Experimental Determination of the Rate Law
for the Reaction of Iodine with Acetone1
Objective
The rate law for the reaction of iodine with acetone will be determined by using the method of
initial rates. In this method, the reaction is run several times with differing initial concentrations of
acetone, iodine, or hydrogen ion. For each reaction, the initial rate of reaction is determined by
measuring the time for the color of iodine to disappear from the reaction mixture. The reaction
orders with respect to acetone, iodine, and hydrogen ion, respectively, will be determined by
finding the dependence of the reaction rate on each of the initial concentrations.
Introduction
Experimentally, it has been found that the rate of a chemical reaction is equal to an expression
called a rate law. Rate laws tend to have the general form of:
Rate = k [A]m[B]n[C]p
where the species A, B, or C may be a reactant, product, or even a catalyst in the overall
chemical reaction. The rate exponents (m, n, and p) are generally small whole numbers, such as
0, 1, or 2. In some cases, the rate exponents are negative numbers (–1 or –2) or half-integers (1/2
or –1/2).
In general, the value of the rate exponents bears no relationship to the stoichiometry of the
overall reaction. Therefore, you generally cannot determine the rate law by looking at the
balanced chemical reaction. The rate law must be determined experimentally by running the
reaction and measuring the reaction rate.
One experimental method that is used to determine rate laws of chemical reactions is the
method of initial reaction rates. In this method, the chemical reaction whose rate law is to be
determined is run a number of times while varying the initial concentrations of the species that
are thought to appear in the rate law. The initial rate of reaction is measured each time the
reaction is run. The dependence of the initial rate on the initial concentration of each species
will give the value of the rate exponent for that particular species.
A typical initial rate experiment might generate data like this.
[A] (moles/L)
0.100 M
0.200 M
0.200 M
[B] (moles/L)
0.200 M
0.200 M
0.400 M
Initial rate (M/min)
5.0 x 10-4
2.0 x 10-3
1.0 x 10-3
The strategy is to find two experiments in which only one of the initial concentrations is changing
and find the dependence of the rate on the change in initial concentrations. For example,
comparing experiments 1 and 2, notice that the initial concentration of A changes from 0.10
mol/L to 0.20 mol/L while the initial concentration of B remains constant. The rate law for
experiments 1 and 2 are:
Rate2 = k
[ A]m2 [ B]m2
Rate1 = k
[ A]1m [ B]1m
Divide the rate law for experiment 2 by the rate law for experiment 1. The rate constant, k, will
cancel out as will the [B] since the [B] is the same for both experiments. All that remains is
Rate2
Rate1
=
[ A] m2
[ A]1m
When the experimental data is inserted into the equation, you get
2.0 x 10-3 M/min
5.0 x 10-4 M/min
=
4
2
(0.200 M)m
(0.100 M)m
=
=
2m
m
The order for the reaction with respect to A is second-order. Similar calculations would let you
determine the order of the reaction with respect to B.
Once the order of the reaction for A and B have been determined, you can use any of the
experiments to calculate the rate constant, k.
Rate
[A]m[B]n
=
k
The units on the rate constant vary depending on the overall order of the rate law.
The Experiment
In today’s experiment, you will determine the rate law for the reaction of iodine with acetone to
produce iodoacetone:
CH3COCH3 (aq) + I2 (aq)  CH3COCH2I (aq) + HI (aq)
Using structural formulas for acetone and iodoacetone, this equation is represented as:
O
H3C
O
CH3
+ I2
H+
+ HI
H3C
CH2I
Since this equation is catalyzed by the hydrogen ion, H +(aq), the concentration of H+ appears in
the rate law as well as the concentrations of acetone and I 2.
The general form for the rate law for this equation is:
Rate = k [CH3COCH3]M [I2] n [H+]p
You will determine the value of the rate exponents (m, n, and p) and the value of the rate
constant k using the method of initial rates.
In this method, you will measure the rate of reaction for several runs of the reaction, varying the
initial concentration of one species in the rate law at a time. The rate of reaction will be
determined by measuring the time it takes for the brown color of iodine in aqueous solution to
disappear. You will measure the brown color spectroscopically using a UV-visible spectrometer.
The rate of the reaction will be equal to the average rate of the reaction over the time period
that the reaction is studied. The assumption is that all of the iodine is used up in the course of the
reaction. The negative sign is included because iodine is a reactant.
Average rate = - [I2]final - [I2]initial
timefinal - timeinitial
Average rate = - 0 M - [I2]initial
timefinal - 0 sec
The experiment will be run a total of 7 times with different initial concentrations of acetone,
iodine, and hydrogen ion. The recipes for making the seven different solutions are given in table
below. The initial concentrations of each species are determined by using the dilution formula.
Figure 1 Schematic Diagram for a spectrophotometer
Any solution that is colored will absorb at a particular wavelength of visible light. An instrument
that is used to measure the absorbance of light by a sample is called a spectrophotometer. A
general schematic diagram for a spectrophotometer is shown above in Figure 1.
It begins with a light source or a light bulb. For the instrument that is used in this lab, the light bulb
emits light of wavelengths ranging from 340 nm to 600 nm. The light then travels to a
monochromator. The monochromator separates the light into its individual wavelengths so that
light of a particular wavelength shines towards the sample. A simple monochromator is a glass
prism. In modern instruments, a polished grating is used to separate the light into its individual
wavelengths. After passing through the monochromator, light of a particular wavelength with
intensity equal to I0 is passed through the sample. As the light passes through the sample, part of
the light may be absorbed. This absorption lowers the intensity of the light. The intensity of the
light after it passes through the sample is equal to I. The detector measures this intensity.
The spectrophotometer expresses the amount of light absorbed in one of two ways. The first way
of measuring the amount of light absorbed is called the percent transmittance, %T, which is
defined as follows:
%T = I x 100
Io
The second way of measuring the amount of light absorbed is called the absorbance, A. It is
defined as follows:
A = 2 - log %T
When a sample completely absorbs no light at all (all light is transmitted), %T = 100 and A = 0.000.
As a sample absorbs a larger fraction of light (less light is transmitted), %T will decrease and A
(the amount of light absorbed) will increase.
Laboratory Procedure
This experiment will be performed with lab partners. Although the experimental procedure is
performed in pairs, the calculations should be performed separately.
Figure 2: A typical Spectronics 20 spectrometer
1. First, the Spectronics 20 spectrometer must be calibrated at the proper wavelength.
A picture of the important knobs on the spectrometer is shown in Figure 2 above. Set
the wavelength of the spectrometer at 410 nm. With nothing in the sample-holder,
adjust the power switch knob until the %T reading is equal to zero. Obtain a cell and
fill it between half and three-quarters full with deionized water and place it into the
sample holder. Adjust the transmittance/absorbance knob until the %T reading is
equal to 100%. The spectrometer is then ready for use.
2. Set up four burettes on each lab bench. Label a burette for acetone, for HCl, for
water, and for iodine. Fill each burette with the indicated solution. The burettes will
be shared by all of the people working on a bench.
3. Obtain 14 small beakers. As specified in the table below, dispense from the burette
the required amounts of each solution into the beakers. Beaker A contains the
acetone, HCl, and deionized water. Beaker B contains the iodine solution. Be sure to
keep the beakers for each run ordered in the proper order. Label a piece of paper
toweling like the one shown in the picture below to keep the beakers clearly
identified.
Beaker A
Beaker B
1
2
3
4
5
6
7
Table 1
Amounts of each solution to be measured for each run
BEAKER A
Beaker B
Run
Number
1
4.0 M acetone
1.0 M HCl
deionized H2O
0.0050 M I2
3.00 mL
3.00 mL
8.00 mL
4.00 mL
2
6.00 mL
3.00 mL
5.00 mL
4.00 mL
3
9.00 mL
3.00 mL
2.00 mL
4.00 mL
4
5
6
7
3.00 mL
3.00 mL
3.00 mL
3.00 mL
6.00 mL
9.00 mL
3.00 mL
3.00 mL
5.00 mL
2.00 mL
4.00 mL
0.00 mL
4.00 mL
4.00 mL
8.00 mL
12.00 mL
4. Take the cell out of the spectrometer and empty the contents. For run #1, pour the
contents of beaker A into beaker B and quickly mix the solution by stirring. As one
person begins mixing the solutions, the other partner should start the timing the
reaction. Quickly pour a small amount of the mixed solution into the spectrometer
cell and swirl in the cell and discard. Then fill the cell about ¾ full with the solution and
place in the spectrometer.
5. The spectrometer should be reading a certain %T reading dependent on the amount
of iodine contained in the solution. As the reaction proceeds and the iodine is
consumed, the color will disappear and the %T reading will increase. (More light will
be transmitted through a clear solution than a colored solution.) Therefore, the %T
reading on the spectrometer should increase at a fairly constant rate. The reaction
can be considered complete when the %T reading stops increasing. This will usually
occur in the 90%-100% region. Stop timing when the %T reading stops increasing. You
may want to record the earliest time at which you observe % becoming constant
and then continue to observe the change in %T for an addition minute. This will allow
you to confidently record the time it takes for the reaction to go to completion.
Record the time elapsed on your data sheet.
6. Repeat steps 4 and 5 for the seven sets of initial concentrations as prescribed in table
above.
Calculations
1. Calculate the initial concentrations of acetone, iodine, and hydrogen ion by
applying the dilution formula:
MinitialVinitial = MfinalVfinal
2. The initial concentration of acetone, hydrochloric acid, and iodine can be found on
the reagent bottles. The initial volumes are the measured volumes for each run as
specified in table above. The final volume is the total volume of the mixed solution,
which for each run is 18.00 mL.
3. The rate of reaction for each run is calculated by finding the average rate of the
reaction as explained in the introduction.
4. Calculate the value of m, of n, and of p. Round to the nearest whole number. Each
order (m, n, and p) can be calculated two different ways. For example, comparing
runs 1 and 2 will give you a value for m. Comparing runs 1 and 3 will also give you a
value for m. The two values should be within 10% of each other or you should repeat
runs 1-3. Show both possible calculations for m, for n, and for p. Calculate the
percent difference for between each pair of values for m, for n, and for p.
5. Calculate the value of the rate constant k by placing in the initial concentrations
and the initial rates into the rate law and solving for k for each run. You should
calculate a value of the rate constant k for each run and average all 7 values to
come up for a value for your rate constant. Don’t forget to put to appropriate units
on your value of the rate constant.
1Based
on an experiment designed by Dr. Arthur A. Low at Tarleton State University
Department of Geosciences and Chemistry