CHEMISTRY EXPERIMENT (MOD.S05)
SPOKANE FALLS COMMUNITY COLLEGE
RATES OF CHEMICAL REACTIONS
A CLOCK REACTION
DISCUSSION
This experiment involves the study of the rate properties, or chemical kinetics, of the following reaction
between iodide ion and peroxydisulfate ion:
2I- + S2O8-2  I2 + 2SO4-2
(1)
This reaction proceeds reasonably slowly at room temperature, its rate depending on the concentrations of
the I- and S2O8-2 ions according to the rate law.
For this reaction the rate law takes the form:
rate = k[I-]n[S2O8-2]m
(2)
One of the main purposes of the experiment will be to evaluate the reaction orders n and m for this
reaction and perhaps the rate constant k. We will briefly examine the effect of a catalyst on the rate of the
reaction. We will also investigate the manner in which the reaction rate depends on temperature and may
evaluate the activation energy Ea for the reaction.
OUR METHOD for measuring the rate of the reaction involves what is frequently called a "clock"
reaction. In addition to reaction (1), whose kinetics we will study, the following reaction will also be
made to occur simultaneously in the reaction flask:
I2 + 2S2O3-2  2I- + S4O6-2
(3)
As compared with (1) this reaction is essentially instantaneous. The I2 produced in (1) reacts very quickly
and completely with the thiosulfate (S2O3-2) ion present in the solution (3), so that until all the thiosulfate
ion has reacted, the concentration of I2 is effectively zero. As soon as the S2O3-2 ions are used up from the
system, the I2 then produced by (1) remains in the solution and its concentration begins to increase. When
the I2 concentration builds up every so slightly, its presence is then quickly and strikingly confirmed if a
starch indicator has also been added to the reaction mixture, since I2 in small concentrations reacts with
starch solution to produce a deep, deep blue color.
By carrying out reaction (1) in the presence of S2O3-2 ions and a starch indicator, we introduce a built-in
"clock" into the system. Our clock tells us when sufficient I2 has been produced by reaction (1) to use up
all the S2O3-2 ions originally added. Since, however, one mole of I2 is produced for each mole of S2O8-2
reacted, and each mole of I2 that reacts in (3) with two moles of S2O3-2 ions, the color change also occurs
at the time when a certain amount of I- and of S2O8-2 ions have reacted, namely an amount of I- in moles
equal to the number of moles of S2O3-2 initially present in the reaction flask. If we fix the amount of ions
used at a value that is small compared to the amount of I- and S2O8-2 in solution, the color change will
occur before any appreciable amounts of reactants are used up, and the concentrations of reactants and the
rate is (2) will remain essentially constant during the time interval over which the rate is measured.
Page 1
Amount of I- reacted in Reaction (1) in each case will be = Amount of S2O3-2 reacted in Reaction (3)
2
0.050 moles S2 O3x (0.50 mL) =_________ moles S2O3-2 in (3) = ________moles of I- reacted in (1)
1000 mL
_______moles I- reacted  ____Liter reaction mixture = _____M I- reacted = - [I-1] = - [
Average Rate of Reaction = -
]
Molarity I 1 reacted
[ I 1]
[ _____ ]
  ____________ M/s
Time reacted
t
__ s
The reaction between I- and S2O8-2 ions will be conducted under the conditions in the discussion above.
Carefully measured amount of I- and S2O8-2 ions in water solution will be mixed in the presence of a
relative small amount of S2O3-2 ions (0.50 mL of 0. 100 M S2O3-2)1 and a starch indicator. The time it
takes for the solution to turn blue will be measured for several different solutions in which the amounts
and hence concentration of I- and S2O8-2 ions are varied, but in which the amount of S2O3-2 ion is held
constant. Essentially, what we will measure is the time required for the concentration of the S2O3-2 ion to
decrease by a predetermined amount. Since the rate of a reaction is equal to minus the change (small) in
concentration of a reactant divided by the time required for the change to occur, the experimental data will
furnish the information needed to find the rate of reaction in each solution. Calculations of both the rate
constant k and the orders of the reaction with respect to I- and S2O8-2 follow from the dependence of the
reaction rate on reactant concentrations. Calculation of the activation energy Ea for the reaction may be
made from data obtained on the dependence of the reaction rate on temperature.
EXPERIMENTATION
A.
Dependence of Reaction Rate on Concentration
You will make up reaction mixtures by mixing different volumes of reactants, and record the time it took
to complete the clock reaction.
Obtain six new plastic disposable graduated pipets (1- or 3-mL size), designate and label a specific pipet
for each of the six solutions used in this experiment. Keep each pipet for later use. Do not dispose of any
pipet until you have finished with all the reactions.
Since the success of this experiment depends largely upon measuring accurate volumes of each solution,
you must take care of using these graduated pipets properly with consistency. A common mistake is
drawing up air bubbles and a failure to read liquid volume at eye-level with the “bottom of the meniscus.”
To avoid these problems, always practice the technique described as follows.
1. Note the mark on the pipet stem corresponding to the volume you need, say for example, 2.0 mL.
Have the test tube into which you would deliver the solution ready at hand.
2. Squeeze the pipet bulb before inserting its tip into the solution. Let go of the bulb and watch liquid
level rise smoothly up the pipet to above the 2.0 mL mark.
3. Now line up your eyes with the 2.0 mL mark on the pipet stem. Gently squeeze the bulb to lower
the liquid till the bottom of the meniscus is exactly on the 2.0 mL mark. Gently squeeze-andrelease the bulb to adjust liquid level, if needed.
4. Transfer the solution directly into the test tube. If a drop of solution is “hanging” from the tip of
the pipet, touch the tip of your pipet to the side of the container to remove any droplets.
5. Squeeze and completely empty the entire content of the pipet. Now you have transferred exactly
2.0 mL of the solution into the test tube.
Page 2
The table below summarizes the volumes of reactants needed to prepare three different reaction mixtures.
Procedure for carrying out each reaction will be much the same, as we describe it for Reaction 1.
Table of Reaction Mixtures at Room Temperature
Rx
#
1
Tube I
2.00 mL 0.200 M KI
Tube II
2.00 mL 0.200 M (NH4)2S2O8
Reaction Tube
0.50 mL 0.050 M Na2S2O3
Total mL
5.00 mL
0.50 mL starch
2
1.00 mL 0.200 M KI
2.00 mL 0.200 M (NH4)2S2O8
1.00 mL 0.200 M KCl
3
2.00 mL 0.200 M KI
0.50 mL 0.050 M Na2S2O3
5.00 mL
0.50 mL starch
1.00 mL 0.200 M (NH4)2S2O8
0.50 mL 0.050 M Na2S2O3
1.00 mL 0.200 M (NH4)2SO4
0.50 mL starch
5.00 mL
Use your own designated pipet for each solution. Carefully deliver exactly 2.0 mL 0.200 M KI into a
small test tube, which we will call “Tube 1”. Similarly, deliver 2.0 mL 0.200 M (NH4)2S2O8 into another
small test tube, called “Tube II” in the table.
Deliver 0.50 mL of 0.100 M Na2S2O3 into a large test tube. This will be called the “Reaction Tube”. Add
to this 0.50 mL of starch solution. Insert a thermometer into this Reaction Tube. Record the temperature
for Trail #1 only. Use the “stopwatch” available with the LabQuest.
Remove the thermometer. Note the time and pour the contents of Tubes I and II simultaneously into it.
Swirl and mix the content of the Reaction Tube thoroughly. Stop swirling and place reaction tube in test
tube rack. Watch the time and observe the reaction mixture until it suddenly turns deep blue, hopefully
between 30 - 150 seconds (temperature dependent). Record the time at the instant the deep blue color
appears. Also record the average temperature of the solution to ± 0.2C.
Dispose of the blue reaction mixture in the appropriate waste container. Rinse all of the used reaction
tubes carefully with distilled water (at the same temperature as the reaction solutions) after each run and
drained before using them again.
Repeat the same procedure and do at least another run of Reaction #1. Record the time and temperature
for the two runs. The times should be within 5.00 seconds of each other.
Repeat the experiment with the other mixtures in the previous table. Note that each reaction mixture
contains exactly the same amount of 0.100 M Na2S2O3 (why?) and starch solutions. The KCl and
(NH4)2SO4 solutions are used rather than water in diluting the reaction mixture so that the ionic strength
of the mixture, which has some effect on reaction rate, can be kept essentially constant.
B.
Dependence on Reaction Rate on the Presence of a Catalyst
Reaction 4. Metallic cations have a pronounced catalytic effect on the rate of this reaction. Observe this
effect by repeating Reaction 1 (don't forget the 0.50 mL 0.100 M Na2S2O3 and the 0.50 mL of starch
solution) with a catalyst at room temperature. Before mixing the three solutions, add one drop of 0.100 M
CuSO4 to the flask containing the 0.200 M (NH4)2S2O8. Swirl the flask for a minute to mix the catalyst
thoroughly before adding to the reaction TUBE. Then mix the three solutions, and when the color change
occurs, record the time and average temperature.
Page 3
C.
Dependence of Reaction Rate on Temperature
In this part of the experiment the reaction will be carried out at several different temperatures, but with the
same concentrations of all reactants at each temperature. The Table indicates the reaction conditions.
Table of Reaction Mixtures at Different Temperatures
Reaction
Reaction Mixture
Desired Temperature
1
as in Reaction 1
Around 20C
5
as in Reaction 1
Around 40C
6
as in Reaction 1
Around 10C
The rate of reaction at about 20C can be obtained from the data already collected on Reaction 1.
Reaction 5 is carried out by preparing the same solutions in the same volumes as those used in Reaction 1:
2.00 mL 0.20 M KI into Tube I, 2.00 mL of 0.2 M (NH4)2S2O8 into Tube II and 0.50 mL 0.100 M
Na2S2O3 with 0.50 mL of starch indicator into the Reaction Tube.
Instead of mixing the solutions at room temperature, put all three test tubes in water at about 40C in one
large beaker. Make sure that the water in the large beaker is kept at about 40C. Leave the tubes in the
water for five to ten minutes so that they will be at the proper temperature. Put the thermometer in the
Reaction Tube. Note the time and quickly remove Tubes I and II from the large beaker and pour their
entire contents together into the Reaction Tube. Swirl to mix well and note the initial temperature of the
reaction mixture. Record the time at which the color change occurs and the temperature of the mixture at
that point.
Repeat the same procedure and do at least another run of Reaction #5. Record the time and temperature
for the two runs.
Following similar procedures and complete at least two runs of Reaction #6 at around 10C, cooling all
reactants in ice-water bath to that temperature before starting the reaction. Record the time required for
the reaction and the initial and final temperatures of the reaction mixture.
**** Be sure all times are within 5.00 seconds of each other for each trial
Page 4
NAME
CLASS TIME ____________
DATA AND CONCLUSIONS for: Rates of Chemical Reactions: A Clock Reaction
A.
Determination of Orders and Rate Constant of the Reaction
Reaction studied:
________________________________________________________________
[I-]
t
Calculate [I-1] below based on discussions shown in page 2 of this handout.
Rate law for this reaction = k[I-]n[S2O8-2]m = -
Provide sample calculations for finding initial concentrations of reactants [I-] and [S2O8-2] for Reaction 2.
Date Table for Reactions at Room Temperature ( range: _________to __________C )
Reaction
Δt (sec) for color to appear
(record two consistent runs)
Initial Conc in Reaction Mixture
[I-]
(average)
0.080 M
1
[S2O8-2]
(average) Rate of
Reaction - Δ[I-]/Δt
0.080 M
2
3
4 Catalyst
Date Table for Reactions at Different Temperatures
Reaction / Temp
(C)
Δt (sec) for color to appear
(record two consistent runs)
(average)
Initial Conc in Reaction Mixture
[I-]
[S2O8-2]
1. ______
5. ______
6. ______
Page 5
(average) Rate of
Reaction - Δ[I-]/Δt
Consider the relative rates of reactions 1, 2 and 3. The changes in the rates resulted only from differences
in [I-] and in [S2O8-2]. Equation (2) on page 1 can be written in the form:
Relative rate = k[I-]n[S2O8-2]m
(3)
where k is a relative rate constant. Our purpose is to find n and m in (3), above. These are the orders of
the reaction with respect to I- and S2O8-2 respectively.
Reactions 1 to 3 must all obey equation (3), with the same value of k, since they were all carried out at the
same temperature and have relative rates that were calculated on the same basis. One way of doing this is
by substituting the experimental value for Reactions 1 and 2 into Equation (3):
Relative rate 1 =
= k[
]n [
]m
Relative rate 2 =
= k[
]n [
]m
Now, divide the first equation by the second, obtaining:
Take the log of both sides of the equation, then divide to solve for n.
=[
]n
n = _______ rounded value to use = _____ (Ordinarily this should be an integer or whole number
fraction. If your data clearly does not yield a whole number or whole number fraction, you may leave it as
is; keep up to 4 decimal places for calculating the rate constant, k.) This “n” is the order of the reaction
with respect to I-.
Follow similar math procedures and use the data from Reactions 1 and 3 to find the order of the reaction
with respect to S2O8-2.
Relative rate 1 =
= k[
]n [
]m
Relative rate 3 =
= k[
]n [
]m
Now, divide the first equation by the second, obtaining:
Take the log of both sides of the equation, then divide to solve for m.
=[
]m
m = _________ rounded value to use = _____
Now use the n and m as calculated above, determine the relative rate constant k for Reactions 1 to 3
using Equation (3) above. Remember to specify units!! Provide sample calculations below:
Reaction
1
2
3
Average the 3 rate
reactions = kavg
Rate
constant,
k
Page 6
B. Effect of a Catalyst on Reaction Rate - Review the relative rates of Reaction 1 and Reaction 4.
Briefly discuss below how much the rate changed (increased 2-fold? 5.8 times? Or what?) with the use of
a catalyst.
Page 7
C. Effect of Temperature on Reaction Rate and The Activation Energy.
Rate constant k and temperature are related according to the Arrehnius equation: ln k = Ea + constant
RT
Specify what each term means in this equation: ln = __________________________________
k = _______________________________
Ea = __________________________________
R = _______________________________
T = ___________________________________
Complete the following table to summarize your experimental data
Reaction Mixtures
1(kavg from page 6)
5
6
Rate constant (Calculated based on the
rate law found in page 6.)
Natural log of rate constant, ln k
(negative number)
Average temp of reaction mixture (C)
Average Temp in K
Inverse absolute temperature, 1/T, K-1
To evaluate Ea;
1. Prepare a graph of natural logarithm of k, ln k (negative, on y axis) vs. 1/T (positive on x axis).
You may do this in either one of two ways:
a. Use standard graph paper. Set up x, and y scales so that the data points will spread over at
least 75% of the space on this paper. Draw a straight line that best-fits all three data points.
Determine the slope of this line by using rise-over-run. Attach your graph to this report.
b. Use a spreadsheet graphing program (Excel® or other) and obtain a trend-line with linear
equation best-fitted to your three data points. Submit a hardcopy of your graph with the
computer generated linear equation printed on it.
2. Report here the slope of your line obtained from step 1 above = __________________
3. The slope of the line equals -Ea/R, where R = 8.314 J/mole-K . Show calculations below to find
Ea with correct units and significant figures for the reaction: 2I- + S2O8-2  I2 + 2SO4-2
Page 8