Chemical Kinetics: The Iodine-Clock Reaction:
S2O82−
(aq) + 2 I−
(aq) → I2(aq) + 2 SO42−
(aq)
To measure the rate of this reaction we must measure the rate of concentration change of one of
the reactants or products. Here, it is convenient to carry out a clock reaction involving the
product I2. To do this, you will include (to the reacting S2O82− and I−)
i) a small (but accurately known) amount of sodium thiosulfate, Na2S2O3, and
ii) some starch indicator.
The added Na2S2O3 does not interfere with the rate of above reaction, but it does consume the I2
as soon as it is formed (see below):
2 S2O32−
(aq) + I2(aq) → S4O62−
(aq) + 2 I−
(aq)
This reaction is much faster than the previous, so the conversion of I2 back to I− is
essentially instantaneous.
In a typical experiment, known concentrations of S2O8
2- and I− are mixed with a
small amount of S2O3
2- and starch. As reaction the first proceeds, S2O8
2- and I− react to
produce I2, but nothing appears to be happening because the second reaction is
consuming the I2 as quickly as it is formed.
The important moment is when the S2O3
2- runs out, for then the I2 is no longer
consumed (see the second reaction), and the dark blue starch−I2 complex forms. The
stoichiometry of the second reaction indicates that one I2 molecule must have been
generated for every 2 ions of S2O3
2 − initially present in the solution.
[I2 ] generated(from first reaction) = [I2 ] consumed(in second reaction) = {1/2[S2O32-]
initially present}
Note that the change in concentration of I2 (before the blue complex forms) is fixed by
amount of S2O32- included in the initial reaction mixture.
rate dI
dt
SO
t
==
−
[] []
223
2
1
2
Where t = time for blue color to appear.
In this experiment you will:
A. Determine Rate Law and Rate Constant
{k is constant at constant temperature}
B. Vary temperature and estimate EA
{Using the Arrhenius Equation}
A. Determining the rate constant (k) and the rate law for the iodine-clock reaction:
