Name:
AP Chemistry
Date:
Period:
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Lab 11 – The Iodine Clock
Introduction:
In this experiment you will determine the Rate Law for the following reaction:
2 H+(aq) + 2 I (aq) + H2O2(aq)  I2 (aq) + 2 H2O (l)
How fast the reaction produces iodine indicate the rate of reaction through the following expression:
d[I2 ]
Rate 
dt
The rate or speed of the reaction is dependent on the concentrations of iodide ion (I-) and hydrogen peroxide,
H2O2. (The spectator ions are left off the reaction.) Therefore, we can write the Rate Law (concentration
dependence) for the reaction as:
 Rate = k[H+]x[I-]y[H2O2]z
Where: x is the order of the reaction in H+, y is the order of the reaction in I-, z is the order of the reaction in
H2O2, and k is the rate constant. Since the Rate Law is empirical, we have to go to the lab to make
measurements that will enable these values to be calculated.
As with most kinetics, the concentration of reactants or products at any instant is difficult to measure directly,
so in this lab the rate will be determined indirectly. We have a very helpful test for the presence of one of the
products, iodine (I2), namely starch. Starch reacts with iodine to form a blue/black colored complex.
Unfortunately as soon as any iodine is produced it will react to make the complex and the solution will turn
blue/black instantaneously. Thus, using starch as an indicator by itself would not be of much help. It confirms
that some amount of I2 is being formed, but it tells us nothing about what we are trying to measure - the rate
(how long it takes to produce a given quantity of I2.)
To get around this problem we will introduce a side reaction that will remove the initial I2 that is produced by
our main reaction. This will prevent the solution from turning black for long enough so that we can make some
good time measurements. We will use the following side reaction:
I2 (aq) + 2 S2O32—(aq)  2 I (aq) + S4O62— (aq)
S2O32—, thiosulfate ion, reacts with I2, which prevents the solution from turning blue/black. Since we have
carefully measured the amount of the thiosulfate, we know exactly how much iodine it will take to react with
this thiosulfate. As soon as all of the thiosulfate has reacted, iodine will then be left over and then the reaction
will change color. By putting in this "time delay", we can now calculate the rate of the overall reaction using
the equation:
1
[S2O32 ]
Rate of Re action  
2 time for blue to appear
From the rate of the chemical reaction, we can determine the values of x, y, and z. Furthermore, once the overall
order of the reaction is known, the rate constant can be determined.

Pre – Lab
1. Copy the data table and all calculations that will be necessary into your lab notebook.
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AP Chemistry
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Date:
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2. According to the reaction set up in the manual, refer back to the reagent table in order to calculate the number
of moles of iodine ion that have reacted by the time the color of the reaction changes.
3. If the rate of a reaction doubles when the concentration of a certain reactant is doubled, what is the partial
order for the reactant?
4. If the rate of a reaction quadruples when the concentration of a reactant is doubled, what is the partial order
for the reactant?
Materials:
Table 1: Equipment and Materials for Iodine Clock Lab
Materials
Quantity
Volumetric Pipet, 5 mL
2
Volumetric Pipet, 10 mL
2
Erlenmeyer Flask, 125 or
4
250 mL
Beakers, 100 or 150 mL
4
Stop Watch
1
Chemicals
Quantity
H2O2, 0.5000 M
40 mL
HCl, 0.050 M
40 mL
KI, 0.05000 M
40 mL
Na2S2O3, 0.01000 M
50 mL
Starch Solution
5 mL
H2O
As needed
Procedure:
1. Clean and mostly dry four Erlenmeyer flasks.
2. Label them 1A, 2A, 3A, and 4A. This flask will be “Beaker A” for the trial.
3. Use fresh pipettes for each solution. Rinse pipette twice with the solution that you will be measuring and
keep this prepared pipette with the corresponding solution. Add all of the contents for run 1, beaker A to the
flask labeled 1A. The contents of the beaker can be seen in the table below.
Table 2: Beaker Contents for Iodine Clock Reaction
Run
1
2
3
4
KI Volume (mL)
10.00
10.00
10.00
5.00
Beaker A
Na2S2O3
Volume (mL)
10.00
10.00
10.00
10.00
H2O Volume
(mL)
10.00
15.00
15.00
15.00
Beaker B
H2O2 Volume
(mL)
10.00
5.00
10.00
10.00
HCl Volume
(mL)
10.00
10.00
5.00
10.00
4. Rinse and mostly dry 4 beakers and label as 1B, 2B, 3B, and 4B.
5. Using the same technique that was utilized in step 3, prepare the following solutions in clean beakers using
Table 2 above.
6. Place 5 – 10 drops of starch solution in beaker A.
7. Add the contents of beaker 1A to flask 1B. Start the stopwatch as soon as the two solutions meet
8. Swirl the flask to mix and note the time it takes for the color to change. This is Run 1.
9. Add the contents of beaker 2A to flask 2B. Repeat procedure in step 8. This is Run 2.
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Date:
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10. Add the contents of beaker 3A to flask 3B. Repeat procedure in step 8. This is Run 3.
11. Add the contents of beaker 4A to flask 4B. Repeat procedure in step 8. This is Run 4.
NOTE – Make sure that the shade of blue/black color you observe is the same in each of the three runs.
12. If time allows, repeat runs 1 – 4 in order to replicate your data.
Data and Calculations:
Table 3: Results for Iodine Clock Reaction
Time
for
Blue
Color
to
Partial
Order of H2O2, H+
Rate Constant (
)
mol
Run
Rate (
)
Appear (s)
and IL× s
1
2
3
4
1. When you mixed the stock solutions with other chemicals you created a larger overall volume. Thus, you are
diluting the solution. Therefore, you must calculate the new diluted concentration for each solution in each trial
using the equation:
M1V1 = M2V2
Whereby M1 is your initial molarity, V1 is your initial volume, M2 is your molarity after mixing, and V2 is the
total volume of the new solution.
2. The rate of reaction for each of the four runs is determined through the use of the concentration of thiosulfate
and the average time for the blue color to appear in each trial:
[S2O3-2 ]
Rate of Re action =
time for blue to appear
3. We know that the rate law is of the form:
Rate = k[H+]x[I-]y[H2O2]z
Using the data from Runs 1 – 4, determine the values of x, y, and z.
4. Calculate the overall order of the reaction
5. Using the rate law that you determined in calculation 3. Determine the value of k at room temperature (the
temperature at which the experiment was ran.
Discussion Questions
1. From the data and observation made above, what is the molecularity of the rate-limiting step of reaction 1?
2. Would the rate constant be the same if the reaction were run at a higher temperature? Why or why not?
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