mettler-Toledo G20 automated titrator

250 mL volumetric flask Laptop computer

pH meter and pH combination electrode Wash bottle

50 mL burette 400 mL beaker

400 mL beaker

Wash bottle

Magnetic stir bar

Hotplate/Stirrer

 

Procedure Week 1

 

Dilution of Unknown Solution

 

1. Obtain an unknown sample of phosphoric acid. In your lab notebook, record which unknown you received. Pipet a 25.00 mL aliquot of the unknown acid sample into a 250.0 mL volumetric flask and dilute to the mark with RO water. Mix well. Transfer the diluted unknown to a clean, dry storage bottle. Label it and keep it in your drawer at the end of the first week of this experiment.

You will use this diluted unknown solution for both parts of this experiment (manual and automated).

 

Calibration of pH Meter

 

  1. pH combination electrodes should be conditioned for at least 8 hours in pH 7.00 buffer solution (soaking buffer). This has been done for you.
  2. The electrode must be filled to the plug hole with saturated KCI. Treat electrodes with extreme care – do not scratch the bottom
  3. Plug the pH meter into the power outlet and attach the combination electrode to the meter.
  4. The electrode is normally kept sitting in a pH 7 buffer. Buffers having pH 4 and 10 are also available in the lab. Calibrate the pH meter/electrode using each of the three buffers (pH 4, 7, and 10). Rinse off the electrode with your wash bottle and a waste beaker when changing from one buffer to another and before immersing the electrode in your solution.
  5. For each buffer, use the following steps to calibrate the pH meter.
    1. Press mode until you see pH (in the upper left corner).
    2. Press setup until you see clear buffer.
    3. Press enter to accept.
    4. Press std; it will say Standardize.
    5. When it shows stable, press std again.

 

Manual Potentiometric Titration

 

  1. Fill and cap a burette with standardized 0.1 M NaOH. Record the exact concentration of the NaOH in your lab notebook. Adjust the level of NaOH in the burette so that it is exactly 0.00 mL at the start of the titration.
  2. Pipet a 20.00 mL aliquot of the diluted unknown sample into a 400 mL beaker. Place the beaker in the centre of a hotplate/stirrer. Make sure that the heat is turned off.
  3. Rinse the electrode thoroughly by spraying it from a wash bottle containing RO water while holding a waste beaker underneath. Immerse the electrode in the solution containing the sample.
  4. Add a magnetic stir bar to the solution and stir gently (about 200 – 300 rpm). Be careful not to allow the magnetic stir bar to hit the electrode.
  5. If the bottom of the electrode is not fully immersed in the sample, you can add RO water to the beaker to bring the solution level up until the electrode is properly immersed. It is all right to add RO water as long as you do it before you start the titration. Do not add RO water to the beaker once the titration has begun.
  6. Position the burette so that the reagent can be delivered without splashing.
  7. Measure the pH. Record the initial pH as well as the burette volume (0.00 mL).

NOTE: Allow time between additions of titrant for the pH to become constant within 0.05 pH units. Record the pH once it becomes stable. The solution should be stirred continuously.

  1. Make sure your burette is at 0.00 mL at the start of the titration. Begin the titration by adding about 1 mL of titrant. Close the stopcock on the burette. Record the exact volume on the burette. Wait for the pH reading on the pH meter to stabilize. Once it has stabilized, record the pH reading.
  2. Add another increment of titrant that is about 1 mL, record the exact volume on the burette, and record the pH once it has stabilized.
  3. Keep adding 1 mL increments and recording volume and pH after each increment until you see that there is a larger jump in the pH.
  4. Then add smaller increments (0.50 then 0.10 mL) as the first end point is approached.
  5. Finally add one-drop portions between each reading in the immediate vicinity of the first end point.

NOTE: As the end point is approached, the change in pH becomes large even when only small amounts of base are added.

  1. After the first end point is reached (indicated by a large change in pH) continue to add titrant, in portions that are the reverse of those added when approaching the end point, until little change in pH is observed.
  2. When you start to come close to a volume that is twice the volume of the first end point, you should again start adding increments of titrant that are smaller.
  3. Once you are past the second end point, you can increase the size of the increments.
  4. Continue to measure volume and pH until you have gone about 3 mL past the second end point.
  5. If you did not get enough data points around the end points, it may be necessary to repeat the titration. More data points close to the end points of the reaction will produce more accurate results. As in most other labs, part of your grade on this lab will depend on accuracy.
  6. Set the pH meter to Standby. Rinse and the electrode with RO water. Place the electrode in soaking buffer.

 

Week 2

Automated Titration

You will use a Mettler-Toledo G20 automated titrator to automatically run the same titration that you did manually in the other part of the experiment. The automated titrator is connected to a laptop computer, which will receive all of the data (volume and pH readings) for each titration. In this part of the experiment, instead of recording data in your lab notebook, you will e-mail the data files from the laptop computer to yourself and also to your lab instructor.

Start the Computer and Autotitrator

 

  1. There is bottle on the left side of the automated titrator which contains the titrant solution. The exact concentration of the titrant is written on a label on the bottle. Record this exact concentration in your lab notebook. Make sure that there is enough titrant in the bottle before you begin. The bottle should be at least 1/3 full before you begin.
  2. Turn on the laptop computer. After it has finished booting, find the icon for LabX software and double click on that to start the LabX software. (LabX is the name of the software on the laptop that receives data from the G20 automated titrator.)
  3. The LabX software on the computer will show PredefinedUser

Press OK

  1. The software will show a menu with choices such as Activation, New Instrument, etc.

Ignore these and press Close.

  1. Turn on the Mettler-Toledo G20 automated titrator by pressing the button on the right side at the front.
  2. After booting, the G20 will show

User Name PredefinedUser

  1. Press Login
  2. It may show PnP sensor DGi115-SC detected at input SENSOR Press OK

 

Calibrate the Electrode

 

9. Press Methods

10. Press Calib. Calibration pH-Sensor Calibration

  1. Open the rubber stopper on the side of the electrode.
  2. Hold a 400 mL waste beaker under the electrode and rinse the electrode using RO water from a wash bottle. Wipe the electrode with a Kimwipe.
  3. Immerse the electrode in pH 4 buffer.
  4. Press Start
  5. It will show

Number of samples 3

Add sample 1/3 Press OK.

  1. It will show

Add Sample 2/3. Do not press OK yet.

  1. Remove the electrode from the pH 4 buffer. Hold a 400 mL waste beaker under the electrode and rinse the electrode using RO water from a wash bottle. Wipe the electrode with a Kimwipe.
  2. Immerse the electrode in pH 7 buffer.
  3. Press OK.
  4. It will show

Add Sample 3/3. Do not press OK yet.

  1. Remove the electrode from the pH 7 buffer. Hold a 400 mL waste beaker under the electrode and rinse the electrode using RO water from a wash bottle. Wipe the electrode with a Kimwipe.
  2. Immerse the electrode in pH 10 buffer.
  3. Press OK
  4. It will show

SLOPECal -58.93 mv/pH (or a similar number)

ZEROCal 6.953 (or a similar number)

Record these two numbers in your lab notebook.

  1. Press OK
  2. When the yellow rectangle in the top right corner changes to blue, press OK.
  3. The electrode is now calibrated.
  4. Remove the electrode from the pH 10 buffer. Hold a waste beaker under the electrode and rinse the electrode using RO water from a wash bottle. Wipe the electrode with a Kimwipe.
  5. Attach the titration cup to the titration head by tightening the blue ring to hold it in place.
  6. Put the electrode into a hole in the top of the titration head, so that the electrode is in the titration cup.

 

Rinse the Burette

 

To “rinse” the burette means to push the liquid titrant from the burette into the titration cup (before doing a titration). After the burette is rinsed, it automatically refills itself. You would normally rinse the burette 3 or 4 times when you put a new titrant into the bottle. In this procedure you will rinse the burette only once.

  1. Press the Home button (a picture of a house) on the G20.
  2. On the G20, press Manual
  3. Press Burette
  4. Press Rinse
  5. Press Start
  6. Loosen the blue ring holding the titration cup and discard its contents in the sink. Rinse out the titration cup with RO water before putting a sample in it. Hold a waste beaker under the electrode, stirrer, and dispensing tip and use an RO wash bottle to rinse them off and remove all traces of the previous liquid.

 

Titrate Unknown Samples

 

  1. Press the Home button (a picture of a house) on the G20.
  2. Press Methods

39. Press PHOS Phosphoric Acid

  1. Press Start
  2. In the ID 1 field, type in your name followed by 01 for your first titration, 02 for your second titration, etc. For example, if your name is Ravi Shankar and you are now running your second titration on the G20, you would enter

Ravi Shankar 02 in the ID 1 field

  1. Press OK
  2. Press Start
  3. The G20 will tell you to Add Sample 1/1. Do not press OK yet. You must press OK only after you have added the sample.
  4. Pipette 10.00 ml of your diluted unknown into the titration cup. Add enough RO water to bring the level up to about 40 mL in the titration cup. (There are volume markings on the titration cup.)
  5. Attach the titration cup to the titration head by tightening the blue ring to hold it in place. 47. Press OK
  6. The titration will proceed automatically and it will stop once 30 mL of titrant have been added. The data from the titration (volume and pH readings) will be sent automatically to the laptop computer.
  7. After the titration has ended, loosen the blue ring holding the titration cup and discard its contents in the sink. Rinse out the titration cup with RO water before putting the next aliquot of unknown into it.
  8. Hold a large waste beaker under the exposed electrode, stirrer, and dispensing tip and use an RO wash bottle to rinse them off and remove all traces of the titration mixture.
  9. Repeat steps 40 to 50 to do a total of three automated titrations. Remember to give each one a unique name, which should be

Your Name 01

Your Name 02

Your Name 03 in the ID 1 field

  1. Remove the electrode from the titration head and put it into the test tube holding water.

 

E-mail Data to Yourself and your Instructor

 

After you have done the three automated titrations on the G20, the data for each titration will be in the LabX software on the laptop computer. From the LabX software, you will export the data for each titration to an Excel file. You will then e-mail the three files to yourself and also to your lab instructor.

  1. In the LabX software on the laptop, in Data, find and double-click on one of your samples.
  2. Press the Measured Values tab.
  3. Press the Print Data tab.
  4. Press the little down arrow next to Export to.
  5. Select XLSX File
  6. Press OK
  7. For Filename, type in the name that you gave to the sample, e.g., Ravi Shankar 02.
  8. Select Downloads. Select AT method. This is where your file will be saved.
  9. Press Save.
  10. Repeat steps 53 to 61 for your two other titrations.
  11. Click on the WiFi icon in the taskbar at the bottom right corner of the screen.
  12. Click Sheridan Secure Access.
  13. Turn off Connect automatically.
  14. Press Connect.
  15. Log in to the Sheridan network with your Sheridan username and password.
  16. Open Internet Explorer.
  17. Enter the following URL: https:/mail.sheridancollege.ca/owa/auth.owa
  18. Sign in to Outlook Web App.
  19. Press new mail
  20. Click on Insert. Click on Attachment.
  21. Attach the three files that contain your data.
  22. Send the e-mail to yourself and also to your instructor.
  23. Sign out from Outlook.
  24. Click on the WiFi icon in the taskbar.
  25. Click on Sheridan Secure Access.
  26. Press Disconnect.

 

 

 

 

 

 

 

 

Report Format

 

Note: For this lab report, the two partners will together submit one joint lab report. That is, the two partners will work together to write only one lab report. Both partners will get the same grade, based on the one lab report. This is different from other experiments, in which the two partners write separate, independent lab reports.

 

Reference: Textbook, pages 624 – 625 and 388

This report will include both the manual and automated titrations done in the two weeks of this experiment.

The lab report will include:

COVER SHEET

TITLE

PURPOSE

OBSERVATIONS

CALCULATIONS

RESULTS

QUESTION

CONCLUSION

SOURCES OF ERROR

REFERENCES

PHOTOCOPY OF ALL DATA FROM LAB NOTEBOOK

 

The cover sheet, title, purpose and references are the same as in the previous reports.

 

OBSERVATIONS

Manual Titration:

State which unknown you received (A, B, C, or D).

State the standardized concentration of the titrant used in the manual titration to four significant figures.

The numerical observations for the manual titration must be entered into a spreadsheet (shown below) which will be used to plot the manual titration curve and the first derivative of the manual titration curve.

 

You will be plotting:

  • Titration curve
  • First derivative of titration curve

Make up a spreadsheet in Excel using your volume and pH readings from the manual titration as the first and third columns (these are your observations). Six columns are required.

 

The column headings for the spreadsheet are:

  • V (volume of titrant (mL)) observation
  • ∆V
  • pH observation
  • ∆ pH
  • ∆ pH / ∆V
  • Average volume

 

Set up the spreadsheet to calculate and fill in the remaining 4 columns.

For each row in the spreadsheet, starting in the second row, these are the meanings of the variables:

  • ∆V – the difference in volume between this row and the row above, i.e., the volume in this row minus the volume in the row above
  • ∆ pH – the difference in pH between this row and the row above, i.e., the pH in this row minus the pH in the row above
  • ∆ pH / ∆V – the ∆ pH in this row divided by the ∆V in this row
  • Average volume – the average of the volume in this row and the volume in the row above

An example of a spreadsheet with typical data might look like this near the beginning.

V

Titrant vol

(mL)

∆V mL

pH

 

∆ pH

 

∆ pH

∆V

Average volume, mL

0.00

2.06

1.02

1.02

2.13

 

0.07

0.069

(1.02+0.00)/2

= 0.51

2.03

1.01

2.19

0.06

0.059

(2.03+1.02)/2

= 1.52

 

Fill in the Excel spreadsheet shown above using your volume and pH data in the first and third columns. Have Excel calculate the changes, the ratio of changes, and the average volume needed in the other columns. Include the spreadsheet in the Observations section of your lab report.

From the above spreadsheet, plot the following 2 graphs using volume as the horizontal axis. Include the plots in the Observations section of your report. Each plot must show the individual data points, not just a smooth line.

Graph 1 will plot pH vs titrant volume (titration curve).

Graph 2 will plot ∆pH/∆V (column 5) vs average volume (first derivative curve).

Notice that in Graph 1, the actual volume readings (from the first column of your spreadsheet) are used for the horizontal axis, but for Graph 2, the average volume (from the sixth column of your spreadsheet) is used for the horizontal axis.

Each of these graphs should fill a page. If a graph is too small, the end point volumes cannot be accurately read.

Your 2 plotted curves should resemble those in your textbook on page 625, Figure 21 – 21, with the difference being that your plotted curves will have two end points instead of one. You will find the two end points from the titration curve and you will also find them from the first derivative curve.

 

End points:

TITRATION CURVE

The titration curve should have two steeply – rising parts. The end points are halfway up each steeply – rising part. From the graph, estimate the volumes at the two end points. Label these two volumes on your titration curve.

 

FIRST DERIVATIVE CURVE

The end points occur where the first derivative curve has its maxima (two highest points).

The volume at the second maximum should be about twice the volume at the first maximum. Indicate these volumes on your plot.

 

Automated Titration:

State the standardized concentration of the titrant used in the automated titration to four significant figures.

List the names of the three files containing the automated titration data that you emailed to yourself and to your instructor.

Copy and paste the three Excel files into the Observations section of the report. The Excel files indicate where the autotitrator found the two end points. They are listed as EQP1 and EQP2.

List all of the endpoints found in both types of titration in a single table like the one shown below, where V1 is the volume at the first end point and V2 is the volume at the second end point.

 

V1, mL

V2, mL

V2/V1

Manual titration curve

 

 

 

Manual first derivative curve

 

 

 

Automated titration 1

 

 

 

Automated titration 2

 

 

 

Automated titration 3

 

 

 

 

CALCULATIONS

Molarity of unknown:

Calculate the molarity of your phosphoric acid unknown using the end point volumes V1 determined by the manual titration curve, manual first derivative curve, and each of the automated titrations.

At the first end point, the phosphoric acid molarity can be calculated from

𝑀𝑀𝐵𝐵𝑉𝑉𝐵𝐵

𝑀𝑀𝐴𝐴 =

𝑉𝑉𝐴𝐴

where MB is the molarity of the NaOH titrant, VB is the volume of NaOH titrant needed to reach the end point, and VA is the volume of diluted unknown acid that was titrated.

Remember that VA was 20.00 mL for the manual titration, and it was 10.00 mL for the automated titrations. It is also possible that the molarity of the base that you used in the manual titration was different from the molarity of the base used in the automated titrations. The value of MA that you calculate in this equation is the molarity of your unknown after you diluted it.

Recall that you diluted your unknown sample by pipetting 25.00 mL of your original unknown into a volumetric flask and making it up to 250.0 mL. Therefore the dilution factor is

250.0 𝑚𝑚𝑚𝑚

𝑑𝑑𝑑𝑑𝑑𝑑𝑑𝑑𝑑𝑑𝑑𝑑𝑑𝑑𝑛𝑛 𝑓𝑓𝑓𝑓𝑓𝑓𝑑𝑑𝑑𝑑𝑓𝑓 = 𝑑𝑑𝑓𝑓 = = 10.00

25.00 𝑚𝑚𝑚𝑚

In the Results section, you must report the concentration that your unknown had as you received it, not after you diluted it. The molarity that you must report is

𝑀𝑀𝑢𝑢𝑢𝑢𝑢𝑢𝑢𝑢𝑜𝑜𝑤𝑤𝑢𝑢 = 𝑀𝑀𝐴𝐴 × 𝑑𝑑𝑓𝑓

For the three automated titrations, calculate the mean, standard deviation, and percent relative standard deviation for the molarity of your unknown acid.

 

Determination of Dissociation Constant:

Using only the data from the first automated titration, find the values of pKa1 and pKa2 for phosphoric acid as follows. pKa1 is equal to the pH of the titration curve at a volume which is exactly half way from the start to the volume at the first end point. pKa2 is equal to the pH of the titration curve at a volume exactly half way between the first end point and the second end point.

Find the pH reading corresponding to the volume reading which is closest to half way between zero volume and the volume of the first end point. Take this pH reading to be your value for pKa1.

Find the pH reading corresponding to the volume reading which is closest to half way between the volume of the first end point and the volume of the second end point. Take this pH reading to be your value for pKa2.

 

RESULTS

In a single table, list the following molarities for your unknown phosphoric acid:

  • Molarity from manual titration curve
  • Molarity from manual first derivative curve
  • Molarity from each of the three automated titrations
  • Average molarity for the three automated titrations
  • Standard deviation of the molarities for the three automated titrations
  • Percent relative standard deviation of the molarities for the three automated titrations

In a second table, list the two pKa values that you found for phosphoric acid. Also, in this table, list the known values for pKa1 and pKa2 for phosphoric acid. Find the known values in your textbook or another reliable source. Cite your source in the References.

 

QUESTION

Phosphoric acid is triprotic. We should see three “jumps” in the titration curve, one for each ionization. However, in your automated titrations, you observed only two jumps in the curve, not three. Explain why you did not see a third jump in the titration curve.

 

CONCLUSION

The Conclusion is two sentences which state “From the manual titration curve, the molarity of the unknown phosphoric acid was found to be ____ M. The mean molarity found in the automated titrations was ____ M.”

SOURCES OF ERROR

List at least 3 sources of error are to be discussed here with their effect on the molarity of the unknown acid. Remember to state the type of error.

 

REFERENCES

List all of your references.

 

 

Laboratory 3

Gravimetric Determination of Chloride in a

Soluble Sample

 

Objective

This experiment is performed to determine the percent chloride in an unknown soluble sample using gravimetric methods.

Background

The work must be shared in order to complete the lab in the allotted time.

References: CDROM: 37B1

Analytical Text: Chapter 12

Chemicals Required

Week 1: Week 2:

Concentrated nitric acid 6 M nitric acid

6 M ammonium hydroxide RO water

0.1 M silver nitrate Concentrated hydrochloric acid

6 M nitric acid

RO water

Unknown sample

Equipment Required

Week 1: Week 2:

3-400 mL beakers Rubber policeman

3-Stirring rods Wash bottle

3-Watch glasses

3-Glass filtering crucibles Vacuum filtration set up

 

Procedure

Week 1

Cleaning Filtering Crucibles

1 Set up and properly secure a Buchner flask. Place about 50 mL of RO water in the flask. Turn on the fume vent over the filtration setup.

  1. Place a crucible in a funnel and insert into the Buchner flask.
  2. Add about 5 mL of concentrated nitric acid to the glass filtering crucible and let stand for 5 minutes. Open the vacuum tap and vacuum filter the nitric acid into the Buchner funnel.
  3. Fill the filtering crucible with water and filter using vacuum. Repeat 2 more times and break the suction by turning off the vacuum.
  4. Add about 5 mL of 6 M NH4OH to the crucible, let stand a few minutes and vacuum filter.
  5. Fill the crucible with RO water and filter using vacuum. Repeat 6 – 8 times.
  6. Repeat steps 2 – 6 for the other 2 crucibles.
  7. Mark the crucibles with numbers 1, 2, and 3, place into a large beaker and dry in the oven at 105°C until the end of the lab period.
  8. Just before the end of the lab period, remove the crucibles from the oven, and place them into the desiccator until your next class.

NOTE: Do not touch the crucibles with your fingers after this. Handle only with tongs or kimwipes.

 

Sample Precipitation

  1. Your instructor or technologist will provide you with your unknown sample (which has been dried by the lab technologist). Immediately record in your lab notebook the number of your unknown sample.
  2. Add 7 – 8 mL of 6 M nitric acid to about 300 mL of RO water.
  3. Measure three samples of 0.15 – 0.20 g of unknown into 400 mL beakers by weighing on the analytical balance. Weigh the samples to the nearest 0.0001 g.
  4. Use 100 mL of the diluted nitric acid (from step 2) to dissolve each sample.
  5. Slowly, while stirring, add 45 mL of 0.1 M AgNO3 (using a graduated cylinder) to each solution.
  6. Heat almost to boiling and digest (without boiling) the solutions for about 10 minutes.
  7. Check each solution by adding a few drops of AgNO3. If a white cloud forms, add another 5 mL of AgNO3 and digest another 10 minutes. Repeat until no white cloud forms when AgNO3 is added.
  8. Let the beakers cool. Cover the beakers using Parafilm and store in your locker until the next lab period.

 

Week 2

Filtration and Drying

  1. Retrieve your crucibles from the desiccator and weigh them on an analytical balance. Weigh the crucibles to the nearest 0.0001 g.
  2. Set up a vacuum filtration apparatus.
  3. For each beaker containing AgCl precipitate, decant the supernatant liquid through the weighed filtering crucible.
  4. Make a dilute solution of nitric acid by adding about 2 mL of 6M HNO3 to 500 mL of RO water.
  5. Wash each precipitate (while it is still in the beaker) using 60 – 80 mL of the diluted nitric acid. Decant the precipitate washings through the filtering crucible. Do this 3 or 4 times for each precipitate.
  6. Quantitatively transfer the AgCl into the crucibles using your wash bottle. Use a rubber policeman to scrape any particles adhering to the walls of the beaker.
  7. Continue washing the precipitate in the crucible with diluted HNO3 until the filtrate is free of Ag+. To test the washings for Ag+, collect a small volume from the filtration flask in a small beaker and add a few drops of HCI. Dump out the rest of the liquid in the filtration flask. If no cloudy precipitate forms, no Ag+ is present.
  8. Repeat steps 3 to 7 for your other 2 samples
  9. Dry the precipitates at 105°C for at least 1 hour.
  10. Store the crucibles in a desiccator until cool.
  11. Weigh the crucibles and their contents to the nearest 0.0001 g.
  12. Calculate the % Cl in your sample.

 

Notes

  1. A slightly acidic solution is required for selective precipitation of AgCl.
  2. A slight excess of silver will decrease the solubility of AgCl.
  3. Photo decomposition of finely divided silver can occur and produce Ag and Cl2, leading to low results for your chloride determination. If you see a violet colour in your precipitate, elemental silver is present and an additional reaction can occur, producing AgCl and ClO3, causing high results. These effects can be reduced to a minimum by avoiding exposure of your precipitate to high levels of light. Efficient performance of your precipitation will also reduce the production of elemental silver. Your product should be stored in your locker (in the dark) until weighed.
  4. Use only the amount of silver nitrate that you need and return any unused silver nitrate.
  5. Use a separate stirring rod for each sample.

Lab Report Format

 

This lab report requires:

 

COVER SHEET

TITLE

PURPOSE

OBSERVATIONS

CALCULATIONS

RESULTS

CONCLUSION

POSSIBLE SOURCES OF ERROR

REFERENCES

PHOTOCOPY OF ALL DATA FROM LAB NOTEBOOK

 

COVER SHEET

The cover sheet is to be completely filled in, attached to the report and placed in the appropriate folder on the front bench. Sample number, mean, standard deviation and % relative standard deviation must appear on the cover sheet or a mark of 0 for accuracy and precision will be given.

 

TITLE

The title should be the same as that in your lab manual.

 

PURPOSE

The purpose should be 1 or 2 short sentences describing the purpose of your lab.

 

OBSERVATIONS

This section include both physical and numerical observations.

Physical observations may be written in sentence format or in a table. Each week observations should be a separate paragraph or in separate tables.

Numerical observations are to be in table format. The table must have an appropriate title and heading.

The observed data must include the following: weight of unknown sample, at least two weight readings for cleaned, empty crucible, and at least two weight readings for crucible with precipitate.

 

CALCULATIONS

As in previous labs, a sample calculation should be shown for the first sample. Mean, standard deviation and % relative standard deviation are to be calculated.

 

RESULTS

The results of your calculations are to be summarized into a table. The results table should show the three individual results for % chloride, the average % chloride, the standard deviation, and the % relative standard deviation.

 

CONCLUSION

The Conclusion is a sentence stating “Sample number ___ was found to contain ___ % chloride with a relative standard deviation of ____ %.

 

SOURCES OF ERROR

List at least 3 sources of error. Describe the type of error and whether it will increase or decrease the % Cl in your final answer.

 

REFERENCES

List all references used in your report. Remember your text, class notes and lab techniques notes are all applicable to your lab. All web sites must be easily accessed by the instructor.

Check to make sure they are accurate and can be easily accessed using your reference.

Marks are deducted for poor references.

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