Chem 1B
Dr. White
Saddleback College
1 Experiment 26 - Kinetics
Objectives
 To determine the rate law for the reaction
between iodide and bromate under acidic
conditions
 To investigate the effect of temperature on rate
 To determine the activation energy for the the
reaction between iodide and bromate under
acidic conditions
Introduction
This experiment involves the study of the rate
properties, or chemical kinetics, of the following
reaction between iodide ion and bromate ion under
acidic conditions:
-
-
+
6 I (aq) + BrO3 (aq) + 6 H (aq) →
3 I2(aq) + Br (aq) + 3 H2O(l)
(Rxn 1)
This reaction proceeds at an easily measurable rate
that depends on the concentrations of the I , BrO3 ,
+
and H ions according to the rate law. For this
reaction, the rate law takes the form
- x
- y
+ z
rate = k [I ] [BrO3 ] [H ]
(1)
One of the main purposes of the experiment will be to
evaluate the rate constant, k, and the reaction orders
x, y, and z for this reaction. We will also investigate
the manner in which the reaction rate depends on
temperature and will 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:
2-
-
2-
I2(aq) + 2 S2O3 (aq) → 2 I (aq) + S4O6 (aq)
(Rxn 2)
As compared with Reaction 1, the rate of Reaction €
2
is essentially instantaneous. The I2 produced in
Reaction 1 reacts completely and instantaneously
22with S2O3 so that until the S2O3 is used up the
concentration of I2 is effectively zero. As soon as the
2S2O3 is used up, the I2 remains in solution and is
made evident by the reaction with starch. The
reaction of I2 with starch produces a blue color. The
appearance of the blue color tells us when a given
amount of BrO3 has reacted. We will not calculate
the actual amount of BrO3 that has reacted, but use
an arbitrary value of 1000. This will give us the
€
relative rate which is equal to 1000/t, where t is the
time it takes to turn blue.
In this experiment, we will carry out the reaction
- +
between BrO3 , I , and H ions under different
concentration conditions. Measured amounts of each
of these ions in water solution will be mixed in the
2presence of a constant small amount of S2O3 . The
time it takes for each mixture to turn blue will be
measured. The time obtained for each reaction will
be inversely proportional to its rate. By changing the
concentration of one reactant and keeping the other
concentrations constant, we can investigate how the
rate of the reaction varies with the concentration of a
particular reactant. Once we know the order for each
reactant, we can determine the rate constant for the
reaction. This is illustrated in the example below.
Example: The kinetics of the following reaction was
investigated: 2NO (g) + O2 (g) → 2NO2 (g). The
following data were collected:
Exp.
[O2] (M)
[NO] (M)
Relative Rate
-1
(M⋅s )
1
0.0010
0.0010
7.1
2
0.0040
0.0010
28.4
3
0.0040
0.0030
255.6
To find the rate law for the reaction we need to use
two experiments where the concentration of one
reactant is held constant. In experiments 1 and 2 the
concentration of NO is constant and the
concentration of O2 is changed.
Thus using
experiments 1 and 2 we can find the reaction order
with respect to O2:
rate2 k[O2 ]2x [NO]2y k(0.0040 M) x (0.0010 M) y
=
=
rate1 k[O2 ]1x [NO]1y k(0.0010 M) x (0.0010 M) y
x
28.4 M ⋅ s −1 ⎛ 0.0040 M ⎞
=
⎜
⎟
7.1 M ⋅ s −1 ⎝ 0.0010 M ⎠
4 = 4x
x =1
A similar process is used to determine the rate order
with respect to NO:
rate3 k[O2 ]3x [NO]3y k(0.0040 M) x (0.0030 M) y
=
=
rate2 k[O2 ]2x [NO]2y k(0.0040 M) x (0.0010 M) y
y
255.6 M ⋅ s −1 ⎛ 0.0030 M ⎞
=
⎜
⎟
28.4 M ⋅ s −1 ⎝ 0.0010 M ⎠
9 = 3y
y=2
Now that we know x and y, we can use the rate law
to determine the rate constant:
Chem 1B
Dr. White
2
€
rate = k[O2 ][NO]
rate
k=
[O2 ][NO]2
Using experiment 1,
7.1 M/s
k=
= 7.1x10 9 M -2 s -1
[0.0010 M][0.0010 M]2
In the second part of the experiment we will
investigate how the rate of the reaction depends on
temperature. In general the rate increases sharply
with temperature. By measuring how the rate varies
with temperature we can determine the activation
energy (Ea), the minimum energy required for the
reaction to occur, by making use of the Arrhenius
equation:
ln k = −
E a ⎛ 1 ⎞
⎜ ⎟ + ln A
R ⎝ T ⎠
(2)
In this equation, k is the rate constant at the Kelvin
temperature T, Ea is the activation energy, A is
€ constant for a given reaction and R is the gas
another
constant. By plotting ln(k) against 1/T we should
obtain, by Equation 2, a straight line whose slope
equals -Ea/R. From the slope of that line, we can
easily calculate the activation energy.
Saddleback College
2 Chem 1B
Dr. White
Procedure
Part 1: Dependence of Reaction
Concentration
Rate
on
See the tables below. Except for the indicated
amounts of solutions used, the 5 trials will follow the
same procedure.
1. Use the following procedure for each trial.
Between trials, rinse the flask with DI water
and shake out the extra water.
a. Measure into the 2 different flasks
(Flask I and Flask II) the reagents
shown in the tables. Use graduated
cylinders to measure all the
reagents.
b. Have a thermometer and a
stopwatch ready. Pour the contents
of Reaction Flask II into Reaction
Flask I and immediately start the
stopwatch. Swirl the reaction flask.
The instant the mixture turns blue,
stop the stopwatch.
Record the
elapsed time of the run in seconds.
Measure
and
record
the
temperature.
Table 1: Volumes and reagents for Reaction Flask I
(125 mL)
Reaction
Mixture
1
2
3
4
5
Volume of
0.010 M KI
(mL)
5
10
5
5
4
Volume of
0.0010 M
Na2S2O3
(mL)
5
5
5
5
5
Volume of
H2O (mL)
5
0
0
0
6
Table 2: Volumes and reagents for Reaction Flask II
(125 mL)
Reaction
Mixture
1
2
3
4
5
Volume of
0.040 M
KBrO3 (mL)
Volume of
0.10 M HCl
(mL)
5
5
10
5
2.5
5
5
5
10
7.5
Starch
Indicator
3 drops
3 drops
3 drops
3 drops
3 drops
Part 2:
Dependence of Reaction Rate on
Temperature
1. In this part of the experiment, use reaction
mixture 1 and follow the instructions above.
However, the reaction mixtures will be at the
following temperatures:
about 0°C, about
10°C, and about 40°C. Record the actual
temperature for each trial.
Saddleback College
3 a. 0°C: Use ice to cool your reaction
flasks to 0°C. Leave the flasks in the
ice for about five minutes before you
start the trial and during the reaction.
b. 10°C: Use a plastic tray with cool
water to get the solutions to about
10°C. Leave the flasks in the water
for about five minutes before you
start the trial and during the reaction.
c. 40°C:
Use the one of the
thermostatted baths set up for the
trial at 40°C. Leave the flasks in the
water for about five minutes before
you start the trial and during the
reaction.
2. Prepare an Arrhenius plot on Excel so that the
activation energy can be determined.
Chem 1B
Dr. White
Saddleback College
Name: ___________________________
4 Lab Day/Time: ______________
Partner: ___________________________
Experiment 26: Kinetics Data and Results
Part 1: Rate Law
Reaction
Time, t, for
Mixture
Color
Change (s)
1
2
3
4
5
Relative
Rate
(1000/t)
[I-] (M)
[BrO3-] (M)
[H+] (M)
Temperature
(°C)
Show how you determined the [I-] for Reaction Mixture 1 (it is diluted from the given stock
solution by mixing all the reagents together):
Show how you determined the [BrO3-] for Reaction Mixture 1 (it is diluted from the given stock
solution by mixing all the reagents together):
Show how you determined the [H+] for Reaction Mixture 1 (it is diluted from the given stock
solution by mixing all the reagents together):
Determine the value of the reaction order x. Show your work below.
Determine the value of the reaction order y. Show your work below.
Chem 1B
Dr. White
Saddleback College
Determine the value of the reaction order z. Show your work below.
Write the General Rate Law: ________________________
Determine the Rate constant for each trial and fill in the following table (include the units of k):
Reaction
Mixture
1
2
3
4
5
Average
Show your calculation for trial 1 below:
k
(
)
Briefly explain why k should have nearly the same value for each of the reactions.
For Reaction Mixture 5, use the average k and the appropriate concentrations to predict
(calculate) the relative rate. Then use this value to predict (calculate) the reaction time, t, for
Mixture 5. Show your calculations below.
relative ratepredicted________
tpredicted________
tobserved________
5 Chem 1B
Dr. White
Saddleback College
6 Calculate the percent difference of your reaction times.
% Difference = _________________
Part 2: Temperature Effects
Approximate
Actual
Time t for
Temperature Temperature
Color
(°C)
(°C)
Change (s)
Your Data:
0°C
10°C
20°C*
40°C
*Data from Reaction Mixture 1 in part 1.
Relative
Rate
(1000/t)
k
ln k
1/T (K-1)
Attach a copy of your Arrhenius Plot made in Excel.
Line of best fit from your Arrhenius Plot:______________________
Determine the activation energy (in kJ/mol) and show your calculation below.
________________
Chem 1B
Dr. White
Saddleback College
7 Experiment 26: Kinetics Postlab Questions
1.
A student studied the clock reaction described in this experiment. She set up a reaction mixture by mixing
10.0 mL of 0.010 M KI, 10.0 mL of 0.0010 M Na2S2O3, 20.0 mL of 0.040 M KBrO3 and 10.0 mL of 0.10 M HCl using
the procedure given. It took 40.0 seconds for the color to turn blue.
a.
She found the concentration of each reactant in the reacting mixture by realizing that the number of
moles of each reactant did not change when that reactant was mixed with the others, but that its
concentration did. The volume of the mixture was 50.0 mL. Find the concentration of each reactant.
-
-
+
[I ] = __________M; [BrO3 ] = __________M; [H ] = __________M
b.
What is the relative rate of the reaction (1000/t)?___________________
c. The student did a second trial in which she changed the amount of one reagent. She mixed 10.0 mL of
0.010 M KI, 10.0 mL of 0.0010 M Na2S2O3, 10.0 mL of 0.040 M KBrO3, 10.0 mL of 0.10 M HCl and 10.0 mL
of H2O. It took 81.0 seconds to turn blue.
Find the concentration of each reactant after mixing:
-
-
+
[I ] = __________M; [BrO3 ] = __________M; [H ] = __________M
d. What is the relative rate of the reaction in the second trial(1000/t)?_____________
-
e. Use the information from the two trials above to determine the order with respect to BrO3 .
n = _________(nearest integer)