EXPERIMENT 3.1
CHEMICAL KINETICS
INTRODUCTION
In this experiment, you will study the dependence of the rate of a reaction on
reactant concentration and temperature. The concentration dependence will be
summarized by solving for an empirical rate law. The sensitivity of the reaction
rate to temperature will be indicated by finding the value of the activation energy of
the reaction. The objectives of the experiment are to practice the mathematics of
rate laws and the Arrhenius equation with real experimental data and to review
dilution calculations.
For reactions of the generic form
aA + bB  cC + dD
where A, B, C and D are species with reaction equation coefficients a, b, c and d, it
is expected that the rate law expressing the dependence of the initial rate on the
starting concentrations of reactants will take the form
Rate = k[A]x[B]y.
In this law, the square brackets signify the molar concentration of the species
enclosed, k is a constant of proportionality called the rate constant, and the
exponents are the orders of the rate law with respect to each species concentration.
It is the orders that indicate how influential the concentration of each reactant is in
determining the reaction rate. The greater the order, the greater the change in the
rate when the concentration is changed. To find a rate law for a particular reaction
and temperature is to experimentally determine the values of the rate law orders
and the constant. The result is useful for the practical value of being able to use it
to predict the reaction rate for a set of concentrations and for theoretical studies in
which the orders may give clues to help determine the specific sequence of
microscopic collisions by which the reaction occurs, the mechanism.
The reaction to be studied in this experiment is the redox reaction
6I-(aq) + BrO3-(aq) + 6H+(aq)  3I2(aq) + Br-(aq) + 3H2O(l).
With three reactants, we expect a rate law of the form
Rate = k[I-]x[BrO3-]y[H+]z
The standard approach to solving the law is to measure the rate using an initial set
of concentrations, then to systematically change the concentration of one reactant at
a time and measure the new rates. The effect of each change allows an order to be
found. When all of the orders are known, one can solve for the constant k at the
temperature observed.
Of course, the first requirement in this approach is to have a method for the rate
measurement. That is, we must be able to determine the number of moles or moles
per liter of change in one of the reactants or products over a measured time span
short enough that the reactant concentrations will not change significantly. To
accomplish this, we will introduce a secondary reaction to provide a visual indicator
for the moment that a fixed extent of reaction has occurred. In the presence of a
starch indicator, the molecular iodine, I2, produced by the primary reaction gives a
blue color to the solution. If only starch and the three reactacnts are combined, one
will see this blue color develop immediately. However, if thiosulfate ion, S2O32- , is
also present in the reaction mixture, it will react quickly with the I2 as it is
produced by the reaction
I2(aq) + 2S2O32-(aq)  2I-(aq) + S4O62-(aq).
Thus as long as there is excess thiosulfate, the I2 will be consumed and the blue
color will not develop. If the thiosulfate is used up, the blue will appear. We will
add the same number of moles per liter of thiosulfate to each reaction trial and
make it the limiting reactant. Thus if we measure the time required for the
appearance of the blue color, we can calculate a reaction rate in units of moles per
liter of thiosulfate per second..
After using rate measurements to find the rate law, we will test it by solving for the
rate for a new set of concentrations and comparing our result to experimental
measurement.
The second objective of our study is to find the dependence of the reaction rate
constant k on the temperature. This dependence is indicated by the value of the
activation energy of the reaction, Ea. The activation energy is the increase in
potential energy of the reacting particles necessary to break bonds and allow for the
change from reactants to products. The higher the amount of this energy barrier,
the more the reaction rate will change as temperature changes. We will find its
value by measuring the rate at several temper-atures using constant initial
concentrations and graphing the logarithm of the rate constants versus the
reciprocal of the absolute temperatures. According to the Arrhenius
equation, the slope of such a graph should equal the negative of the activation
energy divided by the gas law constant R.
MATERIALS AND SAFETY
The reactant concentrations that will be used are all less than 0.10 M. All
materials can be disposed of in the sinks after trial measurements are made.
Nevertheless, take care to avoid unnecessary contact with the strong acid HCl and
the oxidizing agent KBrO3. Wear safety goggles at all times.
PROCEDURE
PART A. EMPIRICAL RATE LAW
Refer to the table below for the volumes of reactant solutions and water to be mixed
together for each of five timed trials. Each material will be in a reservoir with a
large and small graduated cylinder dedicated to it. Use the large cylinder to collect
solutions from the reservoir, then pour from the large cylinder into the small
cylinder to measure the amount to add to your reaction mixture. Leave excess
solution in the large cylinder for the next person to use. We do not want the
reaction to begin until all ingredients are together, so the oxidizing and reducing
agents must be kept apart. Use two 250 mL Erlenmeyer flasks for each trial. Have
one team member gather the three components specified for flask 1 while the other
partner gathers the three components for flask 2. Meet back at the drawer area,
measure the temperature in one of the flasks, then pour them together and begin
timing. Record the time required for the blue color to appear.
Trial
1
2
3
4
5
0.01 M
KI
10 mL
20 mL
10 mL
10 mL
7.5 mL
Flask 1
0.001 M
Na2S2O3
10 mL
10 mL
10 mL
10 mL
10 mL
H2O
10 mL
0 mL
0 mL
0 mL
10 mL
0.04 M
KBrO3
10 mL
10 mL
20 mL
10 mL
7.5 mL
Flask 2
0.1 M
HCl
10 mL
10 mL
10 mL
20 mL
15 mL
Starch
3 drops
3 drops
3 drops
3 drops
3 drops
PART B. TEMPERATURE DEPENDENCE
Consider trial 1 from part A as the first data point in this section also. Using the
same reactant amounts as in trial 1, find the times required for at least two
(preferably three) reactions run at different temperatures. To achieve higher
temperatures, collect the ingredients in the two flasks as before. Obtain two plastic
bowls from the cart and fill them with hot tap water. Immerse the flasks in the hot
tap water, holding a thermometer in one of them. Do not transfer the thermometer
between flasks. When the temperature is within a few degrees of your objective,
mix the two solutions, start timing, and note the temperature of the mixture. Try to
hold this temperature for the duration of the reaction. If the temperature begins to
fall, place the flask back in the hot water. If it rises above the set point in the hot
water, take the flask out again. When the color change occurs, record both the time
and the temperature of the trial.
To achieve a lower temperature, mix about four parts cold tap water with one part
ice in the plastic bowls. Immerse the flasks in the ice baths until one reaches the
target temperature (avoid letting the flask temperatures go below 5°C), then
combine. Use the ice bath in the same way as the hot water bath to keep the
reaction temperature as constant as possible once the reaction has been started.
Record the time and temperature when the color change occurs. After completion,
ask if another group can make use of the ice bath before discarding it.
Check the amount of time available to determine the number of trials you can
complete. If time permits, add trials at temperature in the areas of 40°C, 30°C, and
5°C. If you are running out of time, omit the 30°C trial. Clean up and return
materials to the cart or drawer when finished.
CALCULATION NOTES
As the reaction mixtures are prepared and then combined, each of the reactants is
diluted from the concentrations in the reservoirs by the increase in solution volume.
Recall that the resulting concentration can be found by the equation below which is
based on the number of moles of solute staying constant during dilution.
Mconc × Vconc = Mdil × Vdil
Use this equation to find the initial concentrations of reactants in the combined
solutions as directed in practice question 1.
To determine the orders of the rate laws, consider two trials in which only the
concentration of species A changes from [A]1 to [A]2. If the reaction rate changed
from Rate1 to Rate2, then the order of the rate law with respect to [A] is found from
the relation
Rate2  [A]2 


Rate1  [A]1 
x
By taking the logaritm of both sides of the equation we obtain
 Rate2 
log

 Rate1 
x=
 [A]2 
log

 [A]1 
Rate law orders are frequently found to be whole numbers (although not always) so
you may use judgement in determining if it is reasonable to round to nearest whole
numbers based on the calculated results.
PRELAB QUESTIONS
1.
To facilitate the calculation if the rate law orders, calculate the concentration
of each reactant in trials 1 through 5 (part A) after the flask contents have
been combined but before reaction takes place. In the table below [I-] = [KI],
[BrO3-] = [KBrO3], and [H+]=[HCl]. We only list the ions, because they are
what react in the net ionic equation. Enter the concentrations you calculate
into your table in your lab notebook. Create and fill in this table in your
laboratory notebook as part of your prelab:
Trial
1
2
3
4
5
[I-]
____________
____________
____________
____________
____________
[BrO3-]
____________
____________
____________
____________
____________
[H+]
____________
____________
____________
____________
____________
2. Determine the molarity of the thiosulfate ion initially present in the combined
reaction mixture. Have the insructor check your work. This is the number you will
divide by the reaction time in seconds to obtain the relative rate for each trial.
PART A DATA – create the following table in your lab notebook as part of your
prelab.
For each trial, record the reaction temperature (to verify that is does not vary
significantly) and the time required for the reaction to consume the
thiosulfate in seconds. Then, referring to practice question 2, calculate the
rate as moles per liter of thiosulfate per second.
TRIAL
1
2
3
4
5
Temperature
____________
____________
____________
____________
____________
Time
____________
____________
____________
____________
____________
Rate
____________
____________
____________
____________
____________
PART B DATA – Create the following table in your lab notebook as part of your
prelab.
Copy the data from trial 1 of part A as the first data point in this section. Then,
for each trial at different temperature, record the reaction temperature and time
in seconds. Calculate the rate for each trial.
TRIAL
1
6
7
8
Temperature
____________
____________
____________
____________
Time
____________
____________
____________
____________
Rate
____________
____________
____________
____________
POSTLAB CALCULATIONS
PART A
To determine the rate law, you will first find the order with respect to each
reactant. Show work in the space provided and use logarithms to find the real
number value of each order. If appropriate and reasonable based on results, round
to whole numbers.
b) Compare trials 1 and 2 to determine the rate law exponent for [I-]:
Rate law exponent for [I-]?
c) Nearest whole number?
d) Compare trials 1 and 3 to determine the rate law exponent for [BrO3-]:
Rate law exponent for [BrO3-]?
e) Nearest whole number ?
f) Compare trials 1 and 4 to determine the rate law exponent for [H+]:
Rate law exponent for [H+]?
g) Nearest whole number ?
h) Write the complete rate law including all exponents.
Next, you are to find the value and units of the rate law constant, k.
i) Substitute the values for the concentrations and rate for trial 1 and solve for k.
k1 ____________
j) Substitute the values for the concentrations and rate for trial 2 and solve for k.
k2 ____________
k) Substitute the values for the concentrations and rate for trial 3 and solve for k.
k3 ____________
l) Substitute the values for the concentrations and rate for trial 4 and solve for k.
k4 ____________
m) Average of k values ____________
As a test of the rate law, use the average value of k and the concentrations of trial 5
to calculate a predicted rate for that trial. Compare the result to your measured
rate (item (a)) by calculating the percent difference between them.
n) Predicted rate?
o) Percent difference from measured?
POSTLAB CALCULATIONS
PART B
l) For each trial, calculate a rate constant, k (which should vary with the
temperature). Also find the natural logarithm of k, ln(k); convert the
temperature to Kelvins; and calculate the reciprocal of the Kelvin temperature.
TRIAL
k
1 ____________
6 ____________
7 ____________
8 ____________
ln(k)
____________
____________
____________
____________
T (in K)
____________
____________
____________
____________
1/T
____________
____________
____________
____________
On the graph paper in your lab notebook, or using the page that follows, or using a
spreadsheet program plot ln(k) versus 1/T. Choose appropriate scales so that the
minimum and maximum values are well separated in your graph plotting ln(k)
versus 1/T. Use a straight edge and/or calculator or spreadsheet program to draw a
line that best fits the data and measure its slope. If you rely on a calculator or
spreadsheet application for slope determination, make a note in your notebook and
summarize what you did.