See discussions, stats, and author profiles for this publication at: https://www.researchgate.net/publication/268278075
Reaction Kinetics of the Iodination of Acetone
Conference Paper · October 2014
DOI: 10.13140/2.1.4981.7288
CITATIONS
READS
0
104,732
3 authors, including:
Savannah Brooke Carroll
University of Tennessee
3 PUBLICATIONS 1 CITATION
SEE PROFILE
All content following this page was uploaded by Savannah Brooke Carroll on 15 November 2014.
The user has requested enhancement of the downloaded file.
Chemistry 479 Experiment 4
Reaction Kinetics of the Iodination of Acetone
Savannah Carroll
Lab Partner: Alexis Hughes
Department of Chemistry
University of Tennessee, Knoxville, TN
November 5, 2014
1
Abstract
Reaction kinetics for the iodination of acetone, a color changing reaction, in the presence
of an acid catalyst were studied using spectrophotometer constructed in the lab. These results
were equivalent to what may be recorded using a commercial spectrophotometer. The purpose of
this experiment is to determine if the iodination reaction is a zero, first, or second order reaction.
By systematically varying the concentrations of the reactants, the rate law is determined. The
order of the reaction was also determined by creating an absorbance versus time plot with a
linear fit for the rate law.
Introduction
The purpose of this experiment is to determine if the iodination of acetone is a first,
second, or third order reaction. The time of the reaction is the time it takes for the brown color of
the iodine to disappear meaning the concentration of iodine has become zero—this is measured
by the spectrophotometer. The rate at which a chemical reaction occurs is dependent of the
nature of the reaction, concentration of the reactants, temperature, and the presence of catalysts
(in this case HCl). In this experiment the kinetics of the reaction between iodine and acetone in
acid solution were studied, and the reaction proceeds as follows:
(Meyer)
The rate of this reaction is dependent upon concentration of the hydrogen ion and concentration
of the reactants (Slowinski, Wayne, Masterton 1973). The rate law is therefore: Rate = k
2
(acetone) m (I2) n (H+) p where m, n, and p are orders of the reaction and k is the rate constant for
the reaction. The rate law is useful in the description of substances that can influence the rate of
the reaction. These influences can be grouped into two categories: those whose concentrations
change with time, and those whose concentrations do not necessarily change with time. These
influences can be changed between experiments to show how they influence the rate of a
reaction (Tinoco 1978).
The order of the reaction, that describes the way the rate of the reaction depends on
concentration, may be determined by plotting absorbance data from each reaction versus time
and utilizing the most linear function. Also from these plots, if the reaction is zero order with
respect to iodine as the literature suggests, one can determine the rate constant from the negative
slop of the graph. Since this experimental data recorded by the spectrophotometer is proportional
to what can be recorded from the commercial instrument, these conclusions can be confirmed.
The rate may also be expressed by the following equation, which says that the change in
concentration of the iodine divided by the time required for that change determines the rate:
Rate= -∆ (I2) / ∆t (Slowinski 1973).
The spectrophotometer is able to measure this change in concentration by the color loss due to
the presence then the loss of iodine. It is important to note that here, iodine is the limiting
reactant in excess of acetone and HCl. Since the constant rate is being varied by the change in
concentrations, we can use the rate of two reactions to solve for the order, this will be discussed
in the “Results” section of this experiment.
A Zero Order Reaction
A zero order reaction has a constant rate, therefore the rate law is rate=k where k has
units of Ms-1, and the sum of the exponents in the rate law is equal to zero. A change in
3
temperature is the only thing can affect the rate of a zero-order reaction, and if concentration
from a zero order reaction is plotted against time then the result is a straight line with the slope
being the negative of the rate constant, -k.
This experiment utilizes the relationship between concentration of a reactant and time.
The integrated rate law, in which the equation expresses the concentration of a reactant as a
function of time. The integrated rate law of a zero-order reaction is given as:
[A]t = -kt + [A] 0
A First Order Reaction
A first-order reaction has a rate proportional to the concentration of one reactant and can
be expressed as rate=k [A] with units of sec-1. The integrated form of a first order reaction can be
expressed as ln [A]t = -kt + ln [A] 0 or ln [A]t [A] 0 = -kt or [A] = [A] 0 e-kt. A first order reaction
can be determined by plotting a graph of ln [A] vs. time and a straight line with a slope –k should
result (Slowinski 1973).
A Second Order Reaction
A second-order reaction has a rate proportional to the product of the concentrations of
two reactants or to the square of the concentration of one reactant. For example, rate = [A] 2 or
rate =k [A][B] (in other words, the sum of the exponents is two). The integrated rate law of a
second order reaction is:
1
1
=
+ 𝑘𝑡
[𝐴]
[𝐴]0
With the laboratory data, plots produced for [A] versus time, ln [A] versus time, and 1/[A] versus
time, in order to find the function with the linear fit. The iodination of acetone is convenient
because the color of the iodine reactant is readily observable and the change in concentration of
4
the iodine can be easily monitored by our spectrophotometer and the reaction proceeds quickly
enough for multiple trials but slowly enough to obtain measurements. Also, the reaction is zeroorder with respect to the iodine so the change in concentration is linear with time. This also
allows the initial rate of the reaction to be calculated as an average rate of the change in iodine
concentration over relatively long periods of time.
Experimental Procedure
A spectrophotometer that was built in the previous laboratory was assembled with a blue
diode as a light source for the initial step. LABview software was installed on a PC in order to
record data continuously throughout the reaction. Next, a dilute of iodine was prepared by setting
the voltmeter to read 3.5 Volts with about 4 mL of water, then Iodine was added drop wise until
the voltage was about 1.0 Volts. Next, the commercial spectrophotometer was used to record a
spectrum of the dilute from wavelengths 300nm to 800nm. On the commercial spectrum,
markers were placed at the location of the wavelengths of the blue diode. The blue diode was the
best choice because its wavelength is about 470nm, and the commercial iodine spectrum is
expected to produce highest absorbance at these wavelengths. Four trials were taken to obtain
data, and throughout these trials the volume of the mixture of iodine, acetone, HCl, and water
was 3.5 mL for each trial and the amount of each reactant was varied. A table of the
concentrations during each trial is shown below.
Iodine
Acetone
HCl
H2 0
Total Volume
Trial 1
1.0
1.0
1.0
0.5
3.5
Table 1: Reactant Concentrations
Trial 2
Trial 3
1.5
1.0
0.5
1.5
0.5
0.5
1.0
0.5
3.5
3.5
Trial 4
1.5
0.5
1.0
0.5
3.5
5
The reaction in this experiment can proceed very quickly. For this reason, a pipette was
used to transfer each reactant into a small beaker before each trial so that they could be combined
in the test tube and inserted into the spectrophotometer as quickly as possible. Multiple trials
were taken for each concentration combination. The Beer-Lambert law employed in the last
𝐼
laboratory (A= -log (𝐼 )) was used a second time to express the data in terms of absorbance.
0
Results and Discussion
Absorbance
Figure 1: Absorbance vs Wavelength Iodine
2
1.8
1.6
1.4
1.2
1
0.8
0.6
0.4
0.2
0
Blue: 470 nm
300
400
Yellow: 595 nm
500
600
700
800
Wavelength (nm)
The commercial spectrum suggests iodine will have the highest absorption at
wavelengths from about 540-580 nm. This wavelength range is closest to the wavelength omitted
by the yellow diode, suggesting this would have been a better choice. However, the dilution used
in the commercial spectrum was modified following some test reactions. In the future, another
commercial spectrum should be produced for the second dilution because the use of the blue
diode gave desirable results for the purpose of this experiment.
6
As mentioned in the introduction, a reaction is a zero order if when the concentration is
plotted versus the time the result is a straight line representing a constant rate. Therefore, the
iodination of acetone is a zero order reaction and this is demonstrated in experimental plots for
all concentration as well as the commercial spectrum. The slope of this line is the negative of the
rate law constant (-k). The slope and therefore the rate law constant for each reaction as well as
the commercial spectrum are shown in Table 1. Table 2 shows the orders of the reactants
calculated using the equation: Rate x/Rate y = (k [acetone]a[iodine]I[HCl]h)x ÷ (k
[acetone]a[iodine]I[HCl]h)y, an example is shown in Sample Calculation IV. Solving for a
resulted in a value of 0.69, which rounds to one. Solving for h resulted in a value of 0.71 which
also rounds to one. Adding the three exponents result in a pseudo-order of two.
k (1/Ms)
Trial 1
0.0011
Table 2. Rate Law Constants (k)
Trial 2
Trial 3
Trial 4
0.0008
0.0017
0.0018
Average
0.00135
Data from individual plots of absorbance versus time was extrapolated to plot absorbance
versus time for each reaction on one graph for the start and finish of each. The combined data is
shown below, and an example plot can be found in the appendix that shows the absorbance
before the start of the reaction as well as stabilization after the concentration of iodine reached
zero. As can be seen in Figure 2, a linear function with a negative slope resulted from each
reaction all of which varied in concentration of the reactants. Therefore, the reaction can be
considered zero order with respect to iodine from this visual representation. It can also be said
from this data that each reaction lasted between 100-120 seconds. Table 1 shows the rate
constant from each reaction taken from the negative of the slope; these values indicate the rate is
constant. Mathematically, the reaction is determined to be zero order with respect to iodine by
7
finding the rates of two reactions, the corrected concentrations of iodine, and solving for the
coefficient; this is shown in Appendix A. Sample calculations.
In “Reaction 2” acetone and HCl were not present in excess. For this reason, the reaction
took about 180 seconds longer and values for k and the rate was not useful for this experiment.
In the reactions where acetone and HCl are in excess, their concentrations remain mostly
constant throughout the reaction.
Figure 2: Absorbance Vs. Time
0.3
Absorbance
0.25
0.2
0.15
0.1
0.05
0
0
20
40
60
80
100
120
Time (s)
Reaction 1
Reaction 2
Reaction 3
Reaction 4
Figure 3 shows the natural log of the absorbance for each of the four reaction versus
time. The non-linearity of the graph further confirms that the reaction is zero order with respect
to iodine. If the reaction were first order, there would be a linear trend for this plot.
8
Figure 3: ln(Absorbance) Vs. Time
6
5
ln(A)
4
3
2
1
0
0
20
40
60
80
100
Time (s)
ln(A): T1
ln(A): T2
ln(A):T3
ln(A): T4
A plot of the inverse of the absorbance versus time also failed to show a linear trend. This
plot, Figure 4, represents what would be a second order reaction if the trend wasn’t exponential.
This graph tells us that during the timed reaction the inverse of the absorbance increases
exponentially therefore the absorbance must be decreasing.
Figure 4: 1/Absorbance Vs. Time
700
600
1/A
500
400
300
200
100
0
0
20
40
60
80
100
120
140
Time (s)
1/a: T1
1/A: T3
1/A: T4
Error
9
A potentially significant source of error in this experiment is room temperature. As
mentioned in the introduction, temperature is known to affect the rate of the reaction. Since this
experiment occurred over multiple laboratory periods, the room temperature and temperature of
the solutions were likely different each time. In addition to temperature consistency, it would be
beneficial to run more trials and utilize the concentration ratio that produces accurate and precise
results. Other sources of error to be consider include the spectrometer’s voltage readings which
may cause noise in the recorded data.
Conclusion
This reaction is experimentally beneficial for two reasons. First, the color of the iodine
allows the concentration change to be readily followed; second, the reaction turns out to be zero
order. Therefore, the rate of the reaction is independent of the iodine and the rate can be studied
by making the iodine the limiting reactant in an excess of acetone and HCl. The experimental
data supports these conclusions by showing a linear trend for an absorbance versus time plot as
shown in Figure 2, and also by producing similar results for each reaction with varied
concentrations of each reactant. The deductions from these graphical representations of the
reactions can be verified by the kinetics equations. Through these methods, the overall order of
the reaction was found to be zero, and the orders of acetone, HCl, and iodine were found to be
one, one, and zero. Acetone and HCl may be omitted to give the rate law, rate = k [I2]I or rate =
k. The average value for the rate constant which was found for each trial was found to be
0.00135 1/Ms.
10
Appendix
A. Sample Calculations
I)
Calculation of the corrected concentration (Acetone, Trial 4 shown)
M1V1=M2V2
Where M1 is the molarity of the given substance, V1 is the volume used of that
substance, M2 and V2 are the total volume and molarity used
0.5 mL (V1) of a 6 M acetone solution (M1). M2 is the variable in question and V2 is
3.5 mL (total volume used)
M2= [(6M)*(0.5mL)] / 3.5mL
II)
Calculation of The Rate (Trial 4)
Using the equation: Rate= [I2]/s x 105
The concentration of Iodine in the sample (0.05 M) and the time it took for the
reaction to go to completion (t = 116s)
Rate= [(0.05)/(120)] = 0.00041
III)
Determination of the Order
Rate 4/Rate 3 = (k [acetone]a[iodine]I[HCl]h)2 ÷ (k [acetone]a[iodine]I[HCl]h)3
Where Rate 2 = 1.84729E-5 and Rate 4 = 1.23153E-5 and
[Iodine]2= 0.002142857 and [Iodine]3= 0.001428571
(1.84729E-5) ÷ (1.23153E-5) = [0.0021] I / [0.0014] I
1.50= 1.4997I
I= log (1.5) / log (1.4997)
I=1
11
B. Experimental Data
I)
Corrected Concentrations of Reactants and Reaction time
T1 (M1)
Iodine
0.001429
Acetone 1.714286
HCl
0.285714
Time (s)
116
II)
Corrected Concentrations
T2 (M1)
T3 (M1)
T4 (M1)
0.002143 0.001429 0.002143
0.857143 2.571429 0.857143
0.142857 0.142857 0.285714
280
103
116
Rates of the Reactions
Rates (M/s)
Rate T1
Rate T2
Rate T3
Rate T4
Average(M/s)
1.23153E-05 7.65306E-06 1.387E-05 1.8473E-05
9.9989E-06
III)
Intensity vs Time: Trial 4
Intensity Vs. Time: Trial 4
3
2.5
Intensity (V)
2
1.5
1
0.5
0
0
20
40
60
80
100
120
140
160
Time (s)
IV)
Absorbance vs Time: Trial 4
12
Absorbance
Absorbance vs. Time: Trial 4
0.5
0.45
0.4
0.35
0.3
0.25
0.2
0.15
0.1
0.05
0
0
50
100
150
200
Time
Absorbance vs Time: Trial 2
Absorbance vs Time: Trial 2
0.35
0.3
0.25
Absorbance
V)
0.2
0.15
0.1
0.05
0
0
50
100
150
200
250
300
350
Time (s)
13
References
I.
"The Rate Law." - Chemwiki. N.p., n.d. Web. 29 Oct. 2014.
II.
Szwarc, M. "The Dissociation Energy of the C-N Bond in Benzylamine."Proceedings
of the Royal Society of London. Series A Mathematical and Physical Sciences.Vol.
198, No. 1053 (1949): 285-92. Web. 29 Oct. 2014.
III.
Meyer, Earl N., and Flourence F. Lask. "Rate and Activation Energy of the Iodination
of Acetone." Journal of the Chemical Society (2010): 1-5.Minnesota State University
Moorhead. Web. 2 Nov. 2014.
IV.
Slowinski, Emil J., Wayne C. Wolsey, and William L. Masterton. "Experiment 20:
Rate of Chemical Reactions." Chemical Principles in the Laboratory. Philadelphia:
Saunders, 1973. 149, 150-52. Print.
V.
Tinoco, Ignacio, Kenneth Sauer, and James C. Wang. Physical Chemistry: Principles
and Applications in Biological Sciences. Englewood Cliffs, NJ: Prentice-Hall, 1978.
Print. Fifth.
14
View publication stats