The overall reaction is as follows, with Ac representing acetone

26 Jun The overall reaction is as follows, with Ac representing acetone

Question
6

Reaction Kinetics of Iodine with Ketone

Introduction: In this experiment we investigate the kinetics of the iodination of
acetone. The overall reaction is as follows, with Ac representing acetone
−
Ac + I3 −→ AcI + 2I − + H +

(6.1)

Some general information about reaction 6.1 is known. The reaction of acetone
with iodine in dilute aqueous solution is very slow in neutral solutions, but becomes
appreciably faster in the presence of a strong acid. Neutral salts of strong acids do
not increase the rate of reaction. A weak acid (such as acetic) increases the rate of
iodination less than does a strong mineral acid such as HCl. This suggests that H+
is involved in the kinetics; hence we propose a rate law of the form
d −
−
[I ] = k[H + ]a [Ac]b [I3 ]c
dt 3
where k(T ), a, b, and c are to be determined experimentally.
−

6.1

(6.2)

Theory

The kinetic technique of “swamping” or “isolation” can be used to simplify interpretation of the data. Suppose a reaction solution is prepared in which the initial
concentration of I− is much smaller than the initial concentrations of H+ and Ac.
3
During the course of the reaction, I− and Ac are consumed, and H+ is formed ac3
cording to equation 6.1. Because of the large initial concentrations of H+ and Ac,
their concentrations do not change appreciably, while the concentration of I− does
3
change considerably. In other words, [H + ] and [Ac] are effectively constant during
the reaction. In this case, equation 6.2 can be simplified to
d −
−
[I3 ] = k [I3 ]c
dt

(6.3)

k = k[H + ]a [Ac]b
0
0

(6.4)

−
where

This equation can then be integrated directly for a particular choice of c, giving
the following possibilities:
c=0

−
−
[I3 ]0 − [I3 ] = k t

90

(6.5)

−
−
ln [I3 ]0 − ln [I3 ] = k t

c=1

(6.6)

1
1
=kt
− −
−
[I3 ] [I3 ]0

(6.7)

c=2

−
Plots of various functions of [I3 ] vs. t will indicate which of the above equations
is valid, thus establishing c.

Suppose that the reaction is now carried out with a different initial concentration
of H+ (still large enough to satisfy the swamping requirements) and the same concentration of Ac. The data from this kinetics run will confirm the evaluation of c
from the first results, but a different value for the apparent rate constant will be
obtained. The values for the rate constant for the two different initial solutions are
denoted by k1 and k2 . They are related by
k[H + ]a [Ac]b
k1
0,1
0
=
=
+ ]a [Ac]b
k2
k[H 0,2
0

[H + ]0,1
[H + ]0,2

a

(6.8)

Here [H + ]0,1 is the initial concentration of H+ in solution 1, etc. Hence, the two
values k1 and k2 plus the known initial concentrations of H+ allow the determination
of a. Similarly, variation of [Ac]0 allows the determination of b. It is necessary to
vary the initial concentrations by at least a factor of two to adequately characterize
the exponents.
Of the species mentioned thus far, only I− (aq) absorbs visible light, so it is con3
venient to follow reaction progress by using a spectrophotometer to monitor the
disappearance of I− . There is, however, a minor complication: I− is involved in the
3
3
following equilibrium:
−
I3

I − + I2

(6.9)

with Keq ∼ 1.4×10−3 (Ref. 2). I2 both absorbs visible light an acts as an effective
iodinating agent. Hence reaction rate and the amount of light absorbed by a given
solution of I− depend not only on the concentration of I− , but on the concentration of
3
3
I2 as well. Using Beer’s Law (A = c), the total absorbance of a solution containing
both I− and I2 may be expressed as the sum of the individual absorbances:
3
A=

I2

[I2 ] +
91

−
I3

−
[I3 ]

(6.10)

Figure 6.1: Visible absorption spectrum of I2 and I−
3

−
The molar absorptivities, I2 and I3 , are functions of the wavelength of light
being absorbed, as indicated in Figure 6.1. While in general there is no relation
−
between I2 and I3 , note in particular the relationship which exists at λ = 565 nm.
This is called an isosbestic point. Since the two molar absorptivites are equal at
565 nm, the absorbance of a solution of I2 and I− is
3

A=

−
I3

[I2 ] +

−
I3

−
[I3 ] =

−
I3

−
([I2 ] + [I3 ])

(6.11)

That is, the absorbance at this particular wavelength is proportional to the sum of
the concentrations of I2 and I− . Consequently, we can monitor the total concentration
3
of the iodinating agents (assumed to be I− for simplicity of notation) in a solution
3
by following the absorption at 565 nm, provided has been determined for this
wavelength.
The rates constants in the above analysis are constant with respect to concentration, but they vary with temperature, as indicated by the Arrhenius equation
k = Ae−Ea /RT

(6.12)

where A is constant related to the geometry of the molecules involved in the
reaction and Ea is the activation energy. Both values can be determined if k is
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measured at several temperatures. Data analysis typically includes plotting ln(k) vs.
1/T , so that the slope and intercept readily provide the desired values.

6.2

Experimental

The following solutions should be carefully prepared using volumetric glassware:
1. 25 mL of 0.047 M I− (0.047 M in I2 and 0.47 M in KI) note 10-fold excess
3
of KI
2. 100 mL containing 7.0 mL of reagent grade acetone (calculate resulting molarity)
3. 100 mL of 1.0 M HCl
Keep solutions stoppered to avoid evaporation. All waste should go in the marked
container in the hood.
A minimum of five kinetic runs is necessary for determination of the rate law and
the temperature dependence. The rate law is typically determined at 25o C. Select
the set of reaction conditions with the smallest absorbance change for the additional
kinetic runs at 30o C and 35o C.
Possible reaction mixture:
1. 10.0 mL acetone solution,
2. 5.0 mL 1.0 M HCl,
3. 8.0 mL water,
to which 2.0 mL of I− solution will be added at t = 0 Record the delay before the
3
initial spectrometer reading using a timer. (Note the initial concentrations; to what
extent is the swamping satisfied?)
Possible procedure:
1. Carefully make up a reaction solution of the appropriate amounts of H+ and
Ac in a 25 mL volumetric flask, and add any water necessary. Equilibrate
this solution and the stoppered I− solution in a water bath maintained at the
3
desired temperature.
2. Use the instructions given in the appendix at the end of this lab to prepare the
spectrometer for data collection.
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When everything is ready, transfer the appropriate amount of the I− solution
3
to the reaction tube, and start a timer. Mix the solution vigorously. Transfer
a small amount of the reaction mixture to the cuvette to rinse it out, then
fill the cuvette with a fresh sample of the reaction mixture, and immediately
begin data acquisition with the Execute button. Stop the timer when data
acquisition begins.
3. To determine the molar absortivity of the iodine solution, use absorbance
values for at least three solutions of known concentration. Standard solutions
should have A565nm between 0.1 and 1.5. Data collection is accomplished using
the same general procedure as for the kinetic measurements but the delay time,
integration time and end time can be adjusted.

6.3

Data Analysis

1. Determine the molar absortivity of the I− /I2 mixture by constructing a Beer’s
3
Law plot with the data from your standards.
2. Use your

−
to convert A to [I3 ] for each run.

−
3. Determine the reaction order with respect to [I3 ] using your best data set and
plotting integrated equations for c = 0, 1, or 2. Running regression analysis
and plotting residuals may help.

4. Obtain k for each run.
5. Determine k, a, b, and c for the reaction.
6. Obtain the Arrhenius parameters by plotting ln(k) vs. 1/T .

6.4

Discussion

1. There are two other isosbestic points in Figure 6.1. Why is 565 nm chosen for
analysis?
2. A proposed mechanism for this reaction is given in Figure 6.2. Using the
notation presented there, derive an expression for the overall rate law. What
assumptions must be made for it to match your experimentally determined rate
law?
3. Does the activation energy seem consistent with the proposed mechanism?
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Figure 6.2: Proposed mechanism for acid-catalyzed iodination of acetone

4. How do your data compare with literature values?
5. In this experiment, we have made the assumption that I− and I2 have equal
3
reactivity for ketone iodination. Is this valid?
6. Is the disappearance of I− a good indicator of the overall reaction rate?
3

95

References
1. Birk, J. P.; Walters, D. L. J. Chem. Ed. 1992, 69, 585.
2. Jones, G.; Kaplan, B. B. J. Am. Chem. Soc. 1928, 50, 1845.
3. Rice, F. O.; Kilpatrick, M., Jr. J. Am. Chem. Soc. 1923, 45, 1401.
4. McQuarrie, D. A.; Simon, J. D. Physical Chemistry; University Science Books:
Sausalito, 1997.

6.5

Appendix

HP-8453 Kinetics Measurement
1. Start-up procedure
• Power on computer (if off), spectrometer, and temperature control box
• On desktop, open the Instrument 1 online program
• Select Chem125 as a username and click enter
• After the program opens, power on the printer
• Note the lamp icons in the lower left corner and check that the UV lamp is
off and the visible lamp is on (visible lamp represented by the smaller bulb
icon), turn either lamp on or off by clicking the bulb icon and selecting
lamp off or lamp on
• Note the direction of the light path on the spectrometer and remove the lid
above sample holder. Be sure that the lever on the sample holder is up to
insert or remove a sample cell and down when taking measurements. The
lid can be replaced once sample cell is inserted to stabilize temperature.
• Create a folder with your name to save all data to; this folder can be saved
to the Chem125 folder.
2. Temperature Control
• The current temperature is displayed on the screen that reads Cell Temp
= ##
• Move between screens using the [Up] and [Down] keys
• Find the Set Temp screen and hit [Enter]. This should initiate a flashing
number
96

• Enter the desired temperature using keypad and select Enter
• Return to Cell Temp display to monitor the current temperature reading
3. Take I− spectra at known concentrations
3
• Check that mode is set to standard (top, center pull down menu)
• Run blank spectrum by selecting blank (near the lamp icons), close the
window showing your blank spectrum—make sure to use the same solvent
your sample is in
• Note absorbance value at specific wavelength by selecting fixed wavelengths
from the scroll down menu under the task box on the left side of the screen,
then select setup and enter a wavelength in one of the use wavelength(s)
boxes. Select OK and check the Sample/Result Table under your spectrum for the absorbance at the indicated wavelength
• To save spectrum, click on desired trace then select file | save selected
spectrum as | CSV format, then save to your personal folder
4. Time dependent spectra
• Change mode to Kinetics
• Open set-up under the time and calculations box on the left side of the
screen
• Enter use wavelengths as the wavelength of which you would like to
record the absorbance over time
• Enter a run time of 1000 s, a start time of 0 s, and a cycle time of 30 s,
then select OK
• Run blank spectrum by selecting run blank under the lamp icons—make
sure to use the same solvent as your sample
• Run scan by selecting time based measurement located in the bottom
left corner (selecting sample will run a single spectrum rather than start
your time scan)
• Save your file with a unique name (do not overwrite other data) as name.KD
when prompted; save each run with a different name
• Select start in the bottom left corner
• Save by clicking on data point to select time trace, then select file |
save selected spectrum as | CSV format, then save to your personal
folder
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5. Shutdown procedure
• Turn off visible lamp and close program
• Turn off power to temperature control, spectrometer, and printer
• Leave computer on

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