Experiment 16
Kinetics: Iodine Clock Reaction rev 1/12
GOAL:
The purpose of this experiment is to determine how temperature and concentrations affect the speed of a
reaction. You will determine initial rate, rate law, temperature dependence, and activation energy.
INTRODUCTION:
How does temperature affect the rate of a reaction? Increasing
temperature always increases rate. Why does this happen? Look
at Figure 1, which shows the energy profile for a typical reaction.
In order for the reactants shown on the left to become products on
the right, they must have enough collision energy to get over the
large energy barrier in the middle called the activation energy, Ea.
If a reaction is run at a low temperature, very few of the reactant
collisions will have enough energy to get over the activation
energy barrier and thus the reaction will be slow. At a high
temperature, many more reactant collisions will have enough
energy to get over the Ea barrier and thus the reaction will be fast.
We can collect quantitative data that will allow us to calculate the
rate of reaction at different temperatures. From these rates, we
Figure 1: Reaction Energy Profile
can calculate values for k, the rate law constant. Since rate
increases with temperature, so does the value of k. The specifics of how k varies with temperature lets us
calculate the activation energy, Ea. This activation energy for a reaction is constant, no matter how we
vary the temperature or concentrations of reactions.
The mathematical relationship between k, Ea, and T is summarized by the Arrhenius Equation, Eqn 1,
⁄
Eqn 1
where k is the rate law constant at Kelvin temperature T, Ea is the activation energy, R is the gas law
constant in joules (8.314 J/mol K), and A is the collision frequency factor. Essentially, A represents the
fraction of the collisions with proper geometry to react and e-Ea/RT represents the fraction of collisions with
sufficient energy to react. We typically rearrange the Arrhenius equation into the form shown as Eqn 2.
It may appear complicated at first, but we can make it much easier to use by noticing that it is in the form
of a straight line, Eqn 3.
y
=
( )
Eqn 2
m (x) + b
Eqn 3
1
If we graph ln k on the y-axis versus 1/T on the x-axis, we will get a straight line with a slope of –Ea/R and
a y-intercept of ln A.
How does the concentration of the reactants affect the rate of a reaction? This relationship can be
complicated. For some reactions, varying concentration has no effect. In other reactions, rate is directly
proportional to concentration and doubling one doubles the other. In still other reactions, rate is
proportional to concentration raised to some other exponent: rate = k [reactant]exponent. By varying the
concentrations of our reactants and seeing the effect on rate, we can figure out the correct exponent for a
particular reaction and reactant. We will use the method of initial rates to find the exponents and values
for k for our reaction. See Section 13.3 of your textbook for more details.
You will determine how varying concentrations affects the rate of reaction for
2 I1-(aq) + S2O82-(aq) → I2(aq) + 2 SO42-
Eqn 4
Since we have two reactants to observe, we will need to find both exponents for the rate law:
rate = k [I1-]m [S2O82-]n
Eqn 5
How do we measure rate? Since one of our products, I2, is colored, we can watch for it. To make the color
change sudden instead of gradual, we will add some thiosulfate, S2O32- and some starch. The first amounts
of I2 formed by our reaction will react with thiosulfate. Once all the thiosulfate is gone, our reaction mixture
will suddenly turn blue due to the I2.
We will add some spectator ions to our solutions to keep the total concentration of ions the same in all our
trials. This is the purpose of the KCl and K2SO4 solutions you will use.
PRE-LAB ASSIGNMENT:
In Trial 1, you will dilute 5.0 mL of a 0.100 M I1- solution to a total volume of 15.0 mL. Calculate the [I1-]
in the final Trial 1 solution. In Trial 2, you will dilute 2.5 mL of 0.100 M I1- solution to a total volume of
15.0 mL. Calculate the [I1-] in the final Trial 2 solution. Recall that you can use M1V1=M2V2 for dilution
calculations. Show these calculations in your pre-lab notebook entry.
HAZARDS:
Although all of the reagents in this experiment should be treated with care, none are particularly hazardous.
Most of the chemicals in today’s experiment have similar hazards, so you will only be asked to investigate
one in detail. Look up the MSDS for potassium iodide (http://hazard.com/msds/index.php or Trexler 464).
Print out a copy of the MSDS and attach it to your prelab pages when you turn them in. Record the
following information in your notebook hazards section:
 appearance and odor
 flammability
 is it a carcinogen? (look in the toxicological section)
 potential health affects (give a brief summary; look in the Hazards section)
 We are using KI as a dilute aqueous solution, under controlled lab conditions. How does this
affect the likelihood that you could have one of the potential health affects you listed above?
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PROCEDURE:
We need to keep our reactants in separate containers until the last moment. We will then start timing as we
mix them and continue until the solution turns color. Use the same flask and test tube for all the reactions.
Clean them between reactions only by rinsing thoroughly with distilled water and allowing them to drain a
minute or so.
Part 1, Concentration Effect
1. Prepare a small table in your notebook with 4 rows and columns labeled: Reaction number and
time. All of the Part 1 reactions are done at room temperature. See the summary table below. Note
that you will have up to 4 solutions combined in your 50 mL flask and up to 2 solutions in your test
tube, depending upon the trial number. We need to keep our reactants in separate containers until
the last moment before we mix and start timing.
2. For Reaction 1, get the indicated solutions in your 50 mL flask and the separate solutions in your
test tube. Pour the contents of the test tube into the flask. Immediately start timing as your swirl the
flask for several seconds. Record the time in seconds required for the reaction to turn blue. After
each reaction, collect and save the used solution in one large beaker. Rinse your flask and test tube
with distilled water before proceeding on to the next reaction.
3. Repeat the procedure for Reactions 2-4.
50 mL flask
Reaction
number
Volume of
KI, mL
1
2
3
4
5.0
2.5
5.0
2.5
Volume
of KCl,
mL
0
2.5
0
2.5
Volume of
Na2S2O3,
mL
3.0
3.0
3.0
3.0
Test Tube
Volume
of starch,
mL
2.0
2.0
2.0
2.0
Volume of
K2S2O8,
mL
5.0
5.0
2.5
2.5
Volume
of K2SO4,
mL
0
0
2.5
2.5
3
Part 2, Temperature Effect
4. Prepare a small table in your notebook with 3 rows and columns labeled: Reaction number,
Temperature, and Time. In this part of the experiment we need to control the temperatures of our
solutions. Since the solutions have been at room temperature for several days, we can use them
directly for a room temperature reaction. To study our reaction at other temperatures, we will
immerse the solutions in a water bath BEFORE mixing them. This procedure will give the
solutions time to reach the desired temperature. After mixing, we will return the mixture to the
water bath to maintain this temperature.
5. See the summary table below. For Reaction 5, get the indicated solutions in your 50 mL flask and
the K2S2O8 in a large test tube. As before, pour the test tube contents into the flask. Immediately
start timing as you swirl the flask for 5-10 seconds to mix the solutions. Record the time that it
takes for the solution to turn blue. Record the temperature of your solution. Empty your used
solutions into your large beaker of used solutions. Save this to hold all your reaction wastes for the
entire experiment. Rinse your flask and test tube with distilled water.
Reaction
number
5
6
7
50 mL flask
Volume of Volume of
KI, mL
Na2S2O3,
mL
5.0
3.0
5.0
3.0
5.0
3.0
Volume
of starch,
mL
2.0
2.0
2.0
Test tube
Volume of
K2S2O8,
mL
5.0
5.0
5.0
Intended
temperature,
C
RT
0
50
6. For Reaction 6, follow the same procedure as for Reaction 5, EXCEPT before mixing, place the
flask and test tube containing your solutions in a slushy ice bath for 5 minutes. Record the
temperature of the ice bath. Pour the contents of the test tube into the flask, swirl for several
seconds. Hold the flask in the ice bath as you measure the time it takes for the solution to turn blue.
7. For Reaction 7, follow the same procedure as for Reaction 5, EXCEPT use the hot water bath
provided by the instructor. Again, allow the separate solutions to stay in the bath for 5 minutes
before mixing, and hold them in the bath to maintain their temperature as you time the reaction. Be
sure to record the temperature of the bath.
8. When you have completed all seven trials, you should have about 100 mL of used solutions
collected in your large beaker. Add a scoop of solid Na2S2O3 and swirl the beaker until it dissolves.
If the blue color disappears, dispose of the solution down the drain with excess water. If blue color
remains, add another scoop of solid Na2S2O3 and swirl to dissolve. Continue adding solid Na2S2O3
until the blue color disappears.
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RESULTS AND DISCUSSION:
Go to the Gen Chem Lab Manual Webpage and find the Excel template for Experiment 15. Save this Excel
file, naming it: Your Name Exp 15. When you open it in Excel, you will find color-coded tables needed to
complete your report. While the format of the tables has been set, you must enter you own data (in the pink
blocks), and use Excel formulas to calculate values (in the yellow blocks). Sometimes you must enter data
that is read from a graph or other sources (in the blue blocks).
Use the file to do your calculations and record your results. Use Excel formulas. When you finish your
report, email the file to your instructor. Be sure that your name is in the file name!
In the Excel file, enter your data in the indicated pink blocks. Note that the table of volumes above and
initial solution concentration data has already been entered to save you some time.
The first table in the spreadsheet requiring calculations looks like the one below. In Excel you will note that
this block is yellow. That means you must enter formulas to calculate the result. Do not directly enter
numbers into yellow blocks!
Reaction
Number
1
Initial Conditions at Start of Reaction
[I1-], M
[S2O82-], M
Initial Rate,
M/sec
k
The initial [I1-] and [S2O82-] are calculated from the concentrations of the stock solutions, the volumes added,
and the final volume of the reaction mixture. All of these numbers are already entered in the spreadsheet.
Recall that we normally use M1V1 = M2V2 when doing a dilution calculation. The I- is supplied by KI so
use the molarity and volume of that solution as M1 and V1 when calculating [I1-]. The S2O82- is supplied by
K2S2O8 so use the molarity and volume of that solution as M1 and V1 when calculating [S2O82-]. In all cases
the final volume V2 is the total reaction volume, 15.0 mL. You will enter a formula into Excel but
rearranged to solve for M2 and with the appropriate cell references. You want to type in a formula for
Reaction 1 and then copy it down for Reactions 2-7. For example, for Reaction 1, under [I1-] you will enter:
=(cell with conc of I1-) *(cell with vol of I1- for Rxn 1)/(cell with total volume). When you copy this
formula down for reactions 2-4, Excel will automatically choose the volumes of Reactions 1-4. Notice,
however, that the concentration of I1- and the total volume is the same for all trials. You want Excel to
always use the same value for these. Thus, you must use an absolute address when you enter the cell
references. See Experiment 13 Excel Step 3d for help with absolute address.
Calculate initial rate using the equation below. Again, don’t do the calculation by hand. You must enter the
formula to have Excel do the calculation. Note that 4.96 x 10-4 is entered as =4.96*10^-4. As before, enter
the formula for Reaction 1 only, and then copy it down the column for Reactions 2-7.
The value of k will be calculated later after you find the rate law exponents.
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One way to determine the order with respect to a reactant (the exponent in the rate law) is to examine how
the initial rate changes as the concentration of the reactant changes. Assume that the initial rate law is given
by the equation:
Initial rate = k [I1-]m[S2O82-]n
Our data will allow us to find the values of the exponents by comparing pairs of trials, as you have done in
lecture. Basically, you compare two trials where only one of the reactant concentrations has varied. Let’s
assume that this concentration has doubled. If rate remains constant despite doubling concentration, the
exponent is zero. If doubling the concentration causes the rate to double, then the exponent is 1. If doubling
the concentration causes the rate to increase by a factor of four, then the exponent is 2. See your textbook or
your lecture notes for more examples of finding the exponents in the method of initial rates. Record whole
number values for m and n.
In Excel, you will find two tables like the one below, one for each exponent. They are color coded blue,
meaning that you enter values that you have found using other data. Enter the four values for m and n that
you found by comparing the indicated reactions. How can you get two different values for the same
exponent? Experimental error can cause this. So, if you get two different values for a variable, simply
choose the one you think more reasonable and use it to continue your report.
m (from comparing rxn 2 to rxn 1) =
m (from comparing rxn 4 to rxn 3) =
m (to be used in the rate equation) =
In your report, write, “Final rate law equation: rate = k [I1-]m[S2O82-]n,” substituting in the whole number
values of m and n you just found.
Now use this rate law equation to have Excel calculate the value of k for each trial. Note that this goes into
the column we skipped in the earlier yellow table. Again, enter a formula for the k of Reaction 1 and then
copy it down for Reactions 2-4. All the values you need are already in the spreadsheet, so you will have a
formula that is all cell address, not specific numbers. You will again need to use absolute address for some
cell references. You will also be raising a number to an exponent. The correct form for doing this is: (cell
with concentration)^(cell with exponent).
All your calculations to this point have used variations in concentration to find the rate law. You also did
several trials that varied the temperature. The same exponents hold at all temperatures, but the value of k is
temperature dependent. The exact nature of this dependence can be used to find Ea, activation energy, an
important characteristic of the reaction.
In your Excel file, continue working down the sheet to complete the table for Part 2 of the experiment.
Complete the first four columns in the same way as you did in the previous table. Use Excel formulas to
calculate T in Kelvin, 1/T and ln(k). Graph ln(k) vs. 1/T with an x-y scatter plot. Be sure that you get ln(k)
on the y-axis. Add a linear trendline and the equation for that line. Enter the values that Excel find for the
slope and intercept in the blue boxes in the spreadsheet.
Reaction Number
5
Measured Temperature, in K
1/T, in K-1
ln(k)
6
Recall that the slope equals -Ea/R and the y-intercept equals ln(A). Use formulas to have Excel calculate Ea
and A from the slope and intercept. Note that you will need the exponential function, written =exp(cell) in
Excel, to calculate A.
Save your Excel file and send it as an attachment via email to your instructor. This Excel file will be
graded.
Print a copy of your Excel spreadsheet.
Your report for this experiment will consist of
1. A cover sheet
2. A printout of your Excel spreadsheet
3. Pages with
a. The final rate law (Eqn 5) with correct values for m and n substituted in
b. Word processed answers to the questions below.
QUESTIONS:
1. What affects the value of k: concentrations, temperature, both, or neither? Explain. Given this,
which of the seven values of k should be the same (assuming no experimental error)?
2. What affects the value of Ea: concentrations, temperature, both, or neither? Explain.
3. In your own words, explain how and why temperature affects rate of reaction. Refer to Figure 1.
Discuss how temperature affects the molecules and how that in turn affects rate. (For help, consult
Figure 13.14 in your textbook.)
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