Energy and Equilibrium in Chemical Systems
Iodine Clock Kinetics Lab
Introduction
Chemical kinetics is the study of the rates of reactions. In this experiment, we will investigate
the effect of changing certain conditions of a reaction on the rate of the reaction. The reaction we are
using is called the Harcourt-Esson reaction, or more commonly the Iodine Clock reaction. The reaction
proceeds in two steps:
Step I (Production of triiodide ion)
H2O2 (aq) + 3 I- (aq) + 2 H+ (aq)  I3- (aq) + 2 H2O (l)
Step II (Restoration of iodide ion)
I3- (aq) + 2 S2O3-2 (aq)  3 I- (aq) + S4O6-2 (aq)
We can visualize the progress of the reaction by introducing an indicator, starch, which turns a
deep blue or purple in the presence of the triiodide ion. Therefore, the reaction starts clear (thiosulfate
is reacting with triiodide to reduce it to iodine) and become purple (thiosulfate has been exhausted). We
will investigate the effects of several changes in solution composition as well as the addition of energy
as heat on this reaction.
Purpose
Students will design and execute experiments to determine what factors affect the rate of the
iodine clock reaction.
Materials
0.05 M KI
1% starch
0.01 M Na2S2O3
1.0 M H2SO4
0.05 M H2O2
Stopwatch
Hot plate
Ice bath
100 mL (or larger) Erlenmeyer flask
Procedure
“Base” reaction
15 mL KI solution
5 mL starch solution
2.5 mL sodium thiosulfate solution
5 mL sulfuric acid solution
15 mL distilled water
15 mL hydrogen peroxide
Total solution volume: 57.5 mL
Part I
Using the materials of the lab, write your own procedure to determine what affects the rate of
reaction, and to quantify the affect you observe. Focus on the concepts underlying a good
scientific experiment – repeatability, thoroughness, control of variables. Check your procedure
with the teacher before performing it.
Add the substances in the order listed to a clean Erlenmeyer flask. Begin timing the reaction
once the hydrogen peroxide first touches the solution. Stop timing the reaction when a visible
color change occurs. Repeat three times to establish a good baseline.
Part II
Perform the procedure you wrote during Part I. Gather data to determine how the reaction rate
is changed over time.
Lab Analysis Questions
1. Calculate the concentration of each substance in the “base” reaction. Starch concentration
may be calculated as a mass percent.
2. Quantify the effect each change you made had on the rate. For example, “When the
concentration of sodium thiosulfate was increased by a factor of 2, the rate decreased by a
factor of 4.”
3. Estimate the order of reaction with respect to each species that had an effect on the
reaction rate and propose a rate law for the reaction.
4. A second proposed mechanism for the reaction is shown below. Which mechanism, the one
below or the one at the beginning of this document, is most likely to be correct according to
your data? Defend your answer in the context of your data and the principles of chemical
kinetics.
Step 1: H2O2 + I-  H2O + IOStep 2: H+ + IO-  HIO
Step 3: H+ + I- + HIO  I2 + H2O
(slow)
Teacher’s Notes
-
One key concept here that keen students may pick out is that the thiosulfate ion is the “fuse” for
this reaction. The blue color cannot appear until the thiosulfate ion is used up. This concept is
critical, and it helps to take steps to ensure each student understands it. In fact, it would be
quite interesting to gently nudge one group into performing the reaction without the
thiosulfate.
-
So long as the thiosulfate ion concentration is kept constant, the kinetics of the first reaction can
be investigated. Changing the thiosulfate ion concentration changes the “delay”, which may
seem to be affecting the first reaction but in reality is not.