Chemistry
Investigatory
Project
Clock Reaction
By, Aairah Nihal Unni, 12B
1
Acknowledgement
I would like to extend my heartfelt and sincere gratitude
towards my chemistry teacher, Dr. Damandeep Kaur, lab
attendant Mrs.Yashoda and my dear lab partners for
their immeasurable guidance and cooperation. Without
their encouragement and help, I would not have been able
to finish this project on time.
I would also like to thank my school for giving me this
opportunity to experiment and use their lab equipment to
finish this project and also incite my curiosity
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Index
S.No. Contents Pg. No.
01 Acknowledgement 03
02 Introduction 04
03 Theory 05 - 10
04 Materials required 11
05 Procedure 12 - 13
06 Observations 14 - 18
07 Graphs 19 - 20
08 Pictures 21
09 Conclusions 22
10 Precautions 23
3
Introduction
The “clock reaction” is a reaction famous for its dramatic
colorless-to-blue color change, and is often used in
chemistry courses to explore the rate at which reactions
take place.
The color change occurs when I2 reacts with starch to
form a dark blue iodine/starch complex. The ability to
record the time at which the blue complex appears allows
the rate of reaction to be determined accurately with a
stopwatch. In this experiment, the rate law for a reaction
is determined using the method of initial rates.
The effect of concentration on the rate of this reaction is
determined by measuring the initial reaction rate at
several reactant concentrations.
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Theory
The Clock Reaction
This method starts with a solution of hydrogen
peroxide and sulfuric acid. To this a solution
containing potassium iodide, sodium thiosulfate,
and starch is added. There are two reactions
occurring simultaneously in the solution.
In the first, slow reaction, iodine is produced:
H2O2(aq) + 2I-(aq) + 2H+(aq )→ I2(aq) + 2H2O(l)
In the second, fast reaction, iodine is reconverted to
two iodide ions by the thiosulfate :
2S2O32- (aq) + I2 (aq)→ S4O62- (aq) + 2I- (aq)
5
After some time the solution changes color to a very
dark blue, almost black.
When the solutions are mixed, the second reaction
causes the iodine to be consumed much faster than it
is generated, and only a small amount of iodine is
present in the dynamic equilibrium. Once the
thiosulfate ion has been exhausted, this reaction
stops and the blue color caused by the iodine–starch
complex appears.
Anything that accelerates the first reaction will
shorten the time until the solution changes color.
Decreasing the pH (increasing H+ concentration), or
increasing the concentration of iodide or hydrogen
peroxide will shorten the time. Adding more
thiosulfate will have the opposite effect; it will take
longer for the blue color to appear.
6
The primary reaction to be studied is the oxidation
of I- by S2O32- (persulfate) in aqueous solution:
2I- (aq) + S2O82-(aq) → I2(aq) + 2SO4 2-(aq)
(slow, rate determining) Equation 1
This reaction will be run in the presence of a known
amount of S2O32- (thiosulfate), which reacts very
rapidly with I2. As long as S2O32- is present, I2 is
consumed by S2O32- as fast as it is formed. This
competing reaction prevents the I2 produced from
our reaction of interest from reacting with starch, so
no color change is observed until the thiosulfate is
completely used up.
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The "clock" reaction is the reaction of a very small
amount of S2O32- (thiosulfate) with the I2 produced
in the primary reaction:
I2(aq) + 2S2O32-(aq) → 2I- (aq) + S4O62-(aq) (fast)
Equation 2
The “clock” reaction will signal when the primary
reaction forms a specific amount of I2. The amount
of I2 formed before the color change can be
calculated from the known amount of S2O32- added
using the molar ratio in Equation 2. To find the rate
of Equation 1, the change in the concentration of I2 is
monitored over time. Below, [I2] is the change in the
concentration of I2, and t represents the change in
time:
∆[I2]
Rate = /∆t
8
As soon as all of the S2O32- ions have reacted, the I2 still
being formed (Equation 1) starts to accumulate and reacts
with starch. Starch serves as an indicator to help us “see”
the I2, since the interaction between starch and I2 forms a
blue starch-iodine complex. Thus, "∆t" is simply the time
elapsed between mixing the reagents and the appearance
of the blue color. Because the S2O32- ion concentration in
the reaction mixture is known, you can calculate "∆[I2]"
using the stoichiometry of the “clock” reaction. Since the
same amount of S2O32- is added to each run, ∆[I2] is also
the same for each run. However, the amount of time for
the appearance of the blue color varies with initial
reactant concentrations, with temperature, and in the
presence of a catalyst, so ∆t is not constant.
Reaction Rate:-
9
The rate law for Equation 1 will be determined by
measuring the initial rate of reaction with varying initial
reactant concentrations. The concentration of S2O32- in the
reaction mixture is very small compared to the other
reactants present, such that the measured rate is the initial
rate of the reaction. The rate law for Equation 1 is given
below:
Rate = k[I-]x [S2O82-]y
Here, k is the rate constant, and x and y represent the
order of I- and S2O82-, respectively.
Materials Required
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● H 2O
● H2SO4
● KI
● Na2SO4
● H 2O 2
● Starch Solution
● 250ml Beakers
● Measuring cylinder
● Dropper
● Bunsen burner
● Stopwatch
Procedure
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● Take four 250mL beakers and label them as A,B,C
and D
● Add water ( H O), potassium iodide (KI) and sulphuric
2
acid ( H 2
❑
S O 4 ¿ ¿❑ into the beakers as per the following
measurements shown in the table below
Reactant A (mL) B (mL) C (mL) D (mL)
H2O 40 20 20 40
H2SO4 10 10 10 10
KI 20 40 20 20
Na2SO3 10 10 10 10
H2O2 20 20 40 20
● Add 10 drops of starch indicator to each of the
beakers.
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● Simultaneously pour unique concentrations of
Na2SO3 and H2O2 into each beaker X and start the
stopwatch as soon as the solution starts turning blue-
black
● Turn the stopwatch off once the solution turns
completely opaque and note down the your
observations
Observations
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Beaker Timing (Min:sec)
A 13:45
B 21:30
C 15:40
D 13:30
We observe that we can convert the solution colorless to
blue faster by adding less water.
If you increase the concentration of one of the reactants
(like potassium iodide or hydrogen peroxide), the time
taken for the color change to occur usually decreases,
indicating a faster reaction.
Rate Law
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1. Convert Timing to Rate:
● The rate of reaction is inversely proportional to the
time taken for the reaction to complete.
● Calculate the rate for each beaker as
1
Rate = /Time
Times converted to seconds:
● Beaker A: 13:45 = 825 seconds
● Beaker B: 21:30 = 1290 seconds
● Beaker C: 15:40 = 940 seconds
● Beaker D: 13:30 = 810 seconds
Now, calculate the rates:
Rate A= 1/825 = 0.00121s-1
Rate B= 1/1290= 0.000775s-1
Rate C= 1/940 = 0.00106s-1
Rate D= 1/810 = 0.00123s-1
2. Compare the Conditions to Determine Orders x and y:
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● The rate law is Rate = k[I-]x [S2O82-]y
● We need to see how changing the concentrations of I-
and S2O82- affects the rate
Compare Beaker A and B:
● Beaker A: [I-] = 20 mL, [S2O82-] = 10 mL
● Beaker B: [I-] = 40 mL, [S2O82-] = 10 mL
Here, the concentration of I- is doubled, and the rate
changes from 0.00121 to 0.000775.
Rate A
/Rate B = 0.00121/0.000775 ≈1.56
Since the rate roughly increases by 1.56 when the
concentration of I- is doubled, this suggests that x is close
to 1 (first order in I-).
Compare Beaker A and C:
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● Beaker A: [I-] = 20 mL, [S2O82-] = 10 mL
● Beaker C: [I-] = 20 mL, [S2O82-] = 40 mL
The concentration of S2O82- is quadrupled, and the rate
changes from 0.00121 to 0.00106.
/Rate D = 0.00121/0.00106 ≈1.14
Rate A
The small increase in rate suggests that the reaction is
likely less than first order in S2O82-.
Compare Beaker A and D:
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● Beaker A: [I-] = 20 mL, [S2O82-] = 10 mL
● Beaker D: [I-] = 40 mL, [S2O82-] = 20 mL
Both concentrations of I- and S2O82- are doubled, and the
rate increases slightly to 0.00123.
/Rate D = 0.00121/0.00123 ≈ 0.983
Rate A
Graphs
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Pictures
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Conclusion
Based on the observations:
● The reaction is likely first-order with respect to I-
(x≈1).
● The order with respect to S2O82- is close to zero or
very small, given the minimal impact on the rate.
The rate law could be approximated as:
Rate ≈ k[I-] [S2O82-]y
where y is very small, likely close to zero.
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Precautions
● Always wear safety goggles, a lab coat, and gloves to
protect yourself from chemical splashes.
● Sulfuric acid is highly corrosive, and hydrogen
peroxide can be a strong oxidizer. Handle these
chemicals with care, avoiding direct contact with skin
or eyes.
● Use precise measuring tools like pipettes or burettes
to ensure the correct volumes of reactants are used.
Inaccurate measurements can lead to incorrect
conclusions about the rate law.
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● Use a stopwatch or timer with good precision to
record the time it takes for the reaction to occur. The
timing should start as soon as the reactants are mixed.
● After completing the experiment, dispose of all the
chemicals safely and clean all equipment thoroughly.
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