CHEM 1411 Spring 2016
Experiment 5: The Iodination of Acetone
Background
Chemical kinetics is the study of the rates of chemical reactions—how fast or slow a reaction is.
A reaction rate is quantitatively expressed as the change in concentration (in M) that occurs for a
reactant or product per unit of time, usually in units of seconds. An average reaction rate for a
reaction can be determined by measuring the change in concentration for a finite period of time.
For example, for the generic reaction given below:
a A(g) + b B(g) →
c C(g)
the average reaction rate could be found by either of the following two equations:
average rate = –Δ[A]
Δt
= Δ[C]
Δt
where Δ[A] and Δ[C] are the changes in concentrations for A and C and Δt is the time interval
(in seconds) over which the change occurred. Reaction rates are always given as positive values,
so a negative sign is included in the first expression because the concentration of the reactant A
is decreasing during the reaction. The instantaneous reaction rate is the value of the rate at any
specific point during the reaction—it is the change in concentration for an infinitely small time
interval (e.g., –∂Δ[A] / ∂t). If the change in reactant concentration is graphed as a function of
time, the instantaneous reaction rate at a specific point is the slope of the tangent line at that
point in the graph.
For a given chemical reaction, the instantaneous rate usually decreases with an decrease in one or
more of the reactant concentrations. The relationship between the reaction rate and reactant
concentrations can be quantitatively expressed by the rate law for the reaction. For example, the
rate law for the generic chemical reaction given above would be expected to take the form:
rate = k [A]m[B]n
where [A] and [B] are the concentrations (in M) for the reactants A and B, and k is a
proportionality constant called the rate constant for the reaction. The value for exponents m and
n in the rate law give the order of the reaction with respect to the reactant concentrations. If m is
1, the reaction is said to be first order with respect to reactant A, and if n = 2, the reaction is
second order with respect to reactant B. The overall order for the reaction is the sum of m and n;
in this example the reaction would be third order overall (m + n = 3). Note that the value for m
and n are not necessarily the same as the stoichiometric coefficients in the balanced equation—m
does not have to equal a in the rate law.
The values for the exponents in a typical rate law are usually 0, 1, 2, or 3. Because the rate of the
reaction depends on its specific mechanism (and not just the identity of the reactants and
products), the values for the exponents in the rate law must be determined experimentally by
measuring the instantaneous rate for the reaction at different reactant concentrations. The
instantaneous rate is usually measured experimentally as an initial rate—the rate at the start of
the reaction (t = 0).
CHEM 1411 Spring 2016
Experiment 5: The Iodination of Acetone
Table 1. Initial rate data for the decomposition of N2O5 at 65.0 °C.
2 N2O5(g)
→ 4 NO2(g) + O2(g)
Trial
[N2O5] (M)
Initial Rate (M/s)
1
0.250
0.00131
2
0.500
0.00266
The experimental determination of the rate law can be demonstrated for the decomposition of
N2O5 into NO2 and O2 gas at 65.0 °C. The rate law for this reaction is:
rate = k [N2O5]m
The initial rate of the reaction is measured for two different initial concentrations of N2O5 (see
Table 1 above). The value for the exponent m can then be determined by using a ratio of the rate
law for the two trials (see below). Because the value for k is constant at a given temperature, the
rate constants can be eliminated from the rate law ratio, and the ratio of the initial rates for the
trials is equal to the ratio of the initial N2O5 concentrations raised to the m power. Solving the
ratio expression for m (either by inspection or using logarithms) for this reaction gives an m
value of 1. This reaction is therefore first order with respect to N2O5; the instantaneous rate is
equal to the product of the rate constant and the N2O5 concentration.
ratetrial 2 = k [N2O5]trial 2m
________
____________
ratetrial 1 = k [N2O5]trial 1m
2.03 = 2.00m
0.00266 M/s = (0.500 M)m
____________ __________
0.00131 M/s = (0.250 M)m
m = 1
rate law: rate = k [N2O5]
average value for k = 0.00528 s–1
Once the value for m is known, the experimental rate data for the decomposition of N2O5 given
in Table 1 can also be used to determine an experimental value for the rate constant at 65.0 °C.
Solving the rate law for k (k = rate / [N2O5]) and using the rate and concentration values for
both trials gives k = 0.00524 s–1 for Trial 1 and k = 0.00532 s–1 for Trial 2; the average value for
k for this reaction is therefore 0.00528 s–1. With this information, the instantaneous reaction rate
for this reaction at 65.0 °C can now be calculated for any given concentration of N2O5.
In this experiment, you will experimentally measure the initial reaction rate for the iodination of
acetone (see Equation 1 on the next page) for four different reaction mixtures. You will be able
to use this experimental rate data to determine the rate law for the reaction and to calculate an
average value for the rate constant, k, for the reaction. The rate law for the iodination reaction is:
rate = k [acetone]m[I2]n[H+]p
Note that you will be investigating the unusual possibility that the rate of the reaction depends on
concentration of a product (H+ ions) in addition to the concentrations of the reactants.
CHEM 1411 Spring 2016
Experiment 5: The Iodination of Acetone
Equation 1. The iodination of acetone.
O
O
+
C
CH3
CH3
I2(aq)
+
C
CH3
H+(aq)
+
I–(aq)
CH2I
The iodination reaction of acetone is a convenient reaction to study kinetically because the
instantaneous reaction rate for this reaction can be experimentally measured as an average rate
over the entire course of the reaction. The iodination reaction is zero order with respect to I2
concentration (n = 0); the rate of the reaction does not depend on the concentration of I2. You
will measure the time required (in seconds) for a known initial concentration of I2 to be
completely consumed in the presence of a large excess of acetone and H+ ions. This can be done
visually by monitoring the disappearance of the yellow color of the I2 from the reaction mixture.
The average rate of the reaction can then be found using the equation given below:
rate = –Δ[I2]
Δt
where Δ[I2] the change in concentration for I2 (and is equal in magnitude to the initial
concentration) and Δt is the time required for the reaction to go to completion. Because a large
excess of acetone and H+ is present in the reaction mixtures, their concentrations will remain
essentially constant over the course of the reaction. The instantaneous reaction rate for the
iodination reaction will therefore remain effectively constant for the duration of the reaction and
will be equal to the average reaction rate determined using the above equation. By measuring the
reaction rate for several different reaction mixtures in which the concentrations of acetone, H+,
and I2 are systematically varied, you will be able to determine the values for the exponents m and
p, and verify that the value for n is 0 in the rate law. Once the rate law for the iodination reaction
is known, you will then be able to calculate the average value for the rate constant, k, for this
reaction at room temperature.
CHEM 1411 Spring 2016
Experiment 5: The Iodination of Acetone
Experimental Procedure
•
•
•
Eye protection must be worn at all times during the experiment.
Use caution when handling the hydrochloric acid and iodine solutions, as they can burn or
stain your skin.
Dispose of the reaction mixtures and any excess acetone and iodine solution in the waste
containers in the hood. Dispose of all other waste solutions down the drain with copious
amounts of water.
Table 2. Reaction mixtures for the iodination of acetone trials.
Volume (mL)
1.
2.
3.
4.
Mixture
4.0 M acetone
1.0 M HCl
0.0050 M I2
DI H2O
1
2
3
4
5.0
10.0
5.0
5.0
5.0
5.0
10.0
5.0
5.0
5.0
5.0
10.0
10.0
5.0
5.0
5.0
Obtain about 60 mL of each of the following solutions in clean, dry 100 or 150 mL
beakers: 4.0 M acetone, 1.0 M HCl (a source of H+ ions), and 0.0050 M I2. Obtain several
small test tubes, clean the test tubes (if needed), rinse them with deionized water and shake
them dry. Then fill one of the test tubes about two-thirds full with DI water. Obtain a
stopwatch (if needed) and a digital thermometer.
Using a graduated cylinder to measure the volumes, add 5.0 mL of 4.0 M acetone solution,
5.0 mL of 1.0 M HCl solution, and 10.0 mL of DI water to a clean and dry 125 mL
Erlenmeyer flask (reaction mixture 1 in Table 1).
Condition a separate clean graduated cylinder by rinsing the inside of the cylinder twice
with a few milliliters of the 0.0050 I2 M solution. Then measure out 5.0 mL of 0.0050 I2 M
solution in the cylinder. Pour the I2 solution into the Erlenmeyer flask and at the same time
start the timer on the stopwatch. Quickly swirl the flask to mix and then fill a clean and dry
test tube with the reaction mixture to the same depth as the tube filled with DI water.
Hold the tubes over a sheet of white paper and look down the length of the tubes. As the
iodination reaction proceeds, the yellow color of the iodine solution will fade. Stop the
timer on the stopwatch when this yellow color just disappears—when the color of the
reaction solution matches the color in the test tube filled with DI water. Record the reaction
time for this trial (in seconds) in your Data and Calculations pages. Measure and record the
temperature of the reaction mixture (in °C).
CHEM 1411 Spring 2016
5.
6.
Experiment 5: The Iodination of Acetone
Carry out a second trial for reaction mixture 1, using the test tube containing the reaction
mixture from the first trial as a reference point for the end of the reaction (in place of the
tube filled with DI water). Record the reaction time and temperature for this trial in your
Data and Calculations pages.
Determine the rate of the iodination reaction (in M/s) for reaction mixture 1. Calculate and
record the average time (in seconds) required for the reaction to finish for the two trials.
Then use the equation given below to find the rate of the reaction (where Δt is the average
reaction time). This average reaction rate is also effectively the initial instantaneous rate for
the reaction under these conditions. Note that the initial concentration of I2 in the reaction
mixture is not 0.0050 M, as this concentration was diluted by the addition of the water,
acetone, and HCl solutions.
rate = –Δ[I2]
Δt
7.
8.
Following the procedure given for reaction mixture 1, experimentally determine the
reaction rate for the iodination reaction for reaction mixtures 2–4. Carry out two trials for
each reaction mixture and use the average reaction time for the trials to calculate the rate
for that mixture. Include the data and calculations needed for the determination of the
reaction rate for these reaction mixtures on your Data and Calculations pages.
Use your reaction rate data and the initial concentrations of acetone, I2, and H+ (which
equals the HCl concentration) for the reaction mixtures to determine the rate law for the
iodination of acetone. Compare your measured reaction rates for reaction mixtures 1 and 2
to find the value for m and determine the order of the reaction with respect to the acetone
concentration. Comparison of the rates for mixtures 1 and 3 will allow for the
determination of the value for p, and comparison of the rates for mixtures 1 and 4 should
confirm that the reaction is zero order with respect to I2 (n = 0). Include all of the
calculations needed to determine the rate law for the iodination reaction in your Data and
Calculations pages.
rate = k [acetone]m[I2]n[H+]p
9.
Use your experimental rate law for the iodination reaction and the initial concentrations for
the reaction mixtures to determine the value for the rate constant, k, for each reaction
mixture. Then calculate the average value for k for the iodination reaction. Include all of
the calculations for the average value of k in your Data and Calculations pages.
Note: Inform your instructor if the measured temperature for any of your reaction trials is
substantially different (greater than ±3 °C difference) from the rest of the trials. If time
permits, your instructor may direct you to repeat the trial(s) that took place at the
significantly different temperatures.
CHEM 1411 Spring 2016
Experiment 5: The Iodination of Acetone
Data and Calculations Page
Record all of your group’s experimental data below. Include all data and calculations needed to
determine the reaction rate for each of the four reaction mixtures, the rate law for the reaction
(with the experimental values for the exponents m, n, and p), and the average value for the rate
constant, k, for the iodination of acetone reaction. You are encouraged to organize your data for
the reaction mixtures using one or more tables. Attach additional pages as needed.