Kinetics Rate

Topics: Chemical reaction, Chemical kinetics, Chemistry Pages: 7 (1603 words) Published: June 19, 2013
CHEM 1112
Kinetics of the Persulfate – iodide Clock Reaction

The purpose of this experiment is to determine the rate law and the activation energy for the reaction between persulfate ion, S2O82-, and iodide ion, I-:
S2O82-(aq) + 2 I-(aq)  2 SO42-(aq) + I2(aq)

The rate law can be written as
Reaction rate = (1)

Where m and n are the orders with respect to S2O82- and I-, respectively, and k is the rate constant. Determining the rate law involves determining the values of m and n.
The temperature dependence of the rate constant is given by


Equation (2) is called the Arrhenius Equation, where A is the pre-exponential factor, E is the activation energy with units of J/mol, T is the absolute temperature, and R is the gas constant (R= 8.3145 J K-1 mol-1). Equation (2) can also be written in the form:


According to equation (3) the activation energy can be obtained by measuring the rate constant at several temperatures, and then plotting ln k versus 1/T.

Reaction (1) will be carried out in the presence of thiosulfate ion, S2O32- and starch. The concentration of thiosulfate ion will be maintained at a much lower value than that of either the persulfate ion or the iodide ion. The reactions that occur are:

S2O82-(aq) + 2 I-(aq)  2 SO42-(aq) + I2(aq)
Slow (4)
I2¬(aq) + 2 S2O32-(aq)  2 I- + S4O62-(aq)Fast (5) I2¬(aq) + starch  blue complexSlow (6)

Reaction (4) is much slower than reaction (5), and, as a result, the I2 in formed in reaction (4) is immediately consumed by reaction (5), and the concentration of I2 remains at a very low value as long as thiosulfate ion is present. Once all of the thiosulfate ion is used up, the concentration of I2 from reaction (4) increases. The presence of I2 is detected by the formation of a blue complex resulting from the reaction of I2 with starch, reaction (6). The rate of reaction (6) does not become significant until the concentration of I2 becomes appreciable.

A characteristic of this reaction is that the reaction mixture remains colorless for several minutes after the reactants are mixed. During this time both reactions (4) and (5) are occurring. The solution remains colorless because the I2 from reaction (4) is being consumed by reaction (5), and can’t react with the starch. As soon as the thiosulfate ion is used up, the I2 reacts with the starch, and an abrupt and dramatic color change, from colorless to blue, occurs.

The rate of reaction (4) is the rate of consumption of S2O82- ion. Notice that for every S2O82- ion used up two S2O32- ions are consumed. Thus, we have:
Rate of reaction (4) = (7)

If (S2O32-)o is the initial concentration of S2O32 ion, and t is the time interval from the start of the reaction until the solution changes color, then in equation (7), (S2O32-) = 0 - (S2O32-)o , and if the initial concentration of S2O32- is constant then equation (7) becomes

Rate of reaction (4) = (8)

Combining equations (1) and (8) and taking the logarithms of both sides gives:


From equation (9) we see that if we carry out two or more experiments with a constant concentration of I- we obtain

And a graph of versus will give a straight line with slope = m. Similarly, for two or more experiments with constant concentrations of S2O82-, a graph of versus ln(I-) will give a straight line with slope = n.
To obtain the activation energy, we combine equations (3) and (9):

If the concentrations of reactants are constant then equation (11) becomes:

A graph of versus will give a straight line with slope = -


The following solutions will be provided by the Stockroom:

A0.100 M (NH4)2S2O8
B0.100 M (NH4)2SO4
C0.010 M Na2S2O3
D0.100 M KI in 0.066 M (NH4)2SO4
1 % starch solution...
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