Microscale Chemistry Experiment
Investigation of the overall order of reaction between hydrogen
peroxide and iodide in acidic medium (clock reaction)
Student Handout
Purposes
To determine the kinetic order of the reaction between H2O2(aq) and I-(aq) in acidic
media with respect to
(i)
H2O2(aq),
(ii)
I-(aq) and
(iii)
H+(aq).
Introduction
The kinetics of the reaction:
H2O2(aq) + 2I-(aq) + 2H+(aq)  I2(aq) + 2H2O(l)
can be investigated by the introduction of a small and fixed amount of S2O32-(aq) and
starch indicator.
H2O2(aq) + 2I-(aq) + 2H+(aq)  I2(aq) + 2H2O(l) ……….. main reaction
2S2O32-(aq) + I2(aq)  S4O62-(aq) + 2I-(aq) …….monitor reaction
Starch solution + I2(aq)  blue complex ………… ..indicator reaction
The added S2O32-(aq) consumes the I2(aq) produced from the main reaction. As long as
there are S2O32-(aq) ions in the reaction mixture, I2(aq) formed from the main reaction
will be instantaneously consumed by the S2O32-(aq) ions and the starch indicator will not
be affected. However, when all S2O32-(aq) ions are consumed, the I2(aq) starts to form and
will immediately turn the starch indicator to deep blue.
The overall result is that upon mixing different amounts of H2O2(aq), I-(aq), H+(aq),
S2O32-(aq) and starch indicator, no change will be observed at the start of the experiment,
but the reaction mixture suddenly changes to deep blue after a period of time. The time
elapsed before the development of the blue colour depends on the amount of S2O32-(aq)
used. The greater the amount of S2O32-(aq) is used, the longer will be the time taken for
the development of colour. Thus the reaction responsible by the S2O32-(aq) is also known
as the monitor reaction, as it controls the time taken for the development of colour.
Reactions using the above technique are collectively classified as ‘clock reactions’. If
iodine is used to indicate the reaction time, it is called an iodine clock reaction. Likewise,
if bromine is used, it is called a bromine clock reaction. (The ‘clock’ can be slowed down
by adding more S2O32-(aq).)
Kinetic interpretation of clock reactions
Time elapsed for colour development indicates the time (t) taken for the formation of a
certain amount of iodine from the main reaction. 1/t would be proportional to the rate of
formation of this amount of iodine. 1/t would also be proportional to the initial rate of
decrease in concentration of I-(aq) or H2O2(aq) if the amount of iodine formed is small or
if the amount of S2O32-(aq) used is small. Hence there is a need to use small amount of
1
S2O32-(aq).
Order w.r.t. I-(aq) will be investigated by keeping concentrations of H2O2(aq) and H+(aq)
constant while varying the concentration of I-(aq) in the ratio of 1 : 2 : 4 : 8. If the ratio of
(1/t) doubles each time, the order of reaction w.r.t. to I-(aq) will be determined as 1. If the
ratio of (1/t) remains unchanged, the order can be regarded as zero. The experiment is
then repeated for determining orders for H2O2(aq) and H+(aq).
Safety
Wear safety spectacles and avoid skin contact with the chemicals. Dispose
of chemical waste and excess materials according to your teacher’s
instruction.
EYE PROTECTION
MUST BE WORN
Further information on the chemicals used in the experiment can be found in the
Material Safety Data Sheet (MSDS). Consult your teacher for details.
Materials and Apparatus
About 20 cm3 of each of the following solutions in labelled plastic bottles:
0.60 M H2SO4(aq), 0.60 M KI(aq), 1.50% H2O2(aq)
HARMFUL
HARMFUL/
IRRITANT
0.08 M Na2S2O3(aq), starch solution, deionized water.
Two 8-well reaction strips, micro-tip plastic pipette, stop watch, micro-stirrer or
toothpicks.
Experimental Procedure
(A) Kinetic order w.r.t. iodide ion
1. Using a fresh and clean micro-tip plastic pipette, transfer 1 drop each of 1.5%
H2O2(aq), 0.6 M H2SO4(aq) and starch indicator solution to 4 separate wells of a
8-well reaction strip (call it strip A) so that each well has a total volume of 3 drops
2. Take another 8-well reaction strip (call it strip B), again using a fresh and clean
micro-tip plastic pipette, transfer 1 drop of 0.08 M Na2S2O3(aq) to each of the first 4
wells.
3. Into the same 8-well reaction strip, place 1 drop of 0.6 M KI(aq) to the first well, 2
drops to the second, 4 drops to the third and 8 drops to the fourth. Add 7 drops of
deionized water to the first well, 6 drops to the second and 4 drops to the third so
that the total volume of reactant mixture in each of the 4 well of strip B is 9 drops.
(see Table 1)
4. Stir the solution mixture in each of the wells of strip B with micro-stirrer or
toothpick.
5. Invert strip B and stack it atop strip A so that the first 4 wells of strip B is directly
above the first 4 wells of strip A.
2
6. Hold the two strips firmly together by means of two small pieces of rubber tubing,
one at each end and lower the strip combination suddenly (“shake-down”
technique) so that the two solution mixtures mix thoroughly (see Figure 2). Start the
stop watch at the same time.
7. Turn the strip combination upside down repeatedly and look for the sudden
appearance of a deep blue colour. Record the time taken. Carry on recording time
until all the 4 wells have developed colour in the correct sequence.
8. Clean the reaction strips thoroughly with deionized water and empty the water in
the wells.
Fig. 1 Size of 8-well reaction strip
Fig. 2 The “shake-down” technique
Table (1)
Number of drops
Strip A
Well
1
2
3
4
H2O2(aq) H2SO4(aq)
1
1
Strip B
Starch
solution
I-(aq)
H2O(l)
0.08M
S2O32-(aq)
1
1
2
4
8
7
6
4
0
1
(B) Kinetic order w.r.t. H2O2
9.
Repeat steps (1) to (8) according to Table (2).
Table (2)
Number of drops
Well
1
2
3
4
I-(aq)
1
Strip A
0.08 M
S2O32-(aq)
1
Strip B
Starch
solution
H2O2(aq)
H2O(l)
H2SO4(aq)
1
1
2
4
8
7
6
4
0
1
3
(C) Kinetic order w.r.t. H+
10. Repeat steps (1) to (8) according to Table (3).
Table (3):
Number of drops
Well
1
2
3
4
I-(aq)
1
Strip A
0.08M
S2O32-(aq)
1
Strip B
Starch
H2SO4(aq)
solution
1
2
1
4
8
H2O(l)
H2O2(aq)
7
6
4
0
1
Results
Variable
[I-(aq)]
[H2O2(aq)]
[H+(aq)]
Relative
concentration
t (s)
Rel. initial Deduced
rate,1/t (s-1)
order
1
2
4
8
1
2
4
8
1
2
4
8
(If initial rate doubles when the concentration of a species is doubled,
the reaction is first order with respect to that species)
Discussion Questions
1.
With the help of an appropriate sketch, illustrate the meaning of “initial rate”.
2.
Why is it assumed that in order to obtain initial rate, time (t) has to be small?
3.
Why are initial rate preferred to rates at other times of a reaction, i.e.
instantaneous rates?
4.
Why the amount of S2O32-(aq) added to the reaction mixture has to be small?
5.
From the deduced order of each of the reactants, give a rate equation for the
reaction.
6.
A mechanism for the reaction consists of the following three elementary steps:
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H2O2(aq) + I-(aq)  X + H2O(l) …………. (slow)
H+(aq) + X  Y ……………………. (fast)
Y + H+(aq) + I-(aq)  I2(aq) + H2O(l) ……... (fast)
Suggest species for X and Y in the above elementary reactions in order that the rate
expression for the rate determining step fits with the experimental rate equation.
Reference
D. Ehrenkranz and J. Mauch Chemistry in Microscale, Kendall/Hunt Publishing Co. (1990)
p. 153
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Microscale Chemistry Experiment
Investigation of the overall order of reaction between hydrogen
peroxide and iodide in acidic medium (clock reaction)
Teacher Notes
Kinetic runs are normally performed in sequential order starting with the first one. Time
taken for the whole experiment is the sum of the time spent for each run. One can speed
up the experiment by starting all the kinetic runs at the same time and recording the
time taken for all the kinetic runs simultaneously in one go. Time spent for the whole
experiment is then equal to the time required for the longest run. Thus the experiment
can be completed in a much shorter time. This can be achieved by using micro-scale
equipment together with the “shake-down” technique.
Curriculum Links
*
*
*
*
Rate of chemical reaction
Factors influencing reaction rate
Rate equation and order of reaction
The interpretation of reaction rate at molecular level
Pre-Laboratory Talk
1.
Brief on the purposes, technique and advantages in using the micro-scale
instrument, in particular the “shake-down” technique, perhaps using a
demonstration.
2.
Remind students that there will be ample time in carrying out the experiment and
that there is a need to repeat the experiment for a number of times to obtain
concordant results.
Time Required
Pre-laboratory talk: 15 minutes; Experiment: 40 minutes
Remarks
(i)
Advantages in using the micro-scale technique
Besides the usual advantages enjoyed by performing an experiment in small scale,
the following three advantages cannot possibly be achieved by conventional
methods:
(a)
The first lies in expressing relative concentration in terms of number of drops
of reagent rather than actual concentration in mol dm-3. This greatly simplifies
the calculation.
(b)
The second is that the experiment allows 4 kinetic runs to be performed at the
same time instead of spending time on 4 different experiments. This shortens
the time required for each order determination to as little as 1 minute.
(c)
The layout of the experiment enables students to understand the aim more
coherently as order determinations are completed with minimal steps.
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(ii)
The “1248” method
The method offers a non-graphical approach to study reaction kinetics. Judging
from the name of the method, it literally means varying the concentration of the
reacting species according to a ratio of 1: 2: 4: 8, or doubling the concentration of
the species in succession. If the observed reaction rates also increase in the same
order , i.e. 1: 2: 4: 8, then the kinetic order of the reaction with respect to the
species must be one. On the other hand, if the observed reaction rates increase
according to a ratio of (1)2: (2)2: (4)2: (8)2, then the kinetic order of the reaction with
respect to the species should be two. Lastly, if the reaction rate is insensitive to
change in concentration, the reaction is obviously of zero order with respect to the
species. Reaction examples encountered in AL Chemistry are mainly of first or
zero order kinetics, rarely of second order kinetics.
(iii)
Sources of material
The crucial apparatus of the experiment is the 8-well reaction strip. Strips
employed should be small enough so that contained liquids would not drop when
they are turned upside down. The 8-well reaction strips required for the
experiment are obtainable from: MicroChem Lab (Website: www.mcl.hk)
Sample results
Variable
[I-(aq)]
[H2O2(aq)]
[H+(aq)]
Relative
concentration
t (s)
Rel. initial
rate,1/t (s-1)
1
2
4
8
1
2
4
8
1
2
4
104
55
28
14
115
58
29
14
102
72
51
9.6 x 10-3
1.8 x 10-2
3.6 x 10-2
7.1 x 10-2
8.7 x 10-3
1.7 x 10-2
3.4 x 10-2
7.1 x 10-2
9.8 x 10-3
1.4 x 10-2
2.0 x 10-2
8
26
3.8 x 10-2
order
1
1
---
Since the “shake-down” mixing method is not as effective as conventional pouring
and stirring, repeated “shake-down” is necessary to achieve a close to
homogeneous solution mixture. Thus each run has to be repeated two or three
times to obtain concordant results. For both variations of I-(aq) and H2O2(aq),
doubling the concentration also doubles the reaction rate. The reaction is therefore
first order with respect to I-(aq) and to H2O2(aq). The figures for the H+(aq) ion
variation showed no strong indication for a first order or a zero order kinetics with
respect to H+(aq).
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Suggested Answers to Discussion Questions
1.
With the help of an appropriate sketch, illustrate the meaning of “initial rate”.
Answer:
According to the main reaction,
H2O2(aq) + 2I-(aq) + 2H+(aq)  I2(aq) + 2H2O(l)
formation of iodine is due to the consumption of iodide ions. The formation of a
small amount of iodine is determined by the extent of decrease in concentration of
iodide ions. As shown by the following decay curve, initial rate or the rate of
decrease in concentration of iodide ions at the start of the experiment can be
achieved by considering [I-]o, [I-]t1 and t1 . A longer time interval, t2, gives average
rate.
[iodide] o
[Iodide]
t1
t1
[Iodide]
t2
t2
Time
2.
Why is it assumed that in order to obtain initial rates, t has to be small?
Answer:
From the above sketch, initial rate can be achieved only if a short time interval is
considered. A longer time interval will result in average rates which are not
regarded as instantaneous rates. Rates considered in rate equations are
instantaneous rates.
3.
Why are initial rates preferred to rates at other times of a reaction, i.e.
instantaneous rates?
Answer:
Initial rates rather than instantaneous rates are generally considered. This is because
initial rates can be measured more easily and less reaction complications are
expected at the start of the reaction.
4.
Why the amount of S2O32-(aq) added to the reaction mixture has to be small?
Answer:
The added S2O32-(aq) consumes the iodine formed and controls the initial rate of
decrease of iodide ions of the decay curve. For a decent initial rate of decrease of
iodide ions, amount of S2O32-(aq) used has to be small so that the time taken for the
appearance of the blue colour is short enough to warrant for an initial rate.
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5.
From the deduced order of each of the reactants, give a rate equation for the
reaction.
Answer:
d[I - (aq)]
 k [I - (aq)][ H 2 O 2 (aq)]
dt
6. Answer:
X = OI-(aq),
Y = HOI(aq)
Reference
D. Ehrenkranz and J. Mauch Chemistry in Microscale, Kendall/Hunt Publishing Co. (1990)
p. 153
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