CHEMICAL KINETICS

P. F. P. FLORES
National Institute of Physics, College of Science
University of the Philippines, Diliman, Quezon City 1101, Philippines
Date Submitted: 8 February 2019
Date Performed: 30 January 2019

Abstract

Chemical kinetics studies the different factors of rate of chemical reactions. This experiment
used the iodine clock experiment to determine how the different factors such as, concentration,
−¿¿
temperature, and presence of catalyst, affects the rate of the chemical reaction of I and
2−¿ −¿¿ 2−¿
¿ . Five set-ups of solution with and solution with ¿ with varying
S 2 O8 I S 2 O8
concentrations were mixed. The amount of time when the mixed solution changed color was
recorded. Three of one of the set-ups were made again but the temperature was variated instead of
concentration. Another same set-up was demonstrated again, but with the presence of the Cu 2+ ion.
The results showed that with higher temperature, higher concentration, and presence of catalyst, a
faster rate of reaction occurred. The Arrhenius constant, activation energy, and rate constant (k)
were also calculated, which are 59.2 kJ/mol, 1.29 x107 1/sec, and 3.0 x 10-3, respectively. The study
had the possibility to commit several sources of error that specifically affected the concentrations of
the solutions used.

Introduction In this paper, an iodine clock experiment will

be discussed. It usually used in laboratory
Not all chemical reactions take the same time classes to demonstrate chemical kinetics and
to occur completely. Some take a split second the factors that affect it. It shows how long
or even years to complete a chemical reaction. will the thiosulfate ion will be completely
This is due to the concept of chemical depleted in the solution. The experiment
kinetics. There are several factors affecting determined the following: to describe the
the rate of change in concentration of a kinetics of a chemical reaction, to use the
substance, namely: nature of reactants, initial rate method to determine the rate law
concentration of reactants, temperature and of the reaction, to observe the effect of the
addition of catalyst. These factors are temperature and addition of catalyst on the
determined in order to achieve a certain goal. reaction rate and to calculate the Arrhenius
These are determined to be used in different constant. [2]
applications such as storing milk in a
refrigerator to reduce the rate of the chemical The iodine clock experiment demonstrates
reactions involved. In this case, the the coupled reaction between these 2
temperature was adjusted for the milk to chemical equations:
avoid spoilage. [1]
 2−¿ several factors, namely: nature of reaction,
2−¿ → I 2 +2 SO ¿4 presence of catalyst, and temperature. It is
(1) determined by these equations which are also
−¿+ S2 O¿8
known as the integrated rate law for a first
2 I¿ order:
2−¿
−¿+ S2 O ¿8 −¿¿
2−¿ → 2 I ¿ I
(4) ¿
I 2 +2 S 2 O¿3 '
ln rate=ln k +n ln ¿
wherein ln k’ = k[S2O82-], if [S2O82-] is constant
Due to the following chemical equations
mentioned, the resulting equation in
2−¿
determining the rate of the reaction is:
S 2 O¿8
(5)
2−¿ ¿
¿
S2 O 3 ln rate=ln k ' +m ln¿
¿ wherein ln k’ = k[I-], if [I-] is constant
−¿
I¿ The activation energy describes the excess
¿ energies of molecules. It is determined by this
(2) 2−¿ equation:
S2 O¿8
¿ EA
−∆ ¿ (6) ln k =ln A−
RT
−∆ ¿ wherein,
−∆ ¿ R = 8.314 J/mol K
rate=¿ A = Arrhenius constant
EA = Activation energy
The order of reaction describes the
relationship between a concentration of a The Arrhenius constant describes the
solution and its rate of reaction. It is variation of chemical reactions with respect
determined by this equation: to temperature as Svante Arrhenius
demonstrated himself. Consequently, it is
−¿¿¿
n determined by this equation:
I
−E A
(3) 2−¿ ¿m ¿ (7) RT
k =A
S2 O¿8
rate=k ¿ Methodology [2]
wherein:
2−¿ The following stock solutions were prepared:
m = rate order of S 2 O¿8 0.2 M KI, 0.2 M KCl, 0.1 M K 2S2O8, 0.1 M K2SO4,
¿ 4.0 mM Na2S2O3, 1% (w/v) starch solution,
−¿¿ and 0.01 M CuSO4. Five runs were made to
n = rate order of I determine the effect of Persulfate and Iodide
¿ concentration on the reaction rate. Each run
k = rate constant
required 2 sets of solutions, namely A and B.
The contents of A and B for each run are
The rate constant (k) is dependent on the
found in Table 1. The timer was started as
chemical reaction. This value varies due to
soon as solution A was added to solution B, the initial and final concentration of S2O32-.
and then later stopped when the combined Since the reaction demonstrated the
solutions turned into blue. depletion of S2O32-, the final concentration for
it is 0 M. This is determined due to the
Table 1 Composition of Solutions A and B presence of the blue-black color of the
Run Solution A Solution B (+ 3 drops of solution. This
(mL) starch solution) (mL) -11.0
-1.8 -1.6 -1.4 -1.2 -1 -0.8 -0.6 -0.4 -0.2 0
0.2 M 0.2 M 0.1 M 4.0 mM 0.01 -11.5

ln (rate)
KI KCl K2S2O8 Na2S2O3 M -12.0
CuSO4
1 10.0 0 5.0 5.0 5.0 -12.5
2 5.0 5.0 5.0 5.0 5.0 -13.0
3 2.5 7.5 5.0 5.0 5.0
4 5.0 5.0 7.5 5.0 2.5
ln [I-]
5 5.0 5.0 10.0 5.0 0

happened due to the presence of starch,

Two sets of Run 2 (refer to Table 1) were which acted as an indicator. As soon as the
RUN [S2O82-] [I-] [S2O32-] SO42- was completely depleted, the reaction
1 0.02 0.8 0.0008 formed an iodine starch complex.[1}
2 0.02 0.4 0.0008
3 0.02 0.2 0.0008 Table 2 New Concentration Values of the
Solutions in specific runs
4 0.03 0.4 0.0008
5 0.04 0.4 0.0008 Results from runs 2, 4, and 5 (refer to Table
prepared again to determine the effect of 3), where [I-] was constant, were plotted on a
temperature on the reaction rate. The first ln rate vs ln [S2O82-] graph which showed a
set heated the two solutions up to 50°C in a trend (refer to Figure 1). This graph implied
hot bath before mixing, while the second set that the concentration of thiosulfate is
cooled the two solutions up to 5°C in an ice directly proportional to the rate of reaction.
bath before mixing. The timer started when -10.5
the solutions were mixed, and then stopped -4 -3.9 -3.8 -3.7 -3.6 -3.5 -3.4 -3.3 -3.2 -3.1
as the mixed solution changed color. -11.0
ln (rate)

-11.5
Another set of Run 2 (in Table 1) was
prepared again. Before mixing A and B, 4 -12.0
drops of 0.01 M CuSO4 was added. The timer
-12.5
started when the solutions were mixed, and
then stopped as the mixed solution changed ln [S2O32-]
color.
Figure 1 Graph of ln(rate) vs. ln[S2O32-]
Results and Discussion
The graph represents the equation:
The first part on the experiment, the five runs
were differentiated by changing the y=1.49 x−6.50 (eqn. 5)
concentrations of either the thiosulfate ion or with a value of r2 = 0.96.
iodine ion (refer to Table 2). These data were
obtained by using equation (2). The final This states that the order of reaction for
concentrations of S2O82- and I- were not thiosulfate in this experiment is 1.49.
obtained, but the rate of reaction came from
The results from runs 1, 2, and 3 (refer to were altered (refer to Table 4 for results).
Table 3), where the [S2O82-] was constant, Table 4 showed the relationship between the
were also plotted on a ln rate vs ln [I -] graph temperature and the rate of reaction is
which showed an almost linear trend (refer to directly proportional with each other. The
Figure 2). The graph also implied that the results were plotted on a ln k vs 1/T in Figure
concentration of iodine is directly 3.
proportional to the rate of reaction. 0.00
0.003 0.003 0.003 0.003 0.003 0.004 0.004 0.004
Figure 2 Graph of ln(rate) vs. ln [I-] -2.00
-4.00
The graph represents the equation:

ln(k)
-6.00
y=0.74 x −11.6 (eqn. 4) -8.00
with a value of r2 = 0.98.
-10.00

This states that the order of reaction for 1/T (in K)

iodine in this experiment is 0.74
Figure 3 Graph of ln k vs 1/T
Table 3 Computed Reaction Rates for every
run The activation energy was determined by
RUN Rate (M/s) multiplying the slope to the negative gas
constant, derived from equation (6), which is:
1 7.843 x 10-6
2 4.124 x 10-6 EA = 59.2 kJ/mol
3 2.817 x 10-6
4 9.091 x 10-6 The Arrhenius constant is determined by
5 1.143 x 10-5 raising Euler’s number (e) to the intercept
from the equation formed in the ln k vs 1/T
graph, from equation (7). The resulting
The rate constant (k) was calculated from the
Arrhenius constant from this experiment is:
rate law equation (3) which gave the result of
approximately 3.0 x 10-3.
A = 1.29 x107 1/sec
0.74
−¿ ¿¿ Table 4 Effect of the presence of catalyst in the
I reaction to the rate of the reaction
1.49
(8) 2−¿ ¿ ¿
S2 O¿8 SET-UP CATALYST RATE (M/s)
rate=k ¿ 1 None 4.124 x 10-6
4 Cu2+ 9.302 x 10-6
Table 4 Effect of Temperature to the rate of
reaction Table 4 shows that the rate of the reaction
SET-UP Temp (K) Rate (M/s) increased when the ion Cu2+ is present in the
reaction. CuSO4 served as the catalyst in this
1 298 4.124 x 10-6 experiment. According to Silberberg[3], the
2 323 2.857 x 10-5 presence of a catalyst provides an alternative
pathway for the reaction. The pathway of the
3 278 8.097 x 10-7
reaction with the presence of catalyst has a
lower activation energy to make the reaction
The second part of the experiment repeated complete faster.
the 2nd run in the first part but, temperatures
The most probable mechanism for this factors in order to occur in an ample time.
experiment is: Not all chemical reactions occur at a short
time. Some reactions take years to occur or
(aq) even in a split second. However, there are
2−¿
[ I ⋯ S 2 O8 ](aq) factors that may shorten or lengthen the
complete reaction. These factors may be the
−¿ +S 2 O2−¿
¿ slow ¿ temperature, addition of catalysts, or surface
→
¿ area (for solids). It is important to determine
I (aq) specific kinematic parameters of a specific
2−¿
2 S O 4(aq ) chemical reaction in order to achieve an
2−¿
[I ⋯ S2 O 8 ](aq) fast I 2(aq) +¿ enough amount of time.
→
−¿+¿ The iodine clock experiment showed how
I ¿(aq) these factors specifically can affect the rate of
a chemical reaction. It showed that when
From the determined rate law [equation (8)], there is either a higher amount of
it is concluded that the rate of reaction of the concentration of either of the 2 substances, a
first chemical equation from equation (1) is higher temperature, or a presence of a
slower compared to the second chemical catalyst, the rate of reaction would occur
equation in equation (1). faster.

The results from the previous studies are The rate law for this experiment was determined,
quite similar with the results in this which was:
experiment. The variation of the results −¿ ¿¿0.74
between the two experiments can be affected I
many factors. Since concentration is 2−¿ ¿ 1.49 ¿
dependent on the rate of reaction [1], wet S2 O¿8
apparatuses used like beakers may have
rate=k ¿
caused different concentrations than the ones wherein, the calculated rate constant (k)
recorded. Another factor that can affect the is 3.0 x 10-3.
rate is the preparation of stock solutions
might not be precise enough from the target The Arrhenius constant and activation energy
concentration of the solution. These two were also calculated, which are 59.2 kJ/mol
factors may cause the concentration to lower and 1.29 x107 1/sec, respectively.
down which will lead to a slower rate of
reaction. Another factor that can affect is the Due to the possible amounts of error, this
inconsistency of the observer and the timer. study recommends to always use fresh stock
The observer might have stopped the timer solutions, to be consistent with the definition
when the solution showed a hint of blue, or of change in color (may be either pale blue or
very dark blue in any of the trials. This can blue-black), and to be mindful in using
either lower or higher the rate of reaction. laboratory equipment especially when the
result of the experiment is strictly dependent
Conclusion on the concentration of the solutions.
Chemical kinetic experiments are important
since some chemical reactions have different

References [1] Petrucci, Ralph H. General chemistry:

principles and modern applications. Toronto,
Ont.: Pearson Canada, 2011.
[2] General Chemistry II Laboratory Manual;
Institute of Chemistry, University of the
Philippines Diliman: Diliman, Quezon City,
2018.

[3] Silberberg, M.S. Principles of General

Chemistry, 2nd Edition. McGraw-Hill: 2010.