SCH 4U PERFORMANCE TASKS 09/10
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 two or three – 3 people is the maximum for any group!.
Four days of class time only will be provided for actual completion of the investigation(s) however each
will require preliminary research on your part.
A group report is to be submitted for each investigation. All resources used in your preliminary and any
subsequent research are to be properly referenced in a bibliography. Reports may be hand written
provided they are legible, double-spaced and on one side of the page only. Refer to Appendix A4, Lab
Reports, in your text for guidance on the format of the report. The reports will be due on the Monday
following the allotted lab time.
Number
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
19
INVESTIGATION
Acid Content Of Fruit Juice
Oxidizing Power Of Laundry Bleach
Blood And Le Châtelier's Principle
Electroplating
Entropy Of Reaction
An Analogy To Fluoride Protection Of Tooth Enamel
Mol Wt By Freezing Point Depression
Determination Of The Order Of Reaction
An Oscillating Reactions
Rates Of Reaction
Paper Chromatography Of Food and Candy Dyes
Electrochemical Cells, Thermodynamics and the Equilibrium Constant
The Rate Of Iodination Of 2-Propanone
Spectrophotometric Determination Of An Equilibrium Constant
Kinetic Study of The Reaction of Ki with K2S2O8 Using A Spectrophotometer
Ksp Of Copper (II) Tartrate
Spectrophotometric Determination Of Aspirin
Vitamin C - An Important Antioxidant
Determination of Salt in Potato Chips by the Mohr Method (*)
* May not be completed if this experiment was done in SCH 3U1
Last Revised By R.Tanner and D. Ridge on 09/12/15
ACID CONTENT OF FRUIT JUICE - 1
Introduction:
1. Fruit juices are often used as examples of weak acids. As a weak acid, will more of the acid exist in molecular
form or in ionic form?
2. How might a pH meter be used to determine the amount of acid in ionic form?
3. How might a strong base be used to determine the total amount of acid present?
4. How might this data be used to calculate the degree of ionization?
5. In what range do the percentage ionizations of weak acid generally fall?
Purpose:
The purpose of this experiment is to determine the amount of apparent and total acid in a fruit juice.
Materials/Equipment:
natural fruit juices
250 mL beakers
buffer solution
pipet
magnetic stirrer and stirring bar (optional)
standardized NaOH solution
pH meter
distilled water
burette
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.
Procedure:
1. Check with your teacher and follow their directions regarding the standardization of the pH meter.
2. Pipette 25.0 mL of a fruit juice into a beaker. Measure the pH of the juice and record the value.
3. Place the beaker on the magnetic stirrer. Distilled water may be added so that the stirring bar is submerged.
Make sure that the spinning bar will not hit the electrode of the pH meter.
4. Slowly titrated 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
the pH is above 7 AND the pH values have leveled off.
5. Repeat with the same type of juice two more times.
6. Complete this procedure for at least TWO different types of juice.
7. Plot pH (y-axis) as a function of the volume of NaOH added (x-axis). Determine the equivalence point from
the graph and determine the mL of base required to react with the acid.
8. Repeat for at least one other juice.
Calculations:
1. For each juice sample tested;
(a) Graph the pH data and label the graph fully with all relevant reactions
(b) Calculate the [H3O+] of the juice from the pH value.
(c) Calculate the total acid content of the juice from the titration data.
(d) What percent of the acid was ionized?
(e) Calculate the average percentage ionization.
2. Repeat for each juice.
Questions:
1. Why were the [H3O+] values from the pH reading and the titration calculation different?
2. Of the fruit juices tested, which was the most acidic? least acidic?
3. Why wouldn’t a titration with phenolphthalein be appropriate for a grape juice or tomato juice determination?
Last Revised By R.Tanner and D. Ridge on 09/12/15
Using your understanding from the laboratory analysis answer the following questions about the titration
curve given below:
Concentration of NaOH = 0.200 M
Volume of Acid (HA) = 20.00 mL
Titration of an Unknown Acid
14
12
10
pH
8
6
4
2
0
0
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
Volume of NaOH
Using the graph given above:
1.
2.
3.
4.
Label the graph fully with reactions that explain what is happening throughout the titration.
Determine the %ionization for the acid.
Determine the Ka for the unknown acid.
Determine whether the acid being titrated is formic acid, acetic acid, or valeric acid. Explain your
reasoning fully.
5. Using the Ka value for the unknown acid calculate the pH for the equivalence point and
compare to the experimental equivalence point.
6. Explain why pH at the equivalence point is not equal to 7.
Last Revised By R.Tanner and D. Ridge on 09/12/15
OXIDIZING POWER OF LAUNDRY BLEACH - 2
Introduction:
1. 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. The bleach is reacted with iodide to
produce what product?
2. Acetic acid is added to provide the hydrogen ion needed for the reaction. Give the balanced equation
for this reaction.
3. The amount of product thus formed may then be determined by titration with what reagent solution?
4. Give the balanced equation for this reaction.
5. How and why is a starch solution used to help with the end point detection?
Purpose:
The purpose of this experiment is to determine the NaOCl content of laundry bleaches.
Equipment/Materials:
buret
magnetic stirrer/stirring bar (optional)
250 mL Erlenmeyer flask
250 mL volumetric flask
spoon/spatula
5 mL pipet
wash bottle
vial or small beaker
potassium iodide
beaker for buret waste
100 mL graduated cylinder
0.100 M sodium thiosulfate
10 mL graduated cylinder
balance
buret clamp
ring stand
glacial acetic acid
starch solution
distilled water
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.
Last Revised By R.Tanner and D. Ridge on 09/12/15
Procedure:
1. Use a 5 mL pipet to measure out a sample of a laundry bleach of your choice. Place the bleach in a
vial or a small beaker that has been previously weighed. Determine the mass of the 5 mL of bleach
and its mass per mL. Record these values on your data sheet.
2. Transfer the bleach to a 250 mL volumetric flask. Rinse the original container with distilled water and
add this to the volumetric flask. Bring to volume.
3. Rinse the buret with distilled water. Add two small portions of 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 the
solution in the buret on the data table.
4. Use a 100 mL graduated cylinder to measure out a 50 mL portion of the bleach solution. Place the
sample in an Erlenmeyer flask. Add the stirring bar if one is to be used.
5. Add 2 grams of potassium iodide to the flask. The amount added does not have to be precise. In the
hood, add 10 mL of glacial acetic acid.
6. Titrate with the sodium thiosulfate until the solution is light yellow. At this time add about 5 mL of
the starch indicator.
7. Titrate slowly until the blue color disappears. Record the amount of sodium thiosulfate solution in the
buret at the end of the titration. Record this value in your data table. Note: The blue color may
reappear after the titration has been completed due to air oxidation of the iodide.
8. Titrate three samples of each brand of bleach.
Calculations:
1.
2.
3.
4.
5.
6.
7.
Calculate the moles of sodium thiosulfate used.
Calculate the moles of iodine formed.
Calculate the moles of NaClO in the original sample.
Convert the moles of NaClO to mass of NaClO.
Determine the mass of NaClO per gram of bleach.
Calculate the percent NaClO in the brand of bleach.
Calculate the percentage difference from the manufacturer’s assay.
Questions:
1. For each equation in the introduction section indicate:
(a) species oxidized
(b) species reduced
(c) oxidizing agent
(d) reducing agent
2. Suppose the density of a bleach was found to be 1.07 g/mL. How would the results of an experiment
be affected if a correction for this were not made?
Last Revised By R.Tanner and D. Ridge on 09/12/15
BLOOD AND LE CHÂTELIER'S PRINCIPLE - 3
(Taken from Chemistry 12 by McGraw-Hill Ryerson Publishers)
Introduction:
1.
2.
3.
4.
What is a buffer?
Why are buffer systems extremely important to human health?
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?
5. Give the balanced equations for the blood buffer equilibrium systems.
6. One of the differences between a buffer system in the body and a simple buffer system in a lab is that the body
can enhance the buffer's power to resist pH changes. For example, the lungs regulate the amount of carbon
dioxide that enters and leaves the body. 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.
7. 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 H2CO3
(c) To express your answer in terms of CO2, assume that all undissociated carbonic acid is present as dissolved
CO2
Experiment:
1) Because blood is buffered, it resists changes in pH. To model blood's resistance to changes in pH, design an
investigation that compares the effects of adding an acid or a base to buffered and non-buffered systems. Your
teacher will tell you what materials you may use. The following tips and suggestions will help you get started.
a) You can prepare a simple buffer system by mixing equal volumes of 0.1 mol/L acetic acid, CH3COOH, and
0.1 mol/L sodium acetate, NaCH3COO. Alternatively, you may wish to try more closely simulating the
buffer system in blood. Use soda water, which contains dissolved CO2, and sodium hydrogen carbonate
(baking soda), NaHCO3-. Your teacher may suggest other buffer systems.
b) Consider carrying out your investigation on the micro scale.
c) You will need to decide how to monitor pH changes. You may choose to use a pH meter or a universal
indicator, for example.
d) Decide what type and concentration of acid and base you will add to your systems. Dilute solutions of
strong acids and bases work well.
e) Include a step in which you determine the initial pH of your buffer solution and the unbuffered system.
f) Include all appropriate safety precautions. Obtain your teacher's approval for your procedure, and carry out
your investigation.
2) In your report be sure to compare and contrast the system you used in your investigation with the carbonic
acid/carbonate buffer system in the blood.
a) How are the systems similar?
b) How are they different?
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.
Last Revised By R.Tanner and D. Ridge on 09/12/15
ELECTROPLATING - 4
(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.
(c) In order to prevent burning out the multimeter what should the multimeter range selector 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.
(c) In order to prevent burning out the multimeter what should the multimeter range selector 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. You
will repeat your procedure using three different currents, and then compare your final products. When
designing 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.
Question:
What conditions work best for electroplating zinc onto a metal object? Does changing the current affect the quality
of the finished object?
Hypothesis:
Predict whether the value of the current used to carry out the plating affects the quality of the finished object.
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.
Analysis
1. What problems did you encounter when carrying out your investigation? How did you solve these problems?
2. Compare the three finished objects that you electroplated. What conditions worked best to electroplate zinc
onto a metal object?
3. Were you satisfied with your procedure? What improvements would you make if you did the
experiment again?
4. 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.
Last Revised By R.Tanner and D. Ridge on 09/12/15
ENTROPY OF REACTION - 5
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?
3. How can the above equation be used to determine the minimum entropy change needed to bring about
a spontaneous reaction?
4. Give the basic heat calculation formula.
5. Define specific heat capacity.
6. What are the units for specific heat capacity?
7. What is the specific heat capacity of water?
8. State the Principle of Heat Exchange as it applies to this experiment.
9. How will you calculate ΔH?
10. What concentration should all solutions be for this experiment? Why?
11. Write a detailed procedure, which allows for a minimum of three trials per solid used.
Purpose:
The purpose of this experiment is to estimate the minimum entropy change required for a reaction.
Equipment/Materials:
solid samples:
thermometer
NaNO3
NH4Cl
Styrofoam cup calorimeter
NH4NO3, etc.
distilled water
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.
Procedure:
Perform your experiment.
Questions:
1. Write a balanced equation for the reactions you studied (including the heat).
2. Were these reactions spontaneous? How do you know this
3. Many students believe that a reaction must be exothermic to be spontaneous. Comment on this in
terms of this experiment.
Last Revised By R.Tanner and D. Ridge on 09/12/15
AN ANALOGY TO FLUORIDE PROTECTION OF TOOTH ENAMEL - 6
Introduction:
Many manufacturers of toothpaste add fluoride compounds to their formulations as a means of
strengthening the tooth enamel, making it more resistant to decay. In this activity you will examine the
ability of fluoride ion to protect eggshell from attack by acid. The introduction to your lab report should
detail how fluoride protects the enamel on teeth and outline how you will test your hypothesis.
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.
Generalized Procedure:
1. Assemble the apparatus as shown in the diagram below.
i)
2. Use 1.0 g samples of crushed egg shell.
3. Treat the shell with 2.0 mL aqueous fluoride solutions. for 30 minutes. Rinse with distilled water
before using.
4. Use 5.0 mL of 1.0 M acetic acid.
5. Adjust the water level in the pipette to the 3.00 mL mark before beginning the experiment.
6. Seal the system with the pinch clamp. 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 dioxide produced
to displace the water in the pipette up to the 0.00 mL mark.
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 solution
8. Use bar graphs to present your results.
Analysis:
1. Explain your results. Has the fluoride solution protected the eggshell?
2. Describe the reaction that occurs when acetic acid encounters the eggshell.
3. Describe the reactions that are occurring when fluoride solution is used to protect the eggshell.
Last Revised By R.Tanner and D. Ridge on 09/12/15
MOLECULAR WEIGHT BY FREEZING POINT DEPRESSION - 7
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?
Purpose:
The purpose of this laboratory activity is to determine the molecular weight of urea using the technique of
freezing point depression.
Equipment/Materials:
test tube or large vial
250 or 400 mL beaker
table salt
urea
foam coffee cup(s)
ice
thermometer or temperature probe
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.
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 about 10 mL
of distilled water in the test tube or vial, and reweigh. Determine the mass of the water used. Record
the mass 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 more
stability. The ice bath should be deep enough so that it is above the level of the water in the test tube
or vial but well below the top. Take care not to let any of the salt or ice get into the sample of distilled
water.
3. Place a thermometer or temperature probe in the distilled water. Take time-temperature data every
half-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.
Last Revised By R.Tanner and D. Ridge on 09/12/15
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 the
urea 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 to
determine the freezing point. Record the freezing point in the data table.
8. Repeat the procedure (both Parts I and II).
Calculations:
1. Using the change in freezing point, the kilograms of water used, and the freezing point constant for
water, 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.
Questions:
1. What differences would be expected if an ionic compound such as sodium chloride were used instead
of urea?
2. Why is it not necessary to wait for the entire sample of water to freeze in order to determine its
freezing point?
3. Why is it a good idea to measure the freezing point of the water instead of assuming that its freezing
point is exactly 0o C?
4. What would have happened if a two-gram sample of urea were used in this experiment?
Last Revised By R.Tanner and D. Ridge on 09/12/15
DETERMINATION OF THE ORDER OF REACTION – 8
Introduction:
1. Write the balanced net ionic equation for the reaction between magnesium and hydrochloric acid.
2. 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.
5. Since rate can be expressed as the reciprocal of time taken for the reaction, rewrite this rate equation
in terms of time.
6. 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 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.
Procedure:
1. Perform your experiment.
Analysis:
1. Complete your data table and construct your graph as discussed in the introduction.
2. Based on the discussion in the introduction, give the order of this reaction with respect to hydrogen
ions, and give the value of the rate constant.
Last Revised By R.Tanner and D. Ridge on 09/12/15
AN OSCILLATING REACTION - 9
Reference: Journal of Chemical Education, Nov 1988, p1004
Introduction:
What is an oscillating reaction?
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.
Materials:
hydrogen peroxide
potassium iodate
sulfuric acid
malonic acid
manganese sulphate
starch solution
Procedure:
1. Prepare the following:
(a) 6% hydrogen peroxide
(b) 4.3 g KIO3 + 100 mL water + 1.5 mL 6 mol/L sulfuric acid
(c) 1.56 g malonic acid + 0.34 g MnSO4 .H2O + 100 mL H2O + 3 mL starch solution
2. Mix together equal volumes of solutions (a), (b) and (c).
3. Modify the experiment to determine the effects of various changes on the period of oscillation.
Discussion:
1. Attempt to determine the specific balanced equations governing the behaviour of this system.
2. Search periodicals and reference books for information about the mechanisms of oscillating reactions.
Last Revised By R.Tanner and D. Ridge on 09/12/15
RATES OF REACTION - 10
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.
2. How might the colors of permanganate and manganese 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 you
examine this effect?
7. 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.
Purpose:
The purpose of this investigation is for you to explore the factors affecting the rate of the
permanganate/oxalic acid reaction.
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.
Procedure:
1. Perform your experiment as you designed it, keeping a meticulous, detailed account of the
experiment.
Last Revised By R.Tanner and D. Ridge on 09/12/15
PAPER CHROMATOGRAPHY OF FOOD AND CANDY DYES - 11
Introduction:
1. Define "chromatography”.
2. What is the purpose of doing chromatography?
3. How does chromatography work?
4. What is paper chromatography? Contrast and compare thin layer chromatography with the technique
used in class.
5. What is an Rf value and what are they used for?
6. How is Rf calculated?
In this lab, the separation and identification of the dyes in food colors is performed. Then, by comparing
plates and Rf values, the number of dyes used in the food colors can be determined.
Experiment #1: Food Dyes
Equipment / Materials:
food dyes (McCormick assorted)
1% sodium chloride solution
toothpicks (round)
standard FD&C food colors
100 mL beaker
Chromatography paper
250 mL beaker
ruler
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.
Procedure:
1. Cut a piece of chromatography paper that is no longer than 20 cm in length.
2. Place several drops of each colour of food dyes in separate wells on a spot plate.
3. Using a toothpick, place a small dot of the desired color about 1/2 inch from the bottom of the paper.
It is important to make the spots very small. Usually one touch with a toothpick is sufficient.
4. Repeat until all 4 colors are spotted separately across the bottom of the paper.
5. Place a thin layer of 1% sodium chloride solution in the bottom of the 100 mL beaker.
6. Place the spotted plate in the beaker, spotted end down, leaning it on the side of the beaker (spots must
remain above solvent level).
7. Cover with the 250 mL beaker and allow the solvent to rise to near the top of the plate.
8. Remove plate from solution.
9. Determine Rf values and determine which colours are mixtures.
10. If standard dyes are available, spot a plate with the ones you believe are contained in a mixed food
colour. Spot the food colour on the same plate. Then spot the food color on top of a new spot of
standard dye. Determine if the dyes are the same compound.
Questions:
1. What is the purpose of covering the plate with a 250mL beaker (step 6)?
2. Why is an Rf value necessary to identify a substance on a chromatogram and not just
the distance the spot moved? Why does one component of a mixture move further up the plate than
another?
Last Revised By R.Tanner and D. Ridge on 09/12/15
Experiment #2: Candy Chromotography
Purpose:
The purpose of this experiment is to separate and compare dyes found in two different kinds of candy.
Materials:
1-600 ml beaker
vinegar
capillary tubes
pencil
wool fabric (student provided)
prepared sample of candy coating dye
ruler
chromatography paper (4 x 20 cm)
plastic wrap
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.
Procedure:
Preparation of Candy Dye Solutions:
1. Place candy in a beaker with enough vinegar to cover the top of the candy (use all colors and extra
green.)
2. Heat gently until color coating dissolves.
3. Pour dye solution into a beaker, add a piece of wool cloth, and add a little more vinegar.
4. Boil gently for a few minutes while constantly stirring.
5. Rinse the wool with water.
6. In a clean beaker, place the wool and some household ammonia.
7. Test the solution with litmus paper to make sure it is basic.
8. Concentrate the solution by boiling it to its original volume.
Chromatographic Separation
1. Draw a pencil line 1 cm from the bottom of the chromatography paper.
2. Place the point of the capillary tube into the sample; some solution will be drawn up. Gently and
quickly touch the tip of the capillary tube to the line drawn to make a spot. Allow the spot to dry.
3. Spot twice more, allowing the sample to dry in between applications. Label with pencil at the
bottom the location of each sample.
4. Repeat steps 3 & 4 with the dye from the other candy. (optional – spot standard FD&C dyes as well.
Keep spots 1.5 cm apart on the paper.)
5. Obtain a beaker, add enough vinegar to cover the bottom, and cover the beaker with plastic wrap.
6. Place the spotted paper in the beaker, spotted side down, with only the very top of the paper
touching the side of the beaker. The spots should not be in the vinegar. Cover the beaker with the
plastic wrap.
7. Allow the chromatogram to develop until the solvent (vinegar) nearly reaches the top of the beaker.
Remove the sample from the beaker and draw a line across the paper at the furthest point of the
solvent's progression.
8. Measure the distances traveled by the solvent and each dye, and calculate an Rf value for each
observed color dye contained in the sample.
Analysis:
Explain in detail how you have used chromatography to separate and identify the dyes.
Last Revised By R.Tanner and D. Ridge on 09/12/15
ELECTROCHEMICAL CELLS, THERMODYNAMICS AND THE EQUILIBRIUM
CONSTANT – 12
Introduction:
1. Using two of copper, tin or lead 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.
(b) Identify the anode, cathode and indicate the direction of electron flow.
(c) Calculate ΔE°.
2. Give the Nernst Equation and define all terms.
3. How can ΔG be determined from ΔE?
4. In a plot of ΔG vs T, what thermodynamic quantity is the slope equivalent to?
5. If ΔG, ΔS and T are known, how can ΔH be determined?
6. How can the value of the equilibrium constant, Ke be determined?
Objective:
The quantities ∆G, ∆H, ∆S and Ke are to be calculated from the temperature variation of the measured emf.
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.
Procedure:
Salt Bridge Construction
1. 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, stirring
constantly 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 of
the 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 Construction
4. 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 that
the voltmeter you use is sensitive enough to detect these small changes.
Data Collection
5. Record the voltage and the temperature of the cell. If the potential is negative, reverse the connections. Be
certain that the alligator clips make good contact with the metal strips.
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. Remove the Bunsen burner, and 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.
Analysis:
1. Calculate ∆G for the cell at each of these temperatures.
2. Plot ∆G versus temperature.
3. Caculate ∆S .
4. From the values of ∆G and ∆S, calculate ∆H at 298 K.
Calculate the value of Ke at the varying experimental temperatures.
Last Revised By R.Tanner and D. Ridge on 09/12/15
THE RATE OF IODINATION OF 2-PROPANONE1 - 13
Spectrophotometry Theoretical Considerations:
1.
2.
3.
4.
What is spectrophotometry?
What is visible light?
Why do some substances appear colored?
If light is observed after passing through a 3 cm path length of a 0.10 mol/L colored solution versus
0.20 mol/L of the same solution, in which case will the solution seem darker? Explain.
5. If light is observed after passing through a 3 cm path length of a 0.10 mol/L colored solution, versus a
1 cm path length of the same solution, in which case will the solution seem darker? Explain.
6. The Beer-Lambert Law relates the percentage of light absorbed by a sample to its concentration. State
the Beer-Lambert Law.
7. State the significance of all variables and constants in this law.
8. 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.
9. Are all wavelengths of incident light equally effective for measurements of absorbance by a colored
ion complex? If not, describe how you might establish an effective wavelength to use. Path lengths in
our spectrophotometer are measured in cm or mm whereas the wavelengths of the incident and
transmitted light are measured in nanometers. This is a very large difference in accuracy and it makes
a simple numerical analysis determination of concentration impractical. If we can’t make an accurate
measurement of the path length how might we use the data that the instrument can provide to
determine the concentration of an unknown solution?
11. Why would it be important for the cuvette to be rotated to exactly the same position in the instrument
each that time it is inserted for the taking of a measurement?
12. 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?
Introduction:
1. For the general reaction:
aA + bB → cC
(a) One of the factors that affects the rate of a reaction is the concentration of the reactants. How can
this factor be studied?
(b) Give the generalized rate law.
(c) Are the exponents of this equation necessarily the same as the equation coefficients a and b?
(d) What are these exponents called?
(e) What is a first order reaction?
(f) What is a second order reaction?
(g) What is a third order reaction?
(h) What is the overall order of a reaction?
2. Under what conditions can the rate of a reaction be followed spectrophotometrically?
3. Why can absorbance be used in place of concentration?
4. What kind of graph would allow a calculation of the rate of reaction?
1 1
Adapted from: An Investigation of the Rate of the Reaction Between Iodine and
Acetone in Aqueous Solution, Chemistry, Student's Book II, Topics 13-19, Nuffield
Advanced Science, The Nuffield Foundation, 1970
Last Revised By R.Tanner and D. Ridge on 09/12/15
5. As with concentration, there is a quantitative relationship between reaction rate and temperature, but it
is somewhat more complicated. It is based on the concept of activation energy--the sufficient energy
needed for an effective collision which can initiate reaction.
(a) What is the equation relating the rate constant k to the kelvin temperature T and the activation
energy Ea?
(b) How can this equation be used to determine the activation energy? This equation is in the form of a
straight line where y is the natural log of k (ln k) and x is 1/T. The slope is thus interpreted as -Ea/R. By measuring k at
different temperatures we can graphically determine the activation energy. The constant is, of course, the y-intercept,
and is specific for a particular reaction. It is known as the Arrhenius Constant.
6. 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
well as presumably the concentrations of the two reactants. Write theoretical rate law expression for
this reaction.
A convenient way to measure the rate for this reaction is in terms of the disappearance of I2. Since the
reaction turns out to be zero order with respect to iodine concentration (i.e., n = 0), the rate of the reaction
does not depend on I2 and we can study the rate by making iodine the limiting reagent, having the 2propanone and hydrogen ion present in large excesses so that their concentrations do not change
appreciably by the time all the iodine is gone and therefore the rate remains fairly constant. Under such
circumstances, if it takes t seconds for the color of the I2 to disappear, the rate would be:
rate = [I2]0/t
where [I2]0 is the initial iodine concentration in the mixture. In principle simply looking at the solution,
comparing it to water, could do this. There are a couple of problems with this approach. As the reaction
approaches completion, the color fades rather rapidly, becoming pale yellow. It is difficult to distinguish
just when the color is gone. Also, the final mixture is not perfectly clear like water. Therefore we will use
a spectrophotometer to monitor the course of the reaction. We will be measuring the time it takes for the
%T to become 100%. If the time is measured from the moment the solutions are mixed, it is not even
important to have the sample in the colorimeter immediately, so long as it is in place before the %T
becomes 100%.
Materials:
2.0 M 2-propanone solution
6.0 M HCl solution
a set of cuvettes
ice
Last Revised By R.Tanner and D. Ridge on 09/12/15
0.010 M I2 solution
a colorimeter
standard thermometer
Preparing to experiment:
1. Design an experiment to determine the orders of 2-propanone and HCl in the reaction, and confirm the
order (0) of I2.
2. Design an experiment to determine the activation energy for the reaction. The table below gives a
suggested scheme for accomplishing the objectives listed above, but other combinations may be used
2-propanone
(mL)
2.0
4.0
2.0
2.0
6.0
4.0
I2
(mL)
2.0
2.0
1.0
2.0
2.0
2.0
HCl
(mL)
2.0
2.0
2.0
1.0
2.0
1.0
H2O
(mL)
4.0
2.0
5.0
5.0
0.0
3.0
Temperature
(°C)
room
room
room
room
about 10
about 30
Technique:
1. Prepare a "blank"
2. Insert the blank into the spectrophotometer and set the wavelength to 350 nm.
3. 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)
4. Measure out all ingredients except 2-propanone into a 50 mL beaker
5. Measure 2-propanone into a test tube
6. Start the stopwatch as you quickly pour the 2-propanone into the mixture in the beaker; swirl or stir to
mix
7. Pour some of the mixture into a cuvette and place in the colorimeter to measure until %T reaches
100% ( or %A reaches 0)
8. Stop the stopwatch.
9. 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 the
solutions before mixing and then again after the %T is 100%. The average of these temperatures may
be taken as the temperature of the reaction.
Safety:
List any chemicals used or their products that have specific hazards associated with them. Give the
hazards and the precautions you will take.
Analysis:
1.
2.
3.
4.
5.
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.
Write the rate law for the reaction based on your results.
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 additional
values are given below:
k
°C
4.1 x 10-6 5.0
3.7 x 10-5 25.0
6. Use the five values of k to determine Ea by plotting the information as described in the introduction.
Last Revised By R.Tanner and D. Ridge on 09/12/15
SPECTROPHOTOMETRIC DETERMINATION
OF AN EQUILIBRIUM CONSTANT - 14
Spectrophotometry Theoretical Considerations:
1. What is spectrophotometry?
2. What is visible light?
3. Why do some substances appear colored?
4. If light is observed after passing through a 3 cm path length of a 0.10 mol/L colored solution versus 0.20 mol/L
of the same solution, in which case will the solution seem darker? Explain.
5. If light is observed after passing through a 3 cm path length of a 0.10 mol/L colored solution, versus a 1 cm
path length of the same solution, in which case will the solution seem darker? Explain.
6. The Beer-Lambert Law relates the percentage of light absorbed by a sample to its concentration. State the BeerLambert Law.
7. State the significance of all variables and constants in this law.
8. 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.
9. Are all wavelengths of incident light equally effective for measurements of absorbance by a colored ion
complex? If not, describe how you might establish an effective wavelength to use. Path lengths in our
spectrophotometer are measured in cm or mm whereas the wavelengths of the incident and transmitted light are
measured in nanometers. This is a very large difference in accuracy and it makes a simple numerical analysis
determination of concentration impractical. If we can’t make an accurate measurement of the path length how
might we use the data that the instrument can provide to determine the concentration of an unknown solution?
11. Why would it be important for the cuvette to be rotated to exactly the same position in the instrument each that
time it is inserted for the taking of a measurement?
12. 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?
Introduction:
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+.
2. Give the equilibrium constant expression for the above reaction
3. If the concentration of the iron solution is much greater than that of the KSCN solution upon mixing;
(a) Will the reaction establish equilibrium or go to completion?
(b) How can the concentration of the product be determined from the volume and concentration of the KSCN
used in each trial?
4. Describe how the formation of this colored complex can be used to determine the amount of [Fe(SCN)]2+ in an
equilibrium mixture
1.
Purpose:
The purpose of this experiment is to determine a value for the equilibrium constant for the reaction between iron
(III) nitrate and potassium thiocyanate.
Equipment/Materials:
Spec 20 or similar spectrophotometer
0.00200 M KSCN
0.00200 M Fe(NO3)3
burets or pipets
cuvets
Last Revised By R.Tanner and D. Ridge on 09/12/15
0.200 M Fe(NO3)3
0.05 M HNO3
50 mL beakers
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.
Procedure:
Part 1:
Preparation of Standard Solutions
1. The chart below provides the volumes of reactants needed to prepare the standard solutions.
Solution
1
2
3
4
5
0.00200 M KSCN
5.0 mL
4.0 mL
3.0 mL
2.0 mL
1.0 mL
0.200 M Fe(NO3)3
5.0 mL
5.0 mL
5.0 mL
5.0 mL
5.0 mL
0.05 M HNO3
15.0 mL
16.0 mL
17.0 mL
18.0 mL
19.0 mL
Part II:
Preparation of Equilibrium Mixtures
2. Use a buret or 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
1
2
3
4
5
0.00200 M KSCN
1.0 mL
2.0 mL
3.0 mL
4.0 mL
5.0 mL
0.00200 M Fe(NO3)3
5.0 mL
5.0 mL
5.0 mL
5.0 mL
5.0 mL
0.05 M HNO3
4.0 mL
3.0 mL
2.0 mL
1.0 mL
0
Part III: Testing the Solutions.
2. Turn on the spectrophotometer, and allow it to warm up for approximately 15 minutes.
3. Adjust the wavelength to 447 nm..
4. With no cuvet in the sample compartment, set the percent transmittance to zero.
5. 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 notch on the instrument.)
Close the cover.
6. Adjust the absorbance to zero.
7. Obtain absorbance readings for each of the other standard solutions
8. Obtain the absorbance readings of each of the equilibrium solutions.
Calculations:
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.
Questions:
1. Why is the experiment run at a wavelength of 447 nm?
Last Revised By R.Tanner and D. Ridge on 09/12/15
KINETIC STUDY OF THE REACTION OF KI WITH K2S2O8
USING A SPECTROPHOTOMETER - 15
Spectrophotometry Theoretical Considerations:
1. What is spectrophotometry?
2. What is visible light?
3. Why do some substances appear colored?
4. If light is observed after passing through a 3 cm path length of a 0.10 mol/L colored solution versus 0.20 mol/L
of the same solution, in which case will the solution seem darker? Explain.
5. If light is observed after passing through a 3 cm path length of a 0.10 mol/L colored solution, versus a 1 cm
path length of the same solution, in which case will the solution seem darker? Explain.
6. The Beer-Lambert Law relates the percentage of light absorbed by a sample to its concentration. State the BeerLambert Law.
7. State the significance of all variables and constants in this law.
8. 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.
9. Are all wavelengths of incident light equally effective for measurements of absorbance by a colored ion
complex? If not, describe how you might establish an effective wavelength to use. Path lengths in our
spectrophotometer are measured in cm or mm whereas the wavelengths of the incident and transmitted light are
measured in nanometers. This is a very large difference in accuracy and it makes a simple numerical analysis
determination of concentration impractical. If we can’t make an accurate measurement of the path length how
might we use the data that the instrument can provide to determine the concentration of an unknown solution?
11. Why would it be important for the cuvette to be rotated to exactly the same position in the instrument each that
time it is inserted for the taking of a measurement?
12. 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?
Introduction:
1. For the general reaction:
aA + bB → cC
(a) One of the factors that affects the rate of a reaction is the concentration of the reactants. How can
this factor be studied?
(b) Give the generalized rate law.
(c) Are the exponents of this equation necessarily the same as the equation coefficients a and b?
(d) What are these exponents called?
(e) What is a first order reaction?
(f) What is a second order reaction?
(g) What is a third order reaction?
(h) What is the overall order of a reaction?
2. Under what conditions can the rate of a reaction be followed spectrophotometrically?
3. Why can absorbance be used in place of concentration?
4. What kind of graph would allow a calculation of the rate of reaction?
5. Give the balanced equation for the reaction between KI and K2S2O8.
6. Which product will be the colored species monitored in this experiment?
Last Revised By R.Tanner and D. Ridge on 09/12/15
Purpose:
The purpose of this experiment is to determine the rate of the reaction of potassium iodide and potassium
persulfate and to determine the order of each reactant and the overall order of the reaction.
Equipment/Materials:
1 mL pipet (or 10 mL graduated cylinder)
0.020 M KI solution
2 small test tubes
0.020 M K2S2O8
spectrophotometer
cuvettes
distilled water
parafilm
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.
Procedure:
1. Turn on the spectrophotometer.
2. Place 2 mL of KI and 2 mL of water into a cuvet to serve as the blank.
3. Insert the blank into the spectrophotometer and set the wavelength to 350 nm.
4. 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)
5. Place the required amount of KI and water into a cuvette and place the required amount of potassium
persulfate into a small test tube. (See chart of suggested amounts.)
Vol. KI
Vol. H2O
Set # 1
1 mL
1 mL
Set # 2
2 mL
1 mL
Set # 3
1 mL
2 mL
Set # 4
2 mL
0 mL
Vol. K2S2O8
2 mL
1 mL
1 mL
2 mL
6. Quickly pour the contents of the test tube into the cuvette, cover the cuvette with parafilm and invert
three times. Remove the cover from the sample compartment of the spectrophotometer, insert the
cuvette, and recover quickly.
7. Record the absorbance every 10 seconds for 2 minutes. Consider the absorbance at time 0 to be 0.
Begin taking reading 10 seconds after pouring the contents of the test tube into the cuvette.
8. When the run is complete, remove the cuvette and clean thoroughly. Prepare to do the second set.
Repeat steps 4-6 using as many sets as assigned by the instructor.
9. Plot each run on the same graph. Determine the slope of each run and the rate of the reaction.
Calculations:
1. Determine the order of reaction with respect to KI and K2S2O8
2. Determine a value for the rate constant.
3. Write the rate law expression.
Questions:
1. Why is the wavelength set to 350 nm?
Last Revised By R.Tanner and D. Ridge on 09/12/15
KSP OF COPPER (II) TARTRATE - 16
Spectrophotometry Theoretical Considerations:
1. What is spectrophotometry?
2. What is visible light?
3. Why do some substances appear colored?
4. If light is observed after passing through a 3 cm path length of a 0.10 mol/L colored solution versus 0.20 mol/L
of the same solution, in which case will the solution seem darker? Explain.
5. If light is observed after passing through a 3 cm path length of a 0.10 mol/L colored solution, versus a 1 cm
path length of the same solution, in which case will the solution seem darker? Explain.
6. The Beer-Lambert Law relates the percentage of light absorbed by a sample to its concentration. State the BeerLambert Law.
7. State the significance of all variables and constants in this law.
8. 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.
9. Are all wavelengths of incident light equally effective for measurements of absorbance by a colored ion
complex? If not, describe how you might establish an effective wavelength to use. Path lengths in our
spectrophotometer are measured in cm or mm whereas the wavelengths of the incident and transmitted light are
measured in nanometers. This is a very large difference in accuracy and it makes a simple numerical analysis
determination of concentration impractical. If we can’t make an accurate measurement of the path length how
might we use the data that the instrument can provide to determine the concentration of an unknown solution?
11. Why would it be important for the cuvette to be rotated to exactly the same position in the instrument each that
time it is inserted for the taking of a measurement?
12. 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?
Introduction:
1. What is a solubility product constant?
2. If we wish to determine the Ksp of copper (II) tartrate, what kind of solution must we work with? .
3. 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 with respect to
copper (II) tartrate. How must the amount of tartrate ion added compare to that of copper (II) ions present?
Why?
4. Give the dissociation equation for copper (II) tartrate.
5. Give the Ksp expression.
6. How are the concentrations of the copper 2+ and tartrate ions related to each other in a saturated solution?
7. What color are copper 2+ ions in solution.
8. Describe how spectrophotometry can be used to determine the copper 2+ ion concentration in the saturated
solution.
Purpose:
The purpose of this experiment is to determine the solubility product constant of copper (II) tartrate.
Equipment/Materials:
small Erlenmeyer flasks 15-25 mL
spectrophotometer and cuvets
centrifuge and centrifuge tubes
millipore filter and syringe
Beral pipets
copper (II) sulfate (0.100 M)
sodium tartrate (0.100 M)
adjustable pipets (0-2 mL)
volumetric flask (100 mL, 10 mL) (a 10 mL graduated cylinder may be used)
Last Revised By R.Tanner and D. Ridge on 09/12/15
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.
Procedure:
Part I: Preparation of a saturated solution of copper (II) tartrate
1. Place 4.0 mL of 0.100 M copper (II) sulfate and 5.0 mL of 0.100 M sodium tartrate in a 10.00 mL volumetric
flask. Add distilled water to make 10 mL of solution. Mix well.
2. Allow the solution to remain undisturbed for about 15 minutes while other solutions are being prepared. The
solution should form a precipitate.
3. Centrifuge to remove the precipitate. Save the clear blue solution. If this solution shows any cloudiness or
further precipitates, filter it with a millipore filter.
Part II: Preparation of standard copper (II) tartrate solutions
1. Prepare 10 mL of 0.02 M copper (II) tartrate. Measure 2.00 mL of 0.100 M copper (II) sulfate. Add 5.0 mL of
0.100 M sodium tartrate, and dilute until the total volume is 10.00 mL.
2. Prepare 10 mL of 0.018 M copper (II) tartrate. Measure 1.8 mL of 0.100 M copper (II) sulfate. Add 5.0 mL of
0.100 M sodium tartrate, and dilute until the total volume is 10.00 mL.
3. Prepare 10 mL of 0.015, 0.012, 0.010 M copper (II) tartrate in a similar manner.
Part III: Determination of copper (II) ion concentration in the saturated copper (II) tartrate solution
Turn on the spectrophotometer and allow to warm up for 15 minutes.
Adjust the wavelength to 675 nm.
With the sample compartment empty, set the instrument to 0%T.
Zero the absorbance using a blank made by diluting 5.0 mL of 0.1 M sodium tartrate to 10 mL with distilled
water.
5. Determine the absorbance of each of the five standard copper solutions.
6. Place the saturated copper (II) tartrate solution in a cuvette and record the absorbance of this solution.
1.
2.
3.
4.
Calculations:
1. Prepare your calibration curve and determine the concentration of the copper 2+ ion in the saturated copper (II)
tartrate solution.
2. Calculate the value for the Ksp of copper (II) tartrate.
Questions:
1. Explain why a wavelength of 675 nm was used?
2. How does your experimental value compare to the published value for Ksp?
Last Revised By R.Tanner and D. Ridge on 09/12/15
SPECTROPHOTOMETRIC ANALYSIS OF ASPIRIN - 17
This experiment is an adaptation of a lab taken from Experiments in General Chemistry by Weiss, Wismar, and
Greco (MacMillan Publishing Co., 1983).
Spectrophotometry Theoretical Considerations:
1. What is spectrophotometry?
2. What is visible light?
3. Why do some substances appear colored?
4. If light is observed after passing through a 3 cm path length of a 0.10 mol/L colored solution versus 0.20 mol/L
of the same solution, in which case will the solution seem darker? Explain.
5. If light is observed after passing through a 3 cm path length of a 0.10 mol/L colored solution, versus a 1 cm
path length of the same solution, in which case will the solution seem darker? Explain.
6. The Beer-Lambert Law relates the percentage of light absorbed by a sample to its concentration. State the BeerLambert Law.
7. State the significance of all variables and constants in this law.
8. 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.
9. Are all wavelengths of incident light equally effective for measurements of absorbance by a colored ion
complex? If not, describe how you might establish an effective wavelength to use. Path lengths in our
spectrophotometer are measured in cm or mm whereas the wavelengths of the incident and transmitted light are
measured in nanometers. This is a very large difference in accuracy and it makes a simple numerical analysis
determination of concentration impractical. If we can’t make an accurate measurement of the path length how
might we use the data that the instrument can provide to determine the concentration of an unknown solution?
11. Why would it be important for the cuvette to be rotated to exactly the same position in the instrument each time
it is inserted for the taking of a measurement?
12. 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?
Introduction:
1. A colored 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 the balanced equations for these two reactions.
2. Describe how the formation of this colored complex can be used to determine the amount of aspirin in a
commercial aspirin tablet.
Purpose:
The purpose of this lab is to determine the amount of aspirin in a commercial aspirin product.
Equipment / Materials:
6 - 125 mL erlenmeyer flasks
10 mL graduated cylinder
250 mL volumetric flask
100 mL volumetric flask
5 mL pipet
2 cuvettes
analytical balance (opt.)
commercial aspirin product or aspirin the student has made
acetylsalicylic acid
1 M NaOH
0.02 M iron (III) chloride buffer
spectrophotometer
DI water
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.
Last Revised By R.Tanner and D. Ridge on 09/12/15
Procedure
Part I: Preparation of the Buffer
1. Dissolve 6.48g of FeCl3 in 2 L of 0.1M HCl
Part II: Making Standards.
1. Mass 400 mg of acetylsalicylic acid in a 125 mL Erlenmeyer flask. Add 10 mL of a 1 M NaOH solution to
the flask, and heat until the contents begin to boil.
2. Quantitatively transfer the solution to a 250 mL volumetric flask, and dilute with distilled water to the
mark.
3. Pipet a 2.5 mL sample of this aspirin standard solution to a 100 mL volumetric flask. Dilute to the mark
with a 0.02 M iron (III) solution. Label this solution "A," and place it in a 125 mL Erlenmeyer flask.
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 mL of a 1 M NaOH solution to the flask, and
heat until the contents begin to boil.
2. Quantitatively transfer the solution to a 250 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 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 mL Erlenmeyer flask.
Part IV: Testing the Solutions.
Turn on the spectrophotometer and allow the spec to warm up for 15 minutes.
Adjust the wavelength to 530 nm.
With the sample compartment empty, set the instrument to 0%T.
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.
5. Set the absorbance to 0.000.
6. Obtain absorbance readings for each of the five standard solutions
7. Measure and record the absorbance of the unknown.
1.
2.
3.
4.
Analysis:
1. Prepare your calibration curve and determine the concentration of the unknown.
Questions:
1. Explain why the wavelength of 530 nm was used.
2. How did the concentration of your aspirin solution compare to the accepted value?
3. Is it better to buy generic or brand name aspirin? Support your conclusion.
Last Revised By R.Tanner and D. Ridge on 09/12/15
VITAMIN C - AN IMPORTANT ANTIOXIDANT - 18
(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 Wagner’s Reagent.
3. Give the balanced equilibrium equation for the formation of Wagner’s Reagent.
4. Give an equation illustrating the titration reaction, naming all reactants and products, illustrating their
structures and identify the oxidized and reduced species.
5. What color is molecular iodine?
6. What color are iodide ions?
7. How will the end point of a titration between Vitamin C and Wagner’s Reagent be recognized?
8. In storage, the concentration of iodine in solution decreases fairly quickly over time. Why do you
think this happens?
9. Because the iodine solution's concentration is not stable, it should be standardized frequently. How
will you do so?
10. Explain why the Vitamin C used for standardization must be fresh.
11. Design a titration procedure that uses Wagner’s Reagent to determine and compare the concentration
of vitamin C in a variety of fresh fruit juices. You may find it helpful to research appropriate
procedures on the Internet. Include a method for standardizing the iodine solution. Check your
procedure with your teacher, and then carry out your titration.
Safety:
List any chemicals used or their products that have specific hazards associated with them. Give the
hazards and the precautions you will take.
Questions:
1. 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 or increase?
Explain your prediction, in terms of redox reactions.
2. A chemist adds a few drops of deep violet-red iodine solution to a vitamin C tablet. The iodine
solution 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's
observations, in terms of redox reactions.
Further Study:
If time permits, extend your investigation to compare the vitamin C content of fresh juice and juice that
has been left exposed to air for varying periods of time. Ask your teacher to approve your procedure
before carrying out this extension.
Last Revised By R.Tanner and D. Ridge on 09/12/15
DETERMINATION OF SALT IN POTATO CHIPS BY THE MOHR METHOD - 19
Reference: Vogel: Quantitative Inorganic Analysis
INTRODUCTION:
1.
2.
3.
4.
Sodium chloride, an important nutrient, produces what nutritional problem?
What is the average Canadian consumption of sodium chloride?
How does this compare to what is needed for good health?
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 sodium
chloride content of foods. Describe the general steps involved in a Mohr analysis of chloride ions.
5. Describe the reaction that occurs in a Mohr method titration and explain the reaction in terms of
electrochemistry,
6. Given the mineral content in a snack pack can you compare your results to the manufacturer’s claims?
PURPOSE:
The purpose of this investigation is to experimentally determine the amount of salt present in potato chips
compared to the manufacturer’s claim.
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.
PROCEDURE:
1.
2.
3.
4.
5.
6.
7.
8.
9.
Mass two potato chips.
Place them in a beaker with 30 mL of distilled water.
Cover and heat.
Decant the liquid into a flask.
Rinse the residue with a small amount of distilled water.
Decant the rinse water into the flask.
Titrate.
Repeat until 3 consistent trials are completed
Complete this procedure for at least 2 brands of potato chips
ANALYSIS/DISCUSSION:
1.
2.
3.
4.
How much salt is contained in a snack-size bag of chips?
How does one brand compare to another?
How do other foods (eg. bread, popcorn, soup) compare to chips in salt content?
Determine the percentage difference with the manufacturer’s claim.
Last Revised By R.Tanner and D. Ridge on 09/12/15