Chemical Kinetics: Iodine Clock Reaction
Chemistry 102
Sept. 28, 2010
Abstract
The purpose of this experiment was to learn about the effects of concentration,
temperature, and catalysts on reaction rate. To do this, a known amount of thiosulfate was added
to consume the I2 created in the slower primary reaction, thereby creating a clock by which to
control the amount of I2 reacting with starch to turn the solution blue. In part one of the
experiment, it was observed that by changing the concentrations of the reactants; I-, BrO3-, and
H+, the solution turned blue at different rates, with an inverse relationship between concentration
and time thereby proving that higher concentrations of reactants decrease the time of reaction. In
part two, it was observed that by fixing the concentrations and altering temperature, the changes
in overall reaction time indicated an inverse proportionality between temperature and time, or in
other words the higher the reaction temperature, the shorter the reaction time. In part three,
ammonium molybdate was added to the reactants and reaction time was decreased by a factor of
10, showing that a catalyst shortens the reaction time and lowers the activation energy. The
individual reaction orders for the reactants were determined to be: I- = 1, BrO3- = 1, and H+ = 2.
The overall reaction’s order was therefore 4, and the average rate constant was 33.87 1/M3s, and
the activation energy was 46.06KJ.
Introduction
Chemical Kinetics is the study of the rates of reactions and how these rates are affected
by varying physical properties like temperature, concentration, and even catalysis. Catalysis is
the introduction of a molecule into a reaction which lowers the total activation energy of a
reaction without changing what is produced by that reaction. Activation energy is the amount of
energy that molecules in a reaction must have in order to collide in the proper manner to create a
product
In this experiment iodide, bromate, and hydrogen ions were reacted in order to observe
the effects of the three variables; concentration of reactants, temperature, and the introduction of
a catalyst. The equation for this reaction is:
6I-(aq) + BrO3-(aq) + 6H+(aq) → 3I2(aq) + Br-(aq) +3H2O(l)
From this equation we derive the rate law for this reaction which in this case is:
rate = k[I]x[BrO3]y[H+]z
In order to determine the rate orders (x,y, and z), five mixtures of the previous reaction with
varying concentrations were compared and algebraically divided. With the rate orders
determined, the rate constant K is then computed.
As previously stated, a small amount of sodium thiosulfate was used to form a clock
reaction and guarantee that a known amount of I2 would be formed and react with the starch.
Because any amount of I2 reacting with starch would turn the solution blue, it was necessary to
delay the reaction by introducing the sodium thiosulfate. Once the thiosulfate was consumed, the
desired amount of I2 had been formed. The equation of this reaction is:
I2(aq) + 2S2O32-(aq) → 2I-(aq) + S4O62-(aq)
By using this equation along with the first equation, it is possible to measure the amount of time
it takes for the solution to turn blue as a function of the three reactants. The resultant equation
from this is:
(3 mol I2 / 1 mol BrO3-) x (2 mol S2O32- / 1 mol I2) therefore
∆[BrO3-] / ∆ t = 1/6 ∆ [S2O32-]/∆t
Part two of the experiment involved reacting the first mixture from part one at four
different temperatures in order to understand the effect of temperature and to determine the
activation energy, Ea, for the reaction. The activation energy is calculated by using the
Arrhenius equation:
ln k = -Ea / R (1/T) + ln A
The value of Ea was calculated by graphing ln k vs. 1/T, where T is the temperature of the
reaction in Kelvins. The slope of the resulting graph is equal to Ea / R, R being the constant
8.314 J/mol K.
Finally for part three, the catalyst ammonium molybdate is added to the first mixture of
part one and reacted. This should result in a decrease in the time of reaction.
Experimental
First the reaction order must be determined so an experiment was performed solely for
that purpose. In this section, the contents of two flasks were mixed together and swirled until the
solution turned blue. The concentration in each flask varied according to the trial, and the
reaction times and temperatures were documented after each of the 5 trials. Listed below are the
concentration variations:
Flask 1
Flask 1
Flask 1 Flask 2
Flask 2
Flask 2
Mix .010 M KI .00010 M Na2S2O3 H2O
.040 M KBrO3 .10 M HCL Starch
1
10 ml
10 ml
10 ml
10 ml
10 ml
3 drops
2
20 ml
10 ml
0 ml
10 ml
10 ml
3 drops
3
10 ml
10 ml
0 ml
20 ml
10 ml
3 drops
4
10 ml
10 ml
0 ml
10 ml
20 ml
3 drops
5
8 ml
10 ml
12 ml
5 ml
15 ml
3 drops
In a second experiment, the effect of temperature on the reactions rates was observed and
studied. In this section, reaction mixture 1 was used in all trials with varying temperatures. The
temperatures used were approximately 2o, 8o, 22.4o, and 44o Celsius. The flasks were either
heated in a hot water bath or cooled in ice prior to being added together. The time that it took for
the solution to turn blue was then documented.
Finally, the effect of a catalyst was studied by the addition of one drop of 0.5 M
(NH4)2MoO4, ammonium molybdate, to flask 2 of mixture 1 before the solutions were mixed.
The time for this new reaction was compared to the time documented for mixture 1 in the first
experiment, and as predicted previously, the catalyst greatly reduced the time of reaction and the
activation energy.
Results and Discussion
In chemical kinetics, higher reactant concentrations mean that there is a higher
probability that molecules will collide properly, causing a reaction to take less time to occur. As
shown in table 1, concentration has a direct effect on the rate of the reaction.
Table 1
Mix
Time (s)
Rate -∆[BRO3-]/∆t
[I-]
[BrO3-]
[H+]
Temp (c)
=3.33 x 10-5 (mol/L)/∆t
1
148.56
2.24 x 10 -7
.0020
.008
.02
22.4
2
76.15
4.37 x 10 -7
.0040
.008
.02
22.2
3
85.34
3.90 x 10 -7
.0020
.016
.02
22.0
4
36.28
9.18 x 10 -7
.0020
.008
.04
22.4
5
181.13
1.84 x 10 -7
.0016
.004
.03
22.0
From the table it is clear that, when the Iodide ion concentration doubles in mix 2, the reaction
time is cut in half from mix 1. The same happens in mix 3 compared to mix 1 when the bromate
ion concentration doubles, indicating that these are first order reactions. In mix 4, the hydrogen
ion concentration is then increased thereby greatly reducing the reaction time, indicating a higher
order reaction.
Table 2 shows the effects of temperature on the time of reaction:
Table 2
Mix
Time (s)
Rate -∆[BRO3-]/∆t
Temp (c)
Calculated k
=3.33 x 10-5 (mol/L)/∆t
1A
148.56
2.24 x 10 -7
22.4
35 1/M3s
1B
31.47
1.06 x 10 -6
44
165.63 1/M3s
1C
449.87
7.40 x 10 -8
2
11.56 1/M3s
1D
328.41
1.01 x 10 -7
8
15.78 1/M3s
By graphing the natural log of k vs. 1/T we are able to calculate a slope which we said earlier is
the negative activation energy divided by the constant R. In this case the slope of the graph is
-5,539.80 with activation energy of 46.06 kJ.
Below, is the graph from the gathered data:
Natural Log of K vs.Temperature-1
5.5
5
LN[K] (M-3s-1)
4.5
y = -5,539.80x + 22.48
4
3.5
3
2.5
2
0.0031
0.0032
0.0033
0.0034
0.0035
0.0036
0.0037
Temperature-1 (Kelvins)
In third experiment performed, the effect of the catalyst ammonium molybdate on the
reaction was observed to increase the rate of reaction by a factor of 10. Without the catalyst, the
reaction took 148.56 seconds to react while the reaction with the catalyst took only 15.40
seconds. From this drastic drop in the time of reaction we can conclude that there must be a
drastic drop in activation energy.
Conclusion
In this experiment, it was observed that the concentrations of reactants, temperature, and
the introduction of a catalyst all play a large role in chemical kinetics. By varying any of these
factors, we were able to greatly change the overall rate of the reaction and therefore change the
total time to react. By analyzing the data, it becomes clear that an increase in the reaction rate is
a result of higher reactant concentrations, an increase in temperature, or by adding a catalyst. In
this experiment, the individual rate orders for the reactants were; I- = 1, BrO3- = 1, and H+ = 2.
The overall reaction order was determined to be 4, the rate constant was 33.87 1/M3s, and the
activation energy was 46.06 KJ.
There were several sources of error in this experiment that could cause the calculated
activation energy to be different than expected. For one, we assume that temperature for each
reaction in part II is constant. In reality this temperature slowly increases throughout the reaction,
requiring more difficult calculus to properly analyze the effect temperature has on the reaction.
Second, there seemed to be problems with the color change if the starch indicator was not just
right. This leads me to believe that because starch plays such an important role in the indication
of our reaction, that any minor changes to the starch could likely result in changes to the time of
reaction. And lastly is the assumption that all values on the reagent bottles are the actual
molarity values for the solutions inside. If these values are off then our initial concentrations
used to find x, y, and z could lead to inaccurate reaction orders which would throw off the rate
constant and ultimately the activation energy.