Performance Tasks for SCH 4U14U+… · · 2017-12-28following Performance Tasks are ... (not...

Last Revised By D. Ridge and R. Tanner on 17/12/27

SCH 4U PERFORMANCE TASKS 2017-2018

A list of investigations from which you may make your choices follows.

Choices must be made no later than the start of Unit 5.

You are to work in a group of 2, 3, or 4 people– 4 people is the maximum for any group!.

Four days of class time only will be provided for actual completion of the laboratory

investigation(s) however each performance task will require preliminary research.

Answers, by group, to the Introduction questions, Prediction and Safety are due on

the beginning of class on the FIRST day of the performance task. These are to be

submitted to the Classroom.

Every student must individually keep organized and detailed Observations and complete

all Analysis and Discussion questions. Your observations, graphs, analysis, and

discussion will be handed in on the final summative day. These are to be submitted,

by group, on Google Classroom. It is a good practice to create the data tables and

input any data collected on a DAILY basis.

On the final day of the performance task, each person will be required to answer questions

about their specific performance task in both a written calculation test and an interview.

No study aids are permitted for the written test, but notes may be used for the

interview.

Due to the removal of chemicals for the renovation of the storage room, only the

following Performance Tasks are available to choose from:

Number INVESTIGATION

1 Acid Content of Fruit Juice

2 Oxidizing Power of Laundry Bleach

3 Buffers and Le Châtelier's Principle

4 Determination of The Order Of Reaction

5 Rates of Reaction and Rate Law

6 Electrochemical Cells, Thermodynamics and the Equilibrium Constant

7 Spectrophotometric Determination of an Equilibrium Constant

8 Ksp of Copper (II) Tartrate

9 Spectrophotometric Determination of Aspirin

10 Vitamin C - an Important Antioxidant

11 Determination of Salt in Potato Chips by the Mohr Method


1 - ACID CONTENT OF FRUIT JUICE

Introduction: Use a pre-made citric juice sample provided by the teacher. The manufacturer’s concentration is 0.100 M.

1. Obtain the Ka values for citric acid and examples of the pH curve vs NaOH.

2. How can the Ka value(s) be determined experimentally from your pH curve?

3. As the final endpoint is the 3rd equivalence point with a pH of about 9.5, research an appropriate

indicator, (not phenolphthalein), to use with this experiment to aid with the equivalence point

determination.

Safety: Describe any chemicals used or their products that have specific standards associated with them. Identify

any hazardous procedures. State all precautions you will take. Prediction: Questions 1-12 1. Calculate the initial pH of a 0.100 M citric acid solution.

2. Calculate the final equivalence point volume for the titration of 10.00 mL of a 0.100 M citric acid

solution using 0.200 M NaOH.

3. Calculate the pH at the final equivalence point.

4. As the mole ratio is all that changes in the stoichiometry calculations for the 1st and 2nd equivalence

volumes

a) From the final equivalence point volume calculated in (2) above, determine the volume of NaOH

needed to reach the 1st and the 2nd equivalence points.

b) Determine the volumes at the 3 ½ equivalence points.

5. Using the Ka values, calculate the pHs at the 3 half equivalence point volumes.

6. Using the information on the graph below, determine the pH for the 1st and 2nd equivalence points.

7. Summarize, in your report, a table of, the V½ eq. pt., Veq. pt and the pHs for the 3 acidic Hs. These

volumes and pH values will be very important to keep track of when you do your titration.


Calculations that will need to be able to do for your Analysis: 8. Calculate the acid concentration, [H3A], if the initial pH is 2.09.

9. Calculate the percentage error versus the 0.100 M theoretical.

10. Calculate the acid concentration, [H3A], if 17.05 mL of 0.200 M NaOH in required to reach the

equivalence point in the reaction of 10.00 mL of juice.

11. Calculate the percentage error versus the 0.100 M theoretical.

12. Calculate the percentage error between the 2 different [H3A] values, using the most accurate as the

“theoretical value”.

Materials/Equipment:

0.100 M citric acid fruit juice standardized NaOH solution 250 mL erlenmeyer flasks

pH meter pipet burette your chosen (approved) indicator

Procedure: 1. Pipette 10.00 mL of a fruit juice into a 100 mL beaker. Measure the pH of the juice and record the

value.

2. Use the initial pH to determine the actual concentration of your acid and redo predictions.

3. Add 2 – 3 drops of your chosen indicator (approved by your teacher)

4. Slowly titrate with the standardized sodium hydroxide in small increments, reading the buret to 2

decimal places. Try to find each equivalence point - refer to the table in your prediction. After each

increment of NaOH is added, wait for the pH to stabilize, record the total volume added and the pH at

that point. Continue to add the NaOH until the pH values have leveled off.

5. Repeat with the SAME type of juice until a very good graph is obtained. Perform until three good

trials have been obtained.

6. Plot pH (y-axis) as a function of the volume of NaOH added (x-axis). Plot all trials on one graph.

Determine the equivalence point from the graph and the equivalence point volume of base added.

7. Check the procedure for the second method with your teacher before beginning and complete the

second method to determine the concentration of the fruit juice. Analysis and Discussion: 1. Graph the pH data, placing all trials on one graph and label the graph fully with all relevant reactions

2. Similar to the questions in the prediction, calculate the concentration of the juice from the equivalence

points of the graph.

3. From the pH at the 1/2 equivalence point volume (for each equivalence point), determine the “Ka”s

for the juice.

4. Using the Ka1 and the initial pH, determine the acid concentration and compare with the value from

#2.

5. Determine the concentration of the acid tested from the 2nd method

6. Calculate % error(s) from any manufacturer’s values.

7. Calculate % error between the different methods.

8. Describe all design errors that were encountered during the performance task and include an

improvement that was made to account for each error.


2 - OXIDIZING POWER OF LAUNDRY BLEACH Introduction: Many times in chemistry we must use indirect methods to find the concentration of a chemical in a

solution. One such method is as an iodometric analysis. In this experiment the concentration of NaOCl,

the active reagent in laundry bleaches, will be found.

Purchase a bleach solution that can tell you, on the bottle or the company’s website, the % W/W of

NaOCl in the bleach. Try not to buy concentrated bleach, try to get thin liquid bleach. 1. The bleach will react with iodide in a REDOX reaction. Acetic acid is added to provide the hydrogen

ion needed for the reaction. Write the balanced redox reaction of OCl- with I-. Indicate the species

oxidized, species reduced, oxidizing agent and reducing agent.

2. The iodine produced undergoes another redox reaction with the thiosulfate ion. Write this balanced

redox reaction. Indicate the species oxidized, species reduced, oxidizing agent and reducing agent.

3. Combine the equations from step 1 and step 2 to give the overall reaction. In the first reaction, iodide

ion is added in excess. It reacts with the iodine produced and creates a coloured ion. Write the

equation for this reaction.

5. This coloured ion will disappear during the titration and the solution will turn colourless abruptly. To

help with the endpoint determination, spray starch is added. What will the starch react with and what

will the colour be?

6. Why can’t the starch be added until the solution is a light amber, red-orange colour?

7. Does it matter if the starch is added when the solution trials have different colours of light amber? Safety: Describe any chemicals used or their products that have specific hazards associated with them. Identify any

hazardous procedures. State all precautions you will take. Prediction: Questions 1 - 11

Use the formulas from the last page of the 4U data tables (mostly Grade 11) 1. Most bleaches have a density of 1.09 g/mL. Calculate the mass of bleach in a 5.00 mL sample.

2. Using your NaOCl % W/W and the mass determined in (1) calculate the mass of NaOCl in the 5.00

mL bleach sample.

3. From the mass, calculate the molar concentration, C, of the NaOCl in 5.00 mL of bleach.

4. Determine the concentration of the 5.00 mL NaOCl if diluted to a new volume of 100.00 mL.

5. Calculate the equivalence point volume of a 25.00 mL of the diluted NaOCl with 0.150M Na2S2O3

solution.

6. Calculate the mass of Na2S2O3 required to create 100.00 mL of a 0.150 M solution.

Calculations that will need to be able to do for your Analysis: 7. Calculate the concentration of the dilute bleach if 13.50 mL of 0.150M Na2S2O3 solution is required

to reach with 25.00 mL of dilute bleach.

8. Determine the concentration of the NaOCl in the original 5.00 mL bleach if the volume of the diluted

bleach is 100.00 mL (the flask) and the dilute bleach concentration is from (1) above.

9. Using the concentration above, determine the mass of NaOCl in 5.00 mL of original bleach.

10. If the mass of 5.00 mL of original bleach is 5.68 g, calculate the NaOCl % W/W in the original

bleach.

11. The manufacturer’s claim is 5.75 % W/W, calculate the percentage error.


Equipment/Materials: Buret 250 mL Erlenmeyer flask 100.00 mL volumetric flasks

25.00 mL transfer pipet 5.00 mL pipet solid sodium thiosulfate

glacial acetic acid solid potassium iodide spray starch Procedure: Sodium thiosulfate solution 1. Prepare a 0.150 M sodium thiosulfate from the solid using a 100.00 mL volumetric flask.

- Using a funnel, add the measured solid to the 100.00 mL flask, washing the funnel and weigh boat

and adding it to the flask. Add approximately 50 mL to the flask and mix until the solid dissolves.

Fill the flask to the mark with distilled water. Label the flask with masking tape and your group’s

name for use throughout the task. Dilution of Bleach 1. Place a clean 100.00 mL Volumetric flask on the balance and zero.

2. Use a 5.00 mL pipet to measure out a sample of your laundry bleach and add to the flask. Record the

mass of the 5.00 mL of bleach. Calculate the density of the bleach in g/mL. Record.

3. Rinse the 5.00 mL pipet with distilled water and add this to the volumetric flask. Slowly add distilled

water to the mark, swirling regularly (with cap on) to mix the solution.

4. Redo the prediction, #1 - #5 before you do your titration. Titration 5. Rinse the buret with distilled water, check for leaks or clogs. Then rinse with small portions of

sodium thiosulfate solution. Drain these samples through the buret and discard them. Fill the buret

with the sodium thiosulfate solution. Adjust the level and make sure the tip is filled. Record the

initial level (to 2 decimal places, 0.00 or 0.05 L) of the solution in the buret on the data table.

6. Use a transfer pipet to measure out a 25.00 mL portion of the bleach solution. Place the sample in an

Erlenmeyer flask.

7. Add approximately 2.0 grams of potassium iodide to the flask.

8. In the fumehood, add 10. mL of glacial acetic acid (teacher does this).

9. Titrate with the sodium thiosulfate until the solution is very light amber- almost yellow colour. At

this time add several sprays of spray starch as the indicator.

10. Titrate slowly until the solution is colourless. Record the amount of sodium thiosulfate solution (to 2

decimal places) in the buret at the end of the titration. Note: The blue colour may reappear after the

titration has been completed due to air oxidation of the iodide.

11. Repeat titration until you have 3 good trials minimum. To determine a “good” trial, calculate the

average of the trials and ensure that each is within +/-5% of that average.

12. Find the amount of NaOCl in a bleach sample using a SECOND different technique. Analysis and Discussion:

1. Using your data from Method 1, perform the analysis using steps #7 to #11 in the Prediction.

2. The additional methods will require different analysis calculations to determine the mass of NaOCl in

the bleach or % W/W.

3. Calculate % error between the different methods.


improvement that was made to account for EACH error.


3 – BUFFERS AND LE CHÂTELIER'S PRINCIPLE

Introduction: 1. What is a buffer?

2. Why are buffer systems extremely important to human health?

3. What is the principal buffer system in blood serum based upon? The acid from this equilibrium is

unstable and is also in equilibrium with what gas?

4. Give the balanced equations for the blood buffer equilibrium systems.

5. The kidneys help to regulate the pH of blood in several ways, such as increasing or decreasing

excretion of hydronium ions, H3O+, in urine. Explain how the kidneys might respond to the following

conditions.

(a) The blood pH rises to 7.48.

(b) The blood pH sinks to 7.33. Safety: Describe any chemicals used or their products that have specific hazards associated with them. Identify

any hazardous procedures. State all precautions you will take. Prediction: 1. Calculate the pH of a 1:1 acetic acid and sodium acetate buffer.

2. If it takes 40.00 mL of acetic acid to get 25.00 mL of 0.10 M sodium acetate to a 1:1 buffer, what is

the concentration of the acetic acid.

3. After adding 5:00 mL of 0.10 M HCl (aq) to 25.00 mL of your 1:1 buffer assuming 0.10 M acetic acid

and 0.10 M sodium acetate, what is the predicted pH?

4. After adding 5.00 mL of 0.10 M NaOH (aq) to 25.00 mL of your 1:1 buffer, what is the predicted pH?

5. Using research, find the formula for the Buffer Capacity and calculate the theoretical buffer capacity

based on the answers to questions 3 and 4. Procedure: Create a 1:1 mole buffer that using 0.10 M sodium acetate solution and an UNKNOWN concentration of

acetic acid (commercial vinegar). 1. Prepare for the titration by rinsing the buret with distilled water to check for any clogs or leaks.

Once a buret is good, rinse with the unknown acetic acid and drain through the tip. Fill, using a

funnel, to about the 0 or 1 mL mark. Record all buret volumes to 2 decimal places (0.00 or 0.05 L).

2. Determine the concentration of the acetic acid by titrating 25.00 mL of 0.100 M sodium acetate with

the UNKNOWN acetic acid until you have reached the desired pH of the 1:1 mole buffer. Place a pH

meter in the sodium acetate sample and place the unknown acetic acid in the buret.

3. Titrate a 25.00 mL sample of the buffer with 0.10 M HCl (aq) by adding 10.00 mL of HCl (aq) (drop

by drop) and determine whether this is an effective buffer, using the buffer capacity formula for

calculations.

4. Plot a titration curve and label the buffer region

5. Repeat this procedure with a new 25.00 mL sample of buffer and 0.10 M NaOH (aq) and determine

whether this is an effective buffer.

6. A possible second method is to create a DIFFERENT 1:1 mole buffer and test it in a similar manner.

Determine if this buffer is more or less effective than the acetic acid/sodium acetate buffer. Analysis and Discussion: 1. How much 0.10 M HCl or NaOH was needed to make each buffer reach a pH of 7.00? Does that

agree with theoretical (calculated) amount of 0.10 M HCl?

2. Explain why the buffer becomes less effective as more acid (or more base) is added.


3. Compare and contrast the system you used in your investigation with the carbonic acid/carbonate

buffer system in the blood. How are the systems similar? How are they different?




4 – DETERMINATION OF THE ORDER OF REACTION

Introduction: 1. Write the balanced net ionic equation for the reaction between magnesium and hydrochloric acid.How

would you expect the concentration of the acid and the amount of magnesium metal to affect the rate

of reaction?

3. Why would it be important that the length of magnesium used throughout the experiment be kept

constant?

4. Write the general rate equation for this reaction.Since rate can be expressed as the reciprocal of time

taken for the reaction, rewrite this rate equation in terms of time.Rewrite this rate equation after taking

the natural log of both sides of the rate equation.

7. If this equation were graphed, what type of graph would be obtained?

8. How can the order of reaction and the value of the rate constant be obtained from this graph?

9. Design an experimental procedure involving at least six 25.00 mL aliquots of hydrochloric acid or

varying concentration prepared from stock 6.0 M HCl.

Safety: Describe any chemicals used or their products that have specific hazards associated with them. Identify

any hazardous procedures. State all precautions you will take.

Prediction: If the reaction is first order for all reactants, predict the rate law. Research and find the actual rate law for

this reaction.

Procedure: 1. Perform the experiment from the introduction (FIRST method).

2. Determine and test a SECOND different method to determine the rate law.

Analysis and Discussion: 1. Complete your data table and construct your graph as discussed in the introduction.

2. Based on the data, give the order of this reaction with respect to hydrogen ions, the rate law, and the

value of the rate constant (for all methods used).

3. Compare this rate law to the theoretical rate law for this reaction, calculate the % error.




5 – RATES OF REACTION and RATE LAW

Introduction: 1. (a) Permanganate is a strong oxidizing agent. It will oxidize oxalic acid to carbon dioxide. Give the

balanced equation for this reaction.

(b) Rewrite this equation as a net ionic equation.How might the colours of permanganate and

manganese (II) ion be used to determine the rate of reaction?

3. List three factors which might control the rate of this reaction.

4. What is the accepted Rate law for this reaction?

5. When examining concentration effects, why would it be to your advantage to keep the total volume of

water plus sulphuric acid constant during the experiment?

6. How might you examine the effect of temperature on this reaction?

7. This reaction is auto catalyzed; i.e. one of the products (Mn2+) catalyzes the reaction. How might you

examine this effect?

8. Devise and outline a detailed procedure that will give you enough data to plot a graph. What solution

concentrations will you use? Have this procedure approved by your teacher.

Safety: Describe any chemicals used or their products that have specific hazards associated with them. Identify

any hazardous procedures. State all precautions you will take.

Prediction: If the reaction is first order for all reactants, predict the rate law. Research and find the actual rate law for

this reaction. How will an increase in temperature affect the rate? Explain fully.

Procedure: 1. Perform your experiment as you designed it, approved by your teacher, keeping a meticulous,

detailed account of the experiment.

2. Determine and test a SECOND different method to determine the rate of reaction.

Analysis and Discussion: 1) Graph all data and discuss all possible trends.

2) Give the experimental rate law for this reaction (for all methods used).

3) Compare to the theoretical value, calculate the % error.

4) How does increasing concentration of each reactant affect the rate of reaction?

5) How does increasing the temperature of the reactants affect the rate of reaction?

6) Describe all design errors that were encountered during the performance task and include an



6 – ELECTROCHEMICAL CELLS, THERMODYNAMICS AND

THE EQUILIBRIUM CONSTANT Introduction: 1. Using two of copper, zinc or magnesium and the appropriate nitrate or chloride solution;

(a) Illustrate the electrochemical cell that can be constructed with the placement of a voltmeter

included in the illustration. Identify the anode, cathode and indicate the direction of electron flow.

(b) Calculate E°.

2. Give the Nernst Equation and define all terms.

3. How can G be determined from E?

4. In a plot ofG vs T, what thermodynamic quantity is the slope equivalent to?

5. If G, S and T are known, how can H be determined?

6. How can the value of the equilibrium constant, Ke be determined? Safety: Describe any chemicals used or their products that have specific hazards associated with them. Identify

any hazardous procedures. State all precautions you will take. Prediction: Calculate the theoretical values for ∆G, ∆H, and ∆S at SATP for copper, zinc, and magnesium in their

associated nitrate or chloride solutions for each electrochemical cell. Procedure: Cell Construction

1. Place the copper (II) nitrate solution in a ceramic cup with the copper electrode in the solution.

2. Place the ceramic cup in a large beaker or cup

3. Half fill the beaker with zinc nitrate solution and place a zinc electrode in the beaker.

4. Connect with wires and voltmeter as per the introduction.

Data Collection

5. Record the voltage and the temperature of the cell. If the potential is negative, reverse the

connections.

6. Begin heating the water in the 600-mL beaker. Be certain that the test tubes are firmly clamped in

place.

7. Be careful not to move any part of the cell because the voltage will fluctuate if you do so.

8. Heat the cell to approximately 70° C. Record the new temperature and the cell potential.

9. Record the temperature and voltage at 15° C intervals as the cell cools.

10. When the temperature reaches room temperature, replace the hot water bath with an ice-water bath.

Try not to move the cell. After the cell has been in the ice-water bath for about 10 minutes, record the

temperature and the cell potential.

11. Test this method using different temperatures and materials Analysis and Discussion: 1. Calculate ∆G for the cell at each of these temperatures and plot ∆G versus temperature.

2. Calculate ∆S from the values of ∆G and ∆S, calculate ∆H at 298 K.

3. Calculate the value of Keq at the varying experimental temperatures.

4. Do these values match with theoretical values? Compare and discuss.




7 – SPECTROPHOTOMETRIC DETERMINATION

OF AN EQUILIBRIUM CONSTANT Introduction: 1. What is spectrophotometry?

2. The Beer-Lambert Law relates the percentage of light absorbed by a sample to its concentration.

State the Beer-Lambert Law. State the significance of all variables and constants in this law.

3. On a graph of Absorbance versus concentration, what is the slope equal to?

4. For which type of ion; weakly coloured or intensely coloured, would measurement of light absorbance

be most suitable as a means of determining low ion concentrations? Explain.

5. Are all wavelengths of incident light equally effective for measurements of absorbance by a coloured

ion complex? What wavelength would be appropriate for this experiment?

6. Why is nitric acid added to the solution?

7. We want any absorbency measured using the spectrophotometer to be a result of the coloured

complex. How might we correct for absorbance by the cuvette and solvent molecules and other ions

present?

8. Write the balanced chemical equation for the reaction between aqueous iron (III) nitrate, Fe(NO3)3

and potassium thiocyanate, KSCN. They react to produce the blood-red complex [Fe(SCN)]2+.

9. Give the equilibrium constant expression for the above reaction and value.

10. If the concentration of the iron solution is much greater than that of the KSCN solution upon mixing,

will the reaction establish equilibrium or go to completion?

11. How can the concentration of the product be determined from the volume and concentration of the

KSCN used in each trial?

12. Describe how the formation of this coloured complex can be used to determine the amount of iron

thiocyanate ion in solution.

Safety: List any chemicals used or their products that have specific hazards associated with them. Give the

hazards and the precautions you will take. Prediction: these calculations will be needed for your analysis. 1. Rewrite Beer-Lambert Law to include the equilibrium iron (III) thiocyanate ion concentration. Then

rearrange in terms of the ion.

2. To calculate the equilibrium iron (III) ion concentration, consider that the initial iron (III) ion

concentration is equal to equilibrium iron (III) ion plus the equilibrium iron (III) thiocyanate ion

concentration. Write this equation, then rearrange it in terms of equilibrium ion (III) ion concentration

3. The equilibrium thiocyanate ion concentration can be determined in the same manner. Write the

expression for this calculation.

4. What is the initial concentration of iron (III) ion and the thiocyanate ion if 60.00 mL of 0.0020 M

Fe(NO3)3 (aq) is mixed with 60.00 mL of 0.0020 M KSCN (aq), (volumes add)?

5. If the absorbance of the solution in #4 is found to be 0.255 and the slope of the standard curve is

38496 was is the equilibrium concentration of Fe(SCN)2+ and the Keq?

Equipment/Materials: Spectrovis or similar spectrophotometer cuvets

0.00200 M KSCN 0.200 M Fe(NO3)3 0.00200 M Fe(NO3)3

0.05 M HNO3 burets or pipets 50 mL beakers


Procedure: Part 1: Preparation of Standard Solutions

The chart below provides the volumes of reactants needed to prepare the standard solutions.

Solution 0.00200 M KSCN 0.200 M Fe(NO3)3 0.05 M HNO3

1 5.0 mL 5.0 mL 15.0 mL

2 4.0 mL 5.0 mL 16.0 mL

3 3.0 mL 5.0 mL 17.0 mL

4 2.0 mL 5.0 mL 18.0 mL

5 1.0 mL 5.0 mL 19.0 mL

Part II: Preparation of Equilibrium Mixtures

Use a graduated pipet to measure the volumes of the reactants listed below. Note that this set of

combinations uses the more dilute Fe(NO3)3 solution.

Solution 0.00200 M KSCN 0.00200 M Fe(NO3)3 0.05 M HNO3

1 1.0 mL 5.0 mL 4.0 mL

2 2.0 mL 5.0 mL 3.0 mL

3 3.0 mL 5.0 mL 2.0 mL

4 4.0 mL 5.0 mL 1.0 mL

5 5.0 mL 5.0 mL 0

Part III: Testing the Solutions.

1. Setup the Spectrovis and follow the instructions to collect the data.

2. Calibrate the sensor using a blank made with the 0.05 M HNO3 as the blank

3. Obtain absorbance readings for each of the other standard solutions

4. Obtain the absorbance readings of each of the equilibrium solutions.

Part IV: Alternate Procedure:

Determine and test a SECOND different method to find the equilibrium constant.

Analysis and Discussion: 1. Prepare your calibration curve of Absorbance vs [Fe(SCN)2+] ion using the standard solutions.

2. Determine the concentration of [Fe(SCN)2+] for each of the equilibrium trials.

3. From the concentration of [Fe(SCN)2+] produced and the original concentrations of the reactants,

construct tables to determine the equilibrium concentrations of all species.

4. Use these values to calculate the equilibrium constant for each trial.

5. Report the average value for the constant (for all methods used).

6. Compare this equilibrium constant to the theoretical value.

7. Calculate the % error vs theoretical values and calculate the % error between the different methods.




8 – KSP OF COPPER (II) TARTRATE Introduction: 1. What is spectrophotometry?

2. The Beer-Lambert Law relates the percentage of light absorbed by a sample to its concentration.

State the Beer-Lambert Law. State the significance of all variables and constants in this law.



be most suitable as a means of determining low ion concentrations? Explain.

5. What colour are copper (II) ions in solution.





present?

8. What is a solubility product expression?

9. If we wish to determine the Ksp of copper (II) tartrate, what kind of solution must we work with?

10. These solutions can be prepared by adding solutions of sodium tartrate to those of copper (II) nitrate.

The copper (II) tartrate will then form a precipitate. The solution that remains is then saturated with

respect to copper (II) tartrate. How must the amount of tartrate ion added compare to that of copper

(II) ions present? Why?

11. Give the dissociation equation for copper (II) tartrate.

12. Give the Ksp expression and value.

13. How are the concentrations of the copper (II) and tartrate ions related to each other in a saturated

solution? Describe how spectrophotometry can be used to determine the copper (II) ion concentration

in the saturated solution. Safety: Describe any chemicals used or their products that have specific hazards associated with them. Identify

any hazardous procedures. State all precautions you will take. Prediction: 1. Using the theoretical Ksp value, calculate the equilibrium concentrations of the copper (II) ion and the

tartrate ion when the solution is saturated.

2. If the absorbance of the saturated solution is found to be 0.092 and the slope of the standard curve is

4.4282 what is the equilibrium concentration of Cu2+ and the Ksp?

3. Calculate the mass of Cu(NO3)2 required to create 100.00 mL of a 0.100 M solution.

4. Calculate the mass of KNaC4H4O6·4H2O required to create 100.00 mL of a 0.100 M solution. Equipment/Materials: 25.00 mL Volumetric flasks spectrophotometer and cuvets

centrifuge and centrifuge tubes graduated pipets (0-2 mL)

copper (II) nitrate sodium potassium tartrate Procedure: Part I: Preparation of a saturated solution of copper (II) tartrate 1. Prepare 100.00 mL of a 0.100 M copper (II) nitrate solution from the solid.

Using a funnel, add the measured solid to the 100.00 mL flask, washing the funnel and weigh boat and

adding it to the flask. Add approximately 50 mL of distilled water to the flask and mix until the solid


dissolves. Fill the flask to the mark with distilled water. Label the flask with masking tape and your

group’s name for use throughout the task.

2. Prepare 100.00 mL of a 0.100 M potassium tartrate solution from the solid, using the instructions

given in step #1.

3. Using graduated pipets, place 10.00 mL of 0.100 M copper (II) nitrate and 10.00 mL of 0.100 M

sodium potassium tartrate in a 25.00 mL volumetric flask. Add distilled water to make 25.00 mL of

solution. Mix well.

4. Allow the solution to remain undisturbed for about 15 minutes while other solutions are being

prepared. The solution should form a precipitate.

5. Centrifuge to remove the precipitate. Save the clear blue solution. If this solution shows any

cloudiness or further precipitates, filter it until clear.

Part II: Preparation of standard copper (II) tartrate solutions 1. Prepare 25.00 mL (in a volumetric flask) of 0.0200 M copper (II) tartrate by using the following:

Measure (using a graduated pipet), 2.00 mL of 0.100 M copper (II) sulfate. Add (using a graduated

pipet), 5.00 mL of 0.100 M sodium tartrate, and dilute until the total volume is 25.00 mL. See chart

below.

2. Prepare 25.00 mL of 0.0180 M, 0.015, 0.012, 0.010 M copper (II) tartrate in a similar manner (using a

volumetric flask and graduated pipets). See chart below.

Part III: Determination of copper (II) ion concentration in the saturated copper (II) tartrate

solution 1. Setup the Spectrovis and follow the instructions to collect the data.

2. Calibrate the sensor using a blank made by diluting 5.0 mL of 0.100 M sodium potassium tartrate to

10 mL with distilled water.

3. Determine the absorbance of each of the five standard copper solutions.

4. Place the saturated copper (II) tartrate solution in a cuvette and record the absorbance of this solution.

Part IV: Alternative Procedure

Determine and test a SECOND different method to find the solubility constant for copper (II) tartrate. Analysis and Discussion: 1. Prepare your calibration curve and determine the concentration of the copper (II) ion in the saturated

copper (II) tartrate solution.

2. Calculate the value for the Ksp of copper (II) tartrate (for all methods used).

3. Determine the % error versus the theoretical value.

4. Calculate the % error between the different methods.



Solution 0.100 M cupper (II) nitrate 0.100 M sodium

potassium tartrate

0.0200M 5.00 mL 10.00 mL

0.0180 M 4.50 mL 10.00 mL

0.0150 M 3.75 mL 10.00 mL

0.0120 M 3.00 mL 10.00 mL

0.0100 M 2.50 mL 10.00 mL


9 – SPECTROPHOTOMETRIC ANALYSIS OF ASPIRIN

Introduction:

Purchase uncoated, uncoloured aspirin tablets.

1. What is spectrophotometry?

2. The Beer-Lambert Law relates the percentage of light absorbed by a sample to its concentration. State

the Beer-Lambert Law. State the significance of all variables and constants in this law.



be most suitable as a means of determining low ion concentrations? Explain fully.





present?

7. A coloured complex is formed between aspirin and the iron (III) ion when aspirin reacts first with

sodium hydroxide to form the salicylate dianion which is then reacted with acidified iron (III) ion to

produce the violet tetraaquosalicylatroiron (III) complex. Research and give the balanced equations

for these two reactions.

8. Describe how the formation of this coloured complex can be used to determine the amount of aspirin

in a commercial aspirin tablet. Safety: Describe any chemicals used or their products that have specific hazards associated with them. Identify

any hazardous procedures. State all precautions you will take. Prediction: 1. What is the concentration of ASA if 0.500 grams of ASA is added to 250.0 mL of water?

2. What is the concentration of ASA if 2.50 mL is diluted to 50.00 mL, solution “A”?

3. Calculate the ASA concentrations of the Standard solutions below B through to E. Calculations that will need to be able to do for your Analysis: 4. If the absorbance of the aspirin tablet is found to be 0.039 and the slope of the standard curve is

0.0129 what is the concentration of ASA in the tablet in the diluted solution?

5. What is the concentration of the ASA in the undiluted solution?

6. What is the mass of the ASA in the undiluted solution and consequently, the aspirin tablet? Equipment / Materials: 6 – 50.00 mL Volumetric flasks commercial aspirin product

10 mL graduated cylinder acetylsalicylic acid

500.00 mL Volumetric flask 1 M NaOH

250.00 mL Volumetric flask 0.0200 M iron (III) chloride buffer

5 mL pipet spectrophotometer

2 cuvettes graduated pipettes Procedure: Part I: Preparation of the Buffer:

1. Prepare 500.00 mL of a 0.0200 M iron (III) chloride buffer solution from the 1.62 g of the solid.

Using a funnel, add the measured solid to the 500.00 mL flask, washing the funnel and weigh boat and

adding it to the flask. Add approximately 350 mL of distilled water to the flask and mix until the solid

dissolves. Fill the flask to the mark with distilled water. Label the flask with masking tape and your

group’s name for use throughout the task.


Part II: Making Standards:

1. Mass 400 mg of acetylsalicylic acid in a 125 mL Erlenmeyer flask. Add 10.00 mL of a 1.0 M NaOH

solution to the flask, and heat until the contents begin to boil.

2. Quantitatively transfer the solution to a 250.0 mL volumetric flask. Rinse the Erlenmeyer flask and

add to the volumetric flask, then dilute with distilled water to the mark.

3. Pipet a 2.50 mL sample of this aspirin standard solution to a 50.00 mL volumetric flask. Dilute to the

mark with a 0.0200 M iron (III) chloride buffer solution. Label this solution "A".

4. Prepare similar solutions with 2.0, 1.5, 1.0, and 0.5 mL portions of the aspirin standard. Label these

"B, C, D, and E."

Part III: Making an unknown from a tablet:

1. Place one aspirin tablet in a 125 mL Erlenmeyer flask. Add 10.00 mL of a 1.0 M NaOH solution to

the flask, and heat until the contents begin to boil.

5. Quantitatively transfer the solution to a 250.0 mL volumetric flask. Rinse the Erlenmeyer flask and

add to the volumetric flask, then dilute with distilled water to the mark.

2. Pipet a 2.5 mL sample of this aspirin tablet solution to a 50.00 mL volumetric flask. Dilute to the

mark with a 0.0200 M iron (III) chloride buffer solution. Label this solution "unknown".

Part IV: Testing the Solutions:

1. Setup the Spectrovis and follow the instructions to collect the data.

2. Using a Kimwipe, wipe off the cuvet containing the blank this should be a cuvet of iron buffer, and

place this cuvet in the sample compartment, being sure to properly align it. (The line on the cuvet

should match up with the notch on the instrument.) Close the cover.

3. Obtain absorbance readings for each of the five standard solutions

4. Measure and record the absorbance of the unknown.

Part V: Alternative Method:

Determine and test a SECOND different method to find the amount of ASA in a tablet.

Analysis and Discussion:

1. Prepare your calibration curve and determine the concentration of the unknown.

2. Calculate the amount of ASA in one tablet of aspirin from the data (for all methods used).

3. Calculate the % error from manufacturer. 4. Calculate the % error between the different methods.




10 - VITAMIN C: AN IMPORTANT ANTIOXIDANT Introduction: Purchase uncoloured, unbuffered Vitamin C tablets.

1. What is the chemical name and formula for Vitamin C?

2. One way to determine the Vitamin C content of a sample is perform a REDOX reaction by titrating it

with iodine. What is the difficulty in preparing iodine solutions?

3. An alternative method is to prepare the iodine by the reaction of the iodate ion with the iodide ion in

acidic solution. Give the balanced equation for this redox reaction. Assign oxidation numbers and

identify the species undergoing oxidation and the one undergoing reduction.

4. Give the REDOX reaction between Vitamin C and iodine. Identify the species undergoing oxidation

and the one undergoing reduction.

5. Combine the two reactions, eliminating the iodine, to give the overall reaction.

6. The endpoint is determined once the Vitamin C is all gone, by the reaction of excess iodine with the

iodide ion produced above. The resulting triiodide ion reacting with the starch, (added as a spray).

What colour will this endpoint be?

7. In storage, the concentration of iodide ion in solution decreases fairly quickly over time. Why does

this happen?

8. Explain why the Vitamin C used for standardization must be fresh. Safety: List any chemicals used or their products that have specific hazards associated with them. Give the

hazards and the precautions you will take. Prediction: 1. What is the concentration of a 500 mg Vitamin C tablet, that is crushed, added to a 250.00 mL

volumetric flask and dissolved in distilled water?

2. Using the reaction between Vitamin C and iodate ion, (from the introduction #5), calculate the volume

of 0.00200M iodate ion solution will be needed to fully react with a 10.00 mL of the solution made in

#1 above.

3. What mass of potassium iodate is needed to make 100.00 mL of a 0.0020 M solution?

4. What mass of potassium iodide is needed to make 50.00 mL of a 0.600 M solution? Calculations that will need to be able to do for your Analysis: 5. Calculate the concentration of a 10.00 mL Vitamin C solution if 18.50 mL of 0.00200 M potassium

iodate needed to reach the endpoint.

6. What is the mass of Vitamin C in the 250.00 mL volumetric flask? Procedure: Preparation of the Vitamin C solution 1. Prepare the solution from the solid (prediction #1) using a 250.00 mL volumetric flask.

- Using a mortar and pestle, crush a tablet until it is a fine powder. Using a funnel, with a scoopula,

add as much powder as possible to a 250.00 mL Volumetric flask. Rinse the scoopula and the

pestle into the mortar and add to the volumetric flask. Rinse and add as much of the powder into

the flask as possible. Add approximately 100 mL to the flask and mix until the solid dissolves.

Fill the flask to the mark with distilled water. Label the flask with masking tape and your group’s

name for use throughout the task. Preparation of the 0.00200 M potassium iodate solution 1. Prepare the solution from the solid (prediction #3) using a 100.00 mL volumetric flask.


- Weigh out the solid into a weigh boat. Using a funnel, with a scoopula, add the powder as

possible to a 100.00 mL volumetric flask. Rinse the scoopula and the weigh boat into funnel.

Add approximately 50 mL to the flask and mix until the solid dissolves. Fill the flask to the mark

with distilled water. Label the flask with masking tape and your group’s name for use throughout

the task. Preparation of the 0.600 M potassium iodide solution 1. Prepare the solution from the solid (prediction #4) using a 50.00 mL volumetric flask.

- Weigh out the solid into a weigh boat. Using a funnel, with a scoopula, add the powder as

possible to a 50.00 mL volumetric flask. Rinse the scoopula and the weigh boat into funnel. Add

approximately 25 mL to the flask and mix until the solid dissolves. Fill the flask to the mark with

distilled water. Label the flask with masking tape and your group’s name for use throughout the

task. Titration 1. Rinse the buret with distilled water, check for leaks or clogs. Then rinse with small portions of

0.00200 M potassium iodate solution. Drain these samples through the buret and discard them. Fill

the buret with 0.00200 M potassium iodate solution. Adjust the level and make sure the tip is filled.

Record the initial level (to 2 decimal places, 0.00 or 0.05 L) of the solution in the buret on the data

table.

2. Use a transfer pipet to measure out a 10.00 mL portion of the bleach solution. Place the sample in an

Erlenmeyer flask.

3. Into the Erlenmeyer flask, pipet 5.00 mL of 0.600 M potassium iodide solution and about 5 mL of 1.0

M HCl solution.

4. Using the buret, add about 15 mL of 0.00200 M potassium iodate solution. To the Erlenmeyer flask,

pray in the starch indicator. Add enough that the solution has foam in it. Swirl to mix in the indicator.

5. Continue to titrate slowly now, until the first permanent appearance of the blue-black colour.

6. Repeat titration until you have 3 good trials minimum. To determine a “good” trial, calculate the

average of the trials and ensure that each is within +/-5% of that average.

7. Determine and test a SECOND different method to find the vitamin C in a sample. Analysis and Discussion: 1. Calculate the amount of Vitamin C in one tablet and compare it to the theoretical value (for all

methods used).

2. Calculate the % error from manufacturer. 3. Calculate the % error between the different methods.




11 - DETERMINATION OF SALT IN POTATO CHIPS

BY THE MOHR METHOD Reference: Vogel: Quantitative Inorganic Analysis

Introduction:

Purchase uncoloured, unspiced, no sea salt potato chips. Be prepared on the first day to have the

chips so that they sit in water overnight. 1. The Mohr method is a simple way of analyzing for chloride and iodide ions. Since sodium chloride

is the major source of sodium in many foods, the Mohr method may be used to estimate the sodium

chloride content of foods. The Mohr method involves the reaction between silver nitrate solution and

aqueous chloride ions. This is a gravimetric titration as one of the products forms a precipitate. Give

the chemical reaction for this method.

2. The indicator used is potassium chromate solution. It will form an orange precipitate once the

chloride ions have all reacted. Give the chemical reaction for this reaction.

3. Since both reactions involve the formation of precipitates, we need to determine, using Ksp

principles, why the silver chloride precipitates before the silver chromate.

a) For silver chloride, write the dissociation equilibrium of the solid producing the ions. Include the

Ksp expression and value. Using an ICE table, determine the [Ag+] at equilibrium.

b) Another common salt in potato chips is potassium iodide, sea salt. It doesn’t contain sodium ions,

but can react with the silver ion and falsely add to the amount of chloride ion and then the sodium

ion content. For silver iodide, write the dissociation equilibrium of the solid producing the ions.

Include the Ksp expression and value. Using an ICE table, determine the [Ag+] at equilibrium.

c) For silver chromate, the indicator for the end of the titration, write the dissociation equilibrium of

the solid producing the ions. Include the Ksp expression and value. Using an ICE table, determine

the [Ag+] at equilibrium.

d) Rank the 3 anions, as to which has the smallest silver ion concentration and will then react first, to

the last that will react.

e) Explain why you would want chips with a low sea salt content and how that will affect the amount

of sodium chloride determined to be in the chips?

f) What is the difference in the silver ion concentrations between the chromate and chloride

compounds?

g) Explain why you would stop the titration at the earliest colour of light orange detected and how

that will affect the amount of sodium chloride determined to be in the chips? Safety: List any chemicals used or their products that have specific hazards associated with them. Give the

hazards and the precautions you will take. Prediction: 7. If the amount of sodium in a serving of 1 oz or 18 chips, ins 340. mg, what mass of sodium chloride is

in 4 chips?

8. What amount of 0.100 M AgNO3 solution is needed to fully titrate these 2 chips? Procedure:

1. Mass two potato chips. Determine the salt content and the amount of 0.100 M AgNO3 solution

required. If the volume is greater than 20 mL only mass out 1 chip, if the volume is less than 10 mL,

then use 3 or 4 chips until the expected titre is aroun 12 to 15 mL.

2. Place the chip(s) in a 150 mL or 250 mL beaker with 30 mL of distilled water. Cover the potato

chips with a watch glass that fits in the beaker.

3. Leave overnight or cover with a watch glass and heat with a hot plate.


4. Decant or pipet the liquid into a small Erlenmeyer flask, trying to keep a much of the potato chip

mush in the beaker. Push down on the watch glass in the beaker to get as much solution out of the

mush as possible.

5. Rinse the residue with a small amount of distilled water.

6. Decant the rinse water into the flask.

7. Add 3-5 mL of Ag2CrO4 indicator until the solution is a light yellow colour.

8. Prepare for the titration by rinsing the buret with distilled water to check for any clogs or leaks.

Once a buret is good, rinse with the silver nitrate solution and drain through the tip. Fill, using a

funnel, to about the 0 or 1 mL mark. Record all buret volumes to 2 decimal places (0.00 or 0.05 L).

9. Titrate with 0.100 M AgNO3 solution until the colour is a light orange.

10. Repeat until at least 3 consistent trials are completed.

11. Determine and test a SECOND different method to find the amount of salt in 2 potato chips. Analysis and Discussion: 1. Calculate the mass of salt contained in the potato chip(s) for all methods used.

2. Determine the % error with the manufacturer’s claim.

3. Determine the % error between the different methods.



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