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Last Revised By R.Tanner and D. Ridge on 15/5/16

SCH 4U PERFORMANCE TASKS 2014/15

A list of investigations from which you may make your choices follows.Choices must be made no later than the first day of Unit 4.

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

Answers to the Introduction questions and the Prediction are due on the FIRST dayof the performance task.

Every student must individually keep organized and detailed Observations and completeall Analysis and Discussion questions. Your observations, graphs, analysis, anddiscussion will be handed in on the final summative day.

On the final day of the performance task, each person will be required to answer questionsabout their specific performance task. During this day, the student may only bring inONE sheet of paper with tables of data – observations and analysis results as well asany graphs – no calculations or written answers may be included.

Number INVESTIGATION1 Acid Content Of Fruit Juice2 Oxidizing Power Of Laundry Bleach3 Blood And Le Châtelier's Principle4 Electroplating5 Entropy Of Reaction6 An Analogy To Fluoride Protection Of Tooth Enamel7 Mol Wt By Freezing Point Depression8 Determination Of The Order Of Reaction9 An Oscillating Reactions10 Rates Of Reaction11 Electrochemical Cells, Thermodynamics and the Equilibrium Constant12 The Rate Of Iodination Of 2-Propanone13 Spectrophotometric Determination Of An Equilibrium Constant14 Ksp Of Copper (II) Tartrate

15 Spectrophotometric Determination Of Aspirin16 Vitamin C - An Important Antioxidant (*)17 Determination of Salt in Potato Chips by the Mohr Method (*)

* May not be completed if a similar experiment was done in SCH 3U1


1 - ACID CONTENT OF FRUIT JUICE

Introduction:1. Try to purchase a lightly coloured juice that contains only citric acid and Vitamin C (ascorbic acid).2. Fruit juices are weak acids and will exist as an equilibrium in the solution.3. The main objective for this lab is to determine which of the acids is the dominant acid that will suppress the

other acids in the juice.4. Obtain the Ka, pKa values, and examples of their pH curves for the acids in your juice.5. How can the Ka values be determined experimentally using a titration curve with a strong base?6. How can that help to determine the dominant acid present in the juice?7. How can the concentration of that acid be determined experimentally using a titration curve with a strong

base?8. Are there other experiments that could be used to determine the amount of the other acid(s)?

Safety:Describe any chemicals used or their products that have specific hazards associated with them. Identify anyhazardous procedures. State all precautions you will take.

Prediction:Use the Ka values and show how you will do the calculations for the:

i) Initial pH Volume of Base = 0 mLii) pH at the Volume 1/2 way to the equivalence pointiii) pH at equivalence point.iv) the concentration for the dominant acid using first the initial pH and the equivalence point.

Assume that the amount of 0.200 M NaOH necessary to titrate 25.00 mL of your 0.150 M fruit juice containingascorbic acid is 18.75 mL.at the equivalence point.

Materials/Equipment:natural fruit juice (recommend: apple) NaOH solution250 mL beakers pH meterpipet burette

Procedure:1. Standardize your NaOH (aq) with KHP to determine the exact concentration prior to titration***2. Pipette 25.00 mL of a fruit juice into a beaker. Measure the pH of the juice and record the value.3. Slowly titrate with the sodium hydroxide in small increments. 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 thepH is above 7 AND the pH values have leveled off.

4. Repeat with the SAME type of juice at least two more times. Trials should be within 5% of each other.5. Repeat using a DIFFERENT technique (research is required – try the iodate titration method).

Hint - consider that pure citric and ascorbic are available in solid form.

Analysis and Discussion:1. Plot pH (y-axis) as a function of the volume of NaOH added (x-axis). Plot all trials on the same graph.2. From the graph, determine the equivalence point from the graph and determine the mL of base required to react

with the acid.3. From the V1/2 equivalence point, determine the Ka value4. Conclude what the dominant acid is in the juice.5. Using the initial pH determine the amount of that acid present. Confirm by using the equivalence point.6. Determine the amount of the non-dominant acid in the juice.***May be done prior to performance task


2 - OXIDIZING POWER OF LAUNDRY BLEACH

Introduction:

Many times in chemistry we must use indirect methods to find the concentration of a chemical in asolution. One such method is as an iodometric analysis. In this experiment the concentration of NaOCl,the active reagent in laundry bleaches, will be found.1. The bleach will udergo a REDOX reaction with iodide to produce what product? Acetic acid is added

to provide the hydrogen ion needed for the reaction. Write the balanced redox reaction of OCl- with I-

ion.2. What colour will the produced iodine change the solution?3. The iodine produced undergoes another redox reaction with the thiosulfate ion. Write this balanced

redox reaction.4. How and why is a starch solution used to help with the end point detection?Combine the equations

from step 1 and step 2. Use this equation for the calculations in the lab.One of the reactantsdisappears – it is used as an indicator with starch. What is this reactant?


Prediction:

1. Assuming that the bleach has a density of 1.08 g/mL convert to molarity of the bleach in mol/L.2. Using the bleach manufacturer’s W/W % of NaOCl, calculate the molar concentration (molarity) of

the NaOCl in the bleach using the bleach concentration from #1.3. Determine the amount of 0.100 M sodium thiosulfate solution necessary to titrate 2.50 mL of your

bleach solution.

Equipment/Materials:

Buret 250 mL Erlenmeyer flask250.00 and 100.00 mL volumetric flasks graduated mL pipetbalance buret clampsolid sodium thiosulfate glacial acetic acidpotassium iodide spray starch

Procedure:

1. To determine the mass of 1.00 mL of the bleach. This value is required to convert the concentration ofthe NaOCl in the bleach into a % W/W.

2. Using a 100.00 mL volumetric flask, prepare a 0.100M solution of sodium thiosulfate3. Using a 250.00 mL volumetric flask, prepare an acetic iodide solution with 2.0 g of potassium iodide

a 10.0 mL of glacial acetic acid ( teacher needs to dispense the acid) and dilute with water to the mark.4. Transfer 50.00 mL of the acidic iodide solution into a 125 mL Erlenmeyer flask.5. Add 1.00 mL of the bleach to the Erlenmeyer flack and swirl to mix. The solution should be

red/brown.6. Rinse the buret with distilled water. Add two small portions of 0.100M sodium thiosulfate solution to

the buret. 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 of thesolution in the buret on the data table.

7. Titrate with the sodium thiosulfate until the solution is light yellow. At this time add the stray starchuntil foam remains in the flask.

8. Titrate slowly until the blue color disappears. Record the amount of sodium thiosulfate solution in theburet at the end of the titration. Record this value in your data table. Note: The blue color mayreappear after the titration has been completed due to air oxidation of the iodide.

9. Titrate several trials until the data is within 5% (3 trials minimum).9. Find the amount of NaOCl in a bleach sample using a SECOND different technique.

Analysis and Discussion:

1. Calculate the concentration of NaOCl in 1.00 mL of the bleach using stoichiometry data from thetitration with sodium thiosulfate.

2. From the concentration of NaOCl, determine the mass of NaOCl in 1.00 mL of bleach.3. Calculate the % W/W of NaOCl in the bleach using the mass of the bleach obtained in procedure step

#1.4. Calculate the percent NaClO in the brand of bleach for all methods used.5. Calculate the percentage error from the manufacturer’s assay.6. Determine all sources of error and their impact on the % W/W obtained.


3 - BLOOD AND LE CHÂTELIER'S PRINCIPLE(Based on Lab from Chemistry 12 by McGraw-Hill Ryerson Publishers)

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.

6. Healthy blood has a pH of 7.4. Estimate the ratio of [CO2] to [HCO3-] in the blood. Use the following

assumptions and information to help you.(a) Assume that carbonic acid and hydrogen carbonate ions are the only contributors to blood pH.

(b) Use Ka = 4.3 X 10-7 for H2CO3To express your answer in terms of CO2 assuming that all undissociated carbonic acid is present as dissolved CO2Safety:Describe any chemicals used or their products that have specific hazards associated with them. Identify anyhazardous procedures. State all precautions you will take.

Prediction:Calculate the pH of a 1:1 buffer of sodium acetate and acetic acid.

Procedure:Create a 1:1 mole buffer that using 0.10 M sodium acetate solution and an UNKNOWN concentration of acetic acid(commercial vinegar).

Calculate the pH of a 1:1 mole buffer.

Determine the concentration of the acetic acid by titrating 25.00 mL of 0.100 M sodium acetate with theUNKNOWN acetic acid until you have reached the desired pH of the 1:1 mole buffer.

Titrate this buffer with 0.10 M HCl (aq) and determine whether this is an effective buffer.

Plot a titration curve and label the buffer region

Repeat this procedure with a new sample of buffer 0.10 M NaOH (aq) and determine whether this is aneffective buffer.

A possible extension (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 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) How much 0.10 M NaOH was needed to make each buffer reach a pH of 7.00? Does that agree with

theoretical (calculated) amount of 0.10 M NaOH?3) Explain why the buffer becomes less effective as more acid (or more base) is added.4) 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 - ELECTROPLATING(Taken from Chemistry 12 by McGraw-Hill Ryerson Publishers)

Introduction:1. What is electroplating?2. Describe several applications of electroplating.3. For potential difference measurements

(a) Draw the circuit diagram for measuring potential difference.(b) Examine the multimeter provided. Draw a diagram showing the physical connections and switch settings

which must be made.In order to prevent burning out the multimeter what should the multimeter rangeselector be set to initially?

4. For current measurements(a) Draw the circuit diagram for measuring current.(b) Examine the multimeter provided. Draw a diagram showing the physical connections and switch settings

which must be made.In order to prevent burning out the multimeter what should the multimeter rangeselector be set to initially?

5. In this investigation, you are to design an electrolytic cell to plate zinc onto a metal object of your choice. Youwill repeat your procedure using three different currents, and then compare your final products. Whendesigning your procedure, consider and discuss questions such as the following.(a) How will you clean the pieces of metal you want to plate?(b) What concentration of electrolyte will you use?(c) Give the appropriate balanced redox equations(d) What external voltage will you use?(e) What three current values will you try?(f) What will you use for the anode in your cell?(g) What will you use for the cathode in your cell?(h) What time limit will you set for the electroplating process?(i) Make clear which variables are controlled, dependent and independent throughout your experimental

modifications.6. Based upon the above design considerations;

(a) Draw a diagram of your basic cell design and discuss the function of all essential components.(b) Write a detailed procedure.


Prediction:Predict the best conditions for plating Zinc onto Copper in terms of current, voltage, and solutions used.

Procedure:

Use the procedure designed in the Introduction (after approval from teacher)

Find and test a SECOND different technique to electroplate Zinc on to Copper.

Analysis and Discussion:1. Compare the mass of Zinc plated to the theoretical (calculated) mass of Zinc that should have been plated (for

all methods used).2. What problems did you encounter when carrying out your investigation? How did you solve these problems?3. Compare the three finished objects that you electroplated. What conditions worked best to electroplate zinc

onto a metal object?4. Were you satisfied with your procedure? What improvements would you make if you did the experiment again?5. Could you build a galvanic cell using the same materials that.you used in your procedure? If your answer is

yes, explain how the galvanic cell would differ from the electrolytic cell that you made in this investigation.


5 - ENTROPY OF REACTION

Introduction:

1. Give the thermodynamic equation for the determination of free energy of a reaction.2. In order for a reaction to be spontaneous, what must the value be for the Gibb’s Free Energy? How

can the above equation be used to determine the minimum entropy change needed to bring about aspontaneous reaction?

3. State the Principle of Heat Exchange as it applies to this experiment.4. How will you calculate H? How will you calculate S?5. What concentration should all solutions be for this experiment? Why?6. Write a detailed procedure, which allows for a minimum of three trials per solid used.

Safety:Describe any chemicals used or their products that have specific hazards associated with them. Identifyany hazardous procedures. State all precautions you will take.

Prediction:

Predict (through calculation) the temperature change that should occur if 5.0 grams of NaNO3 are added

to 100.0 mL of distilled water. Repeat this calculation for NH4Cl and NH4 NO3


solid samples: NaNO3 NH4Cl NH4 NO3, etc.

thermometer Styrofoam cup calorimeter distilled water

Procedure:

Using a calorimeter, carefully measure the temperature increase or decrease during a reaction of anionic compound with water.

Determine a SECOND method for finding the entropy of reaction and complete that procedure also.

Analysis and Discussion:1. Were these reactions spontaneous? How do you know this?2. Calculate the H for each system and compare it to the theoretical value (for all methods used).3. Write a balanced equation for the reactions you studied (including the heat).4. Many students believe that a reaction must be exothermic to be spontaneous. Comment on this in

terms of this experiment.5. Calculate the S for each system and compare it to the theoretical value (for all methods used).


6 - AN ANALOGY TO FLUORIDE PROTECTION OF TOOTH ENAMEL

Introduction:1) How do sugar and lactic acid play a part in tooth decay?2) Give all common chemical reactions that occur with tooth enamel in a healthy mouth that has recently eaten a

sugary snack.3) Show the chemical reaction that occurs if fluoride ions are present in the human mouth.4) Why is acetic acid a good substitute for lactic acid in this experiment?5) Fluoride has been known to inhibit the growth of bacteria. How is that fact relevant to this experiment?


Prediction:If 1.0 g of eggshell is reacted with 5.0 mL of 1.0 M acetic acid, how much CO2 would be produced at SATP?

Generalized Procedure:1. Assemble the apparatus as shown in the diagram below.

5 mlp ipette

water filledflaskacetic

acid

eggsh ell

pinch clamp

2. Use approximately 1.0 g samples of crushed egg shell. Mass the eggshell before the experiment.3. Treat the shell with 2.0 mL aqueous fluoride solutions or toothpaste for 30 minutes. Rinse with distilled water

before using.4. Adjust the water level in the pipette to the 3.00 mL mark before beginning the experiment.5. Add 5.0 mL of 1.0 M acetic acid in a test-tube. Seal the system. Gently tilt the flask containing the acetic acid

test tube until the acid spills out onto the shell. Swirl gently and record the time taken for the carbon dioxideproduced to displace the water in the pipette up to the 0.00 mL mark.

6. Mass the eggshells after the experiment is complete.7. Modify the procedure in order to examine the effect of different fluoride solutes, concentrations and treatment

times on the degree of protection offered by the fluoride solution8. Use bar graphs to present your results.9. Determine a second method for assessing the fluoride protection on the eggshells. Hint – you may use an CO2

detector for your second method.

Analysis and Discussion:1) Explain your results. Has the fluoride solution protected the eggshell? Calculate the amount of eggshell that

reacted (in every possible way available to you).2) Describe the reaction that occurs when acetic acid encounters the eggshell. Explain in words how fluoride may

inhibit that reaction.


7 - MOLECULAR WEIGHT BY FREEZING POINT DEPRESSION

Introduction:

1. How are the melting points and boiling points of solvents different after the addition of a solute?2. Are these changes related to the number of particles present in the solution or to the type of particle

present?3. What do we call these types of properties?4. Give the freezing point depression equation and define all of its terms.5. How can this equation be used to determine a molecular weight?

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:

Find the formula for predicting molecular weight from freezing point depression. Use the molecularweight of Urea to predict how much the freezing point of 100.0 mL of distill water should be depressed.


test tube or large vial foam coffee cup(s)250 or 400 mL beaker icetable salt thermometer or temperature probeurea


Procedure:

Part I: Freezing Point of Pure Water

1. Obtain a clean, dry test tube or vial. Determine the mass of the test tube or vial. Place 10.00 mL ofdistilled water in the test tube or vial, and reweigh. Determine the mass of the water used. Record themass of the water in the data table.

2. Prepare an ice bath in a foam cup with ice and table salt. Place the cup in a beaker to give it morestability. The ice bath should be deep enough so that it is above the level of the water in the test tubeor vial but well below the top. Take care not to let any of the salt or ice get into the sample of distilledwater.

3. Place a thermometer or temperature probe in the distilled water. Take time-temperature data everyhalf-minute until ice has formed in the test tube or vial. It is not necessary to freeze the entire sample.Record the temperature at which the sample froze.

4. Do not discard the sample of the distilled water, because the sample will be used in Part II.


Part II: Molecular Weight of the Unknown

5. Remove the test tube or vial containing the distilled water from the ice bath. Allow the ice to melt.Placing the test tube or vial in a beaker of tap water can speed up this step.

6. Weigh out approximately 1 gram of urea. Record the mass of the sample in the data table. Add theurea to the distilled water, and stir until it is all dissolved. Return the test tube or vial to the ice bath.Insert the thermometer or temperature probe.

7. Take time-temperature data as in Part I. Again, the sample does not have to be frozen solid in order todetermine the freezing point. Record the freezing point in the data table.

8. Repeat the procedure (both Parts I and II).

Part III:

Find and test a SECOND different method to investigate freezing point depression.


1. Using the change in freezing point, the kilograms of water used, and the freezing point constant forwater, calculate the number of moles of urea used in each trial.

2. Using the mass of urea and the number of moles of urea, calculate the molecular weight for each trial.3. Calculate the average molecular weight for urea and the percent error for the trials (for all methods

used).4. What differences would be expected if an ionic compound such as sodium chloride were used instead

of urea?5. Why is it not necessary to wait for the entire sample of water to freeze in order to determine its

freezing point?6. Why is it a good idea to measure the freezing point of the water instead of assuming that its freezing

point is exactly 0.00o C?7. What would have happened if a two-gram sample of urea were used in this experiment?


8 - DETERMINATION OF THE ORDER OF REACTION

Introduction:

1. Write the balanced net ionic equation for the reaction between magnesium and hydrochloric acid.Howwould you expect the concentration of the acid and the amount of magnesium metal to affect the rateof reaction?

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

4. Write the general rate equation for this reaction.Since rate can be expressed as the reciprocal of timetaken for the reaction, rewrite this rate equation in terms of time.Rewrite this rate equation after takingthe 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. Identifyany 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 forthis 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.


9 - AN OSCILLATING REACTION

Reference: Journal of Chemical Education, Nov 1988, p1004http://web.chem.ucsb.edu/~feldwinn/DemoLibrary/DemoPDFs/Demo005.pdf

Introduction:

1) What is the Briggs-Rauscher oscillating reaction?2) What are some possible uses of oscillating reactions?3) Give all of the possible reactions that will occur when hydrogen peroxide, potassium iodate, water,

sulphuric acid, malonic acid, magnesium sulfate, and starch solution are placed together in a flask.4) Give an example of a different oscillating reaction.


Prediction:

Predict the effects of increased temperature, increased concentration of sulphuric acid, and increasedamounts of malonic acid would on this oscillating reaction.

Materials:hydrogen peroxidepotassium iodatesulfuric acid

malonic acidmanganese sulphate monohydratestarch solution

Procedure:1. Ensure all glassware has been thoroughly washed and rinsed with distilled water ONLY!2. Prepare the following:

(a) 40.0 ml of 30% hydrogen peroxide plus 60.0 ml of distilled water(b) 4.3 g KIO3 + 100.0 mL distilled water + 1.3 mL of 6 mol/L sulfuric acid

(c) Mix 0.30 g of soluble starch with 50.0 mL of boiling distilled water and stir until dissolved(d) 1.6 g malonic acid + 0.34 g MnSO4 H2O + 50.0 mL distilled H2O

3. Mix together (c) and (d) – this is now solution called (x)4. Mix together equal volumes of solutions (a), (b) and (x).5. Modify the experiment to determine the effects of various changes on the period of oscillation.6. Design (but do not test) a procedure to create a different oscillating reaction with balanced chemical

equations included.

Analysis and Discussion:1. Attempt to determine the specific balanced equations governing the behaviour of this system.Search

periodicals and reference books for information about the mechanisms of oscillating reactions.3. Explain fully why the modifications that you made to this experiment had the impacts they did (for

all methods used).


10 - RATES OF REACTION

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 colors of permanganate andmanganese 2+ ion be used to determine the rate of reaction?

3. List three factors which might control the rate of this reaction.4. 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?5. How might you examine the effect of temperature on this reaction?

6. This reaction is auto catalyzed; i.e. one of the products (Mn2+) catalyzes the reaction. How might youexamine this effect?

7. Devise and outline a detailed procedure that will give you enough data to plot a graph. What solutionconcentrations will you use? Have this procedure approved by your teacher.


Prediction:If the reaction is first order for all reactants, predict the rate law. Research and find the actual rate law forthis reaction. How will an increase in temperature affect the rate? Explain fully.

Procedure:1. Perform your experiment as you designed it, 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 and possible trends2) Give the experimental rate law for this reaction (for all methods used).3) How does increasing concentration of each reactant affect the rate of reaction?4) How does increasing the temperature of the reactants affect the rate of reaction?


11 - ELECTROCHEMICAL CELLS, THERMODYNAMICS ANDTHE 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 theillustration. 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?


ProcedureCalculate 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:Salt Bridge Construction1. Boil sufficient 0.1 M KNO3 solution to fill a glass U-tube.

2. Remove the solution, and add 1 g of agar-agar per 100-mL of solution to the boiling solution, stirringconstantly until the agar-agar dissolves.

3. Before the solution cools, fill a U-tube with the solution, leaving about a half inch of air space at each end ofthe U-tube and insert cotton plugs moistened with 0.1 M KNO3 solution on each end. The cotton plugs must

protrude from the ends of the U-tube.Cell Construction4. Construct the cell you designed in the Introduction section using large diameter test tubes for each half-cell.

Both half-cells should be immersed in a beaker of water large enough to hold the half-cells and the salt bridge.Because the voltage changes are of the order of 30 mV for the temperature range you will study, be certain thatthe voltmeter you use is sensitive enough to detect these small changes.

Data Collection5. 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 andthe 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 Ke at the varying experimental temperatures.

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


12 - THE RATE OF IODINATION OF 2-PROPANONE1

Introduction:1) What is spectrophotometry? 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 andconstants in this law.For which type of ion; weakly colored or intensely colored, would measurementof light absorbance be most suitable as a means of determining low ion concentrations? Explain. Areall wavelengths of incident light equally effective for measurements of absorbance by a colored ioncomplex? What is the best wavelength for this experiment? Explain.We want any absorbencymeasured using the spectrophotometer to be a result of the colored complex. How might we correctfor absorbance by the cuvette and solvent molecules and other ions present?For the generalelementary reaction:

aA + bB cC One of the factors that affects the rate of a reaction is the concentration of the reactants. How can

this factor be studied? Give the generalized rate law. Are the exponents of this equation necessarily the same as the equation coefficients a and b? What are these exponents called? What is a first order reaction? What is a second order reaction? What is a third order reaction? What is the overall order of a reaction?

7) Under what conditions can the rate of a reaction be followed spectrophotometrically?8) Why can absorbance be used in place of concentration?9) What kind of graph would allow a calculation of the rate of reaction?

Safety:

List any chemicals used or their products that have specific hazards associated with them. Give thehazards and the precautions you will take.

Prediction:In this experiment you will study the kinetics of the reaction between 2- propanone (acetone) and iodine:

CH3COCH3(aq) + I2(aq) CH3COCH2I(aq) + H+(aq) + I-

(aq)

The rate of this reaction is found to be dependent on the concentration of hydrogen ion in solution as wellas presumably the concentrations of the two reactants. Write theoretical rate law expression for thisreaction.

1 1 Adapted from: An Investigation of the Rate of the Reaction Between Iodine andAcetone in Aqueous Solution, Chemistry, Student's Book II, Topics 13-19, NuffieldAdvanced Science, The Nuffield Foundation, 1970


Materials:2.0 M 2-propanone solution 0.010 M I2 solution6.0 M HCl solution a colorimetera set of cuvettes standard thermometerice

Preparing the experiment:Design an experiment to determine the orders of 2-propanone and HCl in the reaction, and confirm the

order (0) of I2.

Design an experiment to determine the activation energy for the reaction. The table below gives asuggested scheme for accomplishing the objectives listed above, but other combinations may be used

2-propanone(mL)

I2(mL)

HCl(mL)

H2O

(mL)

Temperature(°C)

2.0 2.0 2.0 4.0 room4.0 2.0 2.0 2.0 room2.0 1.0 2.0 5.0 room2.0 2.0 1.0 5.0 room6.0 2.0 2.0 0.0 about 104.0 2.0 1.0 3.0 about 30

Procedure:Prepare a "blank"Insert the blank into the spectrophotometer and set the wavelength to 350 nm.Set the absorbance to zero. (or set transmittance to zero with the sample compartment empty and to 100

with the blank in the sample compartment, then switch to absorbance)Measure out all ingredients except 2-propanone into a 50 mL beakerMeasure 2-propanone into a test tubeStart the stopwatch as you quickly pour the 2-propanone into the mixture in the beaker; swirl or stir to

mixPour some of the mixture into a cuvette and place in the colorimeter to measure until %T reaches 100% (

or %A reaches 0)Stop the stopwatch.For the mixtures which are run below room temperature and above room temperature the best

"guesstimate" of the temperature for the reaction is arrived at by measuring the temperature of thesolutions before mixing and then again after the %T is 100%. The average of these temperatures maybe taken as the temperature of the reaction.

Determine and test a SECOND different method to find the rate of iodination of 2-propanone.

Analysis and Discussion:Calculate the diluted initial concentrations of the 2-propanone, I2 and HCl in each mixture.

Use the results from #1 above and the time for each reaction to determine the rate of each reaction.Use your results to determine the order of reaction for each reactant (for all methods used).Write the rate law for the reaction based on your results (from all methods used).Use your rate law and each rate to determine a value for k, the rate constant, in each reaction. Average the

values for the room temperature runs. Since you had limited time in the experiment, two additionalvalues are given below:


k °C4.1 x 10-6 5.03.7 x 10-5 25.0

Use the five values of k to determine Ea by plotting the information as described in the introduction.


13 - SPECTROPHOTOMETRIC DETERMINATIONOF 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.3. State the Beer-Lambert Law. State the significance of all variables and constants in this law.4. For which type of ion; weakly colored or intensely colored, 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 colored

ion complex? What wavelength would be appropriate for this experiment? 400-500 nm6. We want any absorbency measured using the spectrophotometer to be a result of the colored complex.

How might we correct for absorbance by the cuvette and solvent molecules and other ions present?7. 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+.8. Give the equilibrium constant expression for the above reaction.9. If the concentration of the iron solution is much greater than that of the KSCN solution upon mixing;10. 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 colored complex can be used to determine the amount of

Safety:List any chemicals used or their products that have specific hazards associated with them. Give thehazards and the precautions you will take.

Prediction:Find the theoretical equilibrium constant for the reaction between iron (III) nitrate and potassiumthiocyanate. Calculate the concentrations of the iron (III) ion and thiocyanate ion when the system is atequilbrium.

Equipment/Materials:Spectrovis or similar spectrophotometer0.00200 M KSCN 0.200 M Fe(NO3)3

0.00200 M Fe(NO3)3 0.05 M HNO3

burets or pipets 50 mL beakerscuvets

Procedure:Part 1:Preparation of Standard SolutionsThe chart below provides the volumes of reactants needed to prepare the standard solutions.

Solution 0.00200 MKSCN

0.200 MFe(NO3)3

0.05 M HNO3

1 5.0 mL 5.0 mL 15.0 mL2 4.0 mL 5.0 mL 16.0 mL3 3.0 mL 5.0 mL 17.0 mL4 2.0 mL 5.0 mL 18.0 mL5 1.0 mL 5.0 mL 19.0 mL


Part II: Preparation of Equilibrium MixturesUse a buret or pipet to measure the volumes of the reactants listed below. Note that this set ofcombinations uses the more dilute Fe(NO3)3 solution.

Solution 0.00200 M KSCN 0.00200 MFe(NO3)3

0.05 M HNO3

1 1.0 mL 5.0 mL 4.0 mL2 2.0 mL 5.0 mL 3.0 mL3 3.0 mL 5.0 mL 2.0 mL4 4.0 mL 5.0 mL 1.0 mL5 5.0 mL 5.0 mL 0

Part III: Testing the Solutions.1. Turn on the spectrophotometer, and allow it to warm up for approximately 15 minutes.2. Adjust the wavelength to 447 nm..3. With no cuvet in the sample compartment, set the percent transmittance to zero.4. Use a cuvet filled with the 0.05 M HNO3 as the blank. and place this cuvet in the sample

compartment, being sure to properly align it. (The line on the cuvet should match up with the notchon the instrument.) Close the cover.

5. Adjust the absorbance to zero.6. Obtain absorbance readings for each of the other standard solutions7. Obtain the absorbance readings of each of the equilibrium solutions.

Part IV: Alternate Procedure:Determine and test a SECOND different method to find the equilbrium constant.

Analysis and Discussion:1. Prepare your calibration curve and determine the concentration of [Fe(SCN)2+] for each of the

equilibrium trials.2. 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.3. Use these values to calculate the equilibrium constant for each trial.4. Report the average value for the constant (for all methods used).5. Compare this equilibrium constant to the theoretical value.6. Why is the experiment run at a wavelength of 447 nm?7. Why is precipitation in the cuvettes a possible problem for the standard curve?


14 - 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.3. For which type of ion; weakly colored or intensely colored, would measurement of light absorbance

be most suitable as a means of determining low ion concentrations? Explain.4. Are all wavelengths of incident light equally effective for measurements of absorbance by a colored

ion complex? What wavelength would be appropriate for this experiment?5. We want any absorbency measured using the spectrophotometer to be a result of the colored complex.

How might we correct for absorbance by the cuvette and solvent molecules and other ionspresent?What is a solubility product constant?

7. If we wish to determine the Ksp of copper (II) tartrate, what kind of solution must we work with?8. These solutions can be prepared by adding solutions of sodium tartrate to those of copper (II) sulfate.

The copper (II) tartrate will then form a precipitate. The solution that remains is then saturated withrespect to copper (II) tartrate. How must the amount of tartrate ion added compare to that of copper(II) ions present? Why?

9. Give the dissociation equation for copper (II) tartrate.10. Give the Ksp expression.11. How are the concentrations of the copper 2+ and tartrate ions related to each other in a saturated

solution?12. What color are copper 2+ ions in solution.13. Describe how spectrophotometry can be used to determine the copper 2+ ion concentration in the

saturated solution.

Safety:


Prediction:

Find the theoretical Ksp for copper (II) tartrate. Calculate the equilibrium concentrations of the copper

(II) ion and the tartrate ion when there is a saturated solution.


small Erlenmeyer flasks 15-25 mL spectrophotometer and cuvetscentrifuge and centrifuge tubes millipore filter and syringeBeral pipets copper (II) nitratesodium potassium tartrate adjustable pipets (0-2 mL)volumetric flask (100 mL, 10 mL) (a 10 mL graduated cylinder may be used)


Procedure:

Part I: Preparation of a saturated solution of copper (II) tartrate

1. Make 100.00 mL of a 0.100 M copper (II) nitrate solution from the solid.2. Make 100.00 mL of a 0.100 M sodium potassium tartrate solution from the solid.3. Using graduated pipets, place 10.00 mL of 0.100 M copper (II) nitrate and 5.00 mL of 0.100 M

sodium potassium tartrate in a 25.00 mL volumetric flask. Add distilled water to make 25.00 mL ofsolution. Mix well.

4. Allow the solution to remain undisturbed for about 15 minutes while other solutions are beingprepared. The solution should form a precipitate.

5. Centrifuge to remove the precipitate. Save the clear blue solution. If this solution shows anycloudiness 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 graduatedpipet), 5.00 mL of 0.100 M sodium tartrate, and dilute until the total volume is 25.00 mL. See chartbelow.

2. Prepare 25.00 mL of 0.0180 M, 0.015, 0.012, 0.010 M copper (II) tartrate in a similar manner (using avolumetric flask and graduated pipets). See chart below.

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

1. Setup the Spectrovis and follow the instructions to collect the data.2. Zero the absorbance 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 ProcedureDetermine and test a SECOND different method to find the solubility constant for copper (II) tartrate.


1. Prepare your calibration curve and determine the concentration of the copper 2+ ion in the saturatedcopper (II) tartrate solution.

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

3. How does your experimental value compare to the published value for Ksp?

Solution 0.100 M cupper (II) nitrate 0.100 M sodiumpotassium tartrate

0.0200M 2.00 mL 5.00 mL0.0180 M 1.80 mL 5.00 mL0.0150 M 1.50 mL 5.00 mL0.0120 M 1.20 mL 5.00 mL0.0100 M 1.00 mL 5.00 mL


15 - SPECTROPHOTOMETRIC ANALYSIS OF ASPIRIN

This experiment is an adaptation of a lab taken from Experiments in General Chemistry by Weiss,Wismar, and Greco (MacMillan Publishing Co., 1983).

Introduction:

1. What is spectrophotometry? The Beer-Lambert Law relates the percentage of light absorbed by asample to its concentration. State the Beer-Lambert Law. State the significance of all variables andconstants in this law.

3. For which type of ion; weakly colored or intensely colored, would measurement of light absorbancebe most suitable as a means of determining low ion concentrations? Explain fully.

4. Are all wavelengths of incident light equally effective for measurements of absorbance by a coloredion complex? What wavelength would be appropriate for this experiment?

5. We want any absorbency measured using the spectrophotometer to be a result of the colored complex.How might we correct for absorbance by the cuvette and solvent molecules and other ions present? Acolored complex is formed between aspirin and the iron (III) ion when aspirin reacts first with sodiumhydroxide to form the salicylate dianion which is then reacted with acidified iron (III) ion to producethe violet tetraaquosalicylatroiron (III) complex. Research the balanced equations for these tworeactions.Describe how the formation of this colored complex can be used to determine the amount ofaspirin in a commercial aspirin tablet.

Safety:


Prediction:What is the concentration of ASA if 0.400 grams of ASA is added to 250.0 mL of water?Calculate the amount of 0.200 M NaOH that would be needed to completely react with a 0.500 g tablet ofASA.

Equipment / Materials:

6 - 125 mL erlenmeyer flasks commercial aspirin product or aspirin the student has made10 mL graduated cylinder acetylsalicylic acid250 mL volumetric flask 1 M NaOH100 mL volumetric flask 0.02 M iron (III) chloride buffer5 mL pipet spectrophotometer2 cuvettes DI wateranalytical balance (opt.)

Procedure:

Part I: Preparation of the Buffer:1. Dissolve 1.62 g of FeCl3 (Iron (III) chloride) in 500.0 mL of 0.10 M HCl

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, and dilute with distilled water to

the mark.


3. Pipet a 2.50 mL sample of this aspirin standard solution to a 100 mL volumetric flask. Dilute to themark with a 0.020 M iron (III) solution. Label this solution "A," and place it in a 125 mL Erlenmeyerflask.

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.2. Quantitatively transfer the solution to a 250.0 mL volumetric flask, and dilute with distilled water to

the mark.3. Pipet a 2.5 mL sample of this aspirin tablet solution to a 100.0 mL volumetric flask. Dilute to the

mark with a 0.02 M iron (III) solution. Label this solution "unknown," and place it in a 125 mLErlenmeyer flask.

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

3. Obtain absorbance readings for each of the five standard solutions4. 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.


1. Prepare your calibration curve and determine the concentration of the unknown.2. Calculate the amount of ASA in one tablet or aspirin from the data (for all methods used).3. How did the experimental value for ASA compare to the accepted value?4. Explain why the wavelength of 530 nm was used.5. ASA is an important compound in chemistry - explain why using balanced chemical equations.


16 - VITAMIN C: AN IMPORTANT ANTIOXIDANT(Taken from Chemistry 12 by McGraw-Hill Ryerson Publishers)

Introduction:1. Vitamin C is an antioxidant. Why is this important in human health?2. One way to determine the Vitamin C content of a sample is to titrate it with potassium iodide/iodate

system. Determine a way to make a solution of potassium iodide/iodate for this experiment.3. Give an equation illustrating the titration reaction, naming all reactants and products, illustrating their

structures and identify the oxidized and reduced species.4. What color is molecular iodine?5. How will the end point of a titration between Vitamin C and potassium iodide/iodate system be

recognized?6. In storage, the concentration of iodine in solution decreases fairly quickly over time. Why does this

happen? Because the iodine solution's concentration is not stable, it should be standardizedfrequently. How will you standardize the iodide/iodate solution?

7. 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 thehazards and the precautions you will take.

Prediction:What amount of potassium iodide/iodate solution would be needed to fully react with a 0.500 g tablet ofVitamin C?

Procedure:Perform the titration experiment from the introduction (FIRST method) using potassium iodide/iodate.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. If you titrate orange juice that has been exposed to the air for a week, will the Vitamin C

concentration be different from the Vitamin C concentration in fresh juice? If so, will it decrease orincrease? Explain your prediction, in terms of redox reactions.

3. A chemist adds a few drops of deep violet-red iodine solution to a Vitamin C tablet. The iodinesolution quickly becomes colourless. Then the chemist adds a solution that contains chlorine, C12.

The chemist observes that the violet-red colour of the iodine reappears. Explain the chemist'sobservations, in terms of redox reactions.

Further Study:If time permits, extend your investigation to compare the Vitamin C content of fresh juice and juice thathas been left exposed to air for varying periods of time.


17 - DETERMINATION OF SALT IN POTATO CHIPSBY THE MOHR METHOD

Reference: Vogel: Quantitative Inorganic Analysis

Introduction:

1. Sodium chloride, an important nutrient, produces what nutritional problems?2. What is the average Canadian consumption of sodium chloride? How does this compare to what is

needed for good health?3. The Mohr method is a simple way of analyzing for chloride and bromide ions. Since sodium chloride

is the major source of chloride in many foods, the Mohr method may be used to estimate the sodiumchloride content of foods. Describe the general steps involved in a Mohr analysis of chloride ions.

4. Describe the reaction that occurs in a Mohr method titration and explain the reaction in terms ofelectrochemistry,

5. Given the mineral content in a snack pack, how can you compare your results to the manufacturer’sclaims?

Safety:

List any chemicals used or their products that have specific hazards associated with them. Give thehazards and the precautions you will take.

Prediction:

What amount of 0.10 M AgNO3 (aq) is needed to fully titrate 0.020 grams of sodium chloride?

Procedure:

1. Mass two potato chips.2. Place them in a beaker with 30 mL of distilled water.3. Cover and heat.4. Decant the liquid into a flask.5. Rinse the residue with a small amount of distilled water.6. Decant the rinse water into the flask.7. Titrate.8. Repeat until 3 consistent trials are completed9. Determine and test a SECOND different method to find the amount of salt in 2 potato chips.


1. How much salt is contained in a snack-size bag of chips (calculate for all methods used)?2. Determine the percentage error with the manufacturer’s claim.3. What accounted for this error?

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