Acetone Iodination Kinetics: Determining Reaction Order & Rate

This report investigates the kinetics of acetone iodination, a composite reaction, through experimental analysis. By varying reactant concentrations (acetone and HCl), the reaction's progress was monitored and timed, allowing for the determination of reaction orders and rate constants. Iodine served as an indicator due to its color change during the reaction. The experimental procedure involved mixing hydrochloric acid, acetone, and iodine solutions, recording the time taken for the iodine color to disappear. Concentration values were calculated, and reaction rates were determined at various time intervals. Graphical plots were generated to analyze the relationship between reaction rate and reactant concentrations, ultimately revealing the reaction orders for HCl and acetone. The study concludes that the reactions of HCl and acetone are approximately first order, while the reaction of iodine is of zero order, with the pseudo rate values found to be far less than zero. Desklib provides access to this and other solved assignments.

TITLE: THE KINETICS OF THE IODINATION OF ACETONE, A COMPOSITE
REACTION
ABSTRACT
This report aims at determining the kinematics of iodination of acetone. This is done by
monitoring the rate of chemical reaction and keenly observing its progress as the experimental
data is recorded. Through variation of the concentration of the reactants in every set of the
mixtures as time of chemical reaction is recorded, the reaction order and the rate constant of the
reaction are determined for both the acetone and HCL reactions. From these two parameters, the
rate of chemical reaction can be computed. Iodine, due to its colour, is used as an indicator in
this experiment to show the times of the reaction.
INTRODUCTION
Time is a very key tool in determining how fast or slow a reaction is(Bell & Jones, 2014). A
reaction that takes a lesser time for the reactants to disappear to form product and vice versa, at
equilibrium, is said to have a faster rate of reaction while that which takes more time for the
reactants to disappear is said to have a slower rate of chemical reaction(Brezonik, 2018). Rate of
chemical reaction measures how fast the reactants disappear to form a product. Concentration of
both the reactants and products determines the rate of the reaction per unit time - which can be
in seconds or minutes(Gold, 2016, p. xx). Generally, the rate of a reaction (R) is given by the rate
of disappearance of the reactants i.e.
rate of reaction , R= ∆(Reactants)
∆ t
Alternatively, this rate can also be given as the rate of formation of the product(Dawson & Key,
2016) i.e.
rate of reaction , R= ∆(Products)
∆ t
The change in the rate of the reactants is always negative. This is because, as the reaction
proceeds, the concentration of the reactants reduces(Michael, 2015). On the other hand, the change
in the rate of the product is positive as shown in the above equation since the concentration of the
product increases as the reaction proceeds. This is due to formation of more products on the
product side(Foster, 2015). Dark red solid iodine changes to yellow in aqueous state, in the

presence of an acid when reacted with acetone(Ninham, 2016). As iodine is consumed in the
formation of hydrogen iodide and iodoacetone on the product side, the yellow colour fades. H+
acts as a catalyst for this reaction(Liang, Mi, Wang, Wang, & Zhang, 2004). The chemical equation
below describes the process.
( C H3 ) 2 C=O+ I2 →C H3 ( C H2 I ) C=O+ HI
The equation describing the rate law is therefore expressed as;
Rate=−∆(I2 )
∆ t =k (I 2)x ¿ ¿
METHOD
10ml of Hydrochloric acid and acetone of 10ml were measured and poured inside flask.
Thereafter, iodine solution (of 2ml capacity) was placed into the flask containing acetone and
HCL solutions and stop watch started immediately. The time when the yellow color of iodine
disappeared was then recorded as in results. The above procedure was repeated at various values
of time with the total volume of the mixture kept at 30 ml at every instant.
RESULTS
Table 1.1: Experimental Results
Volume of
Acetone\ml Volume of HCL\ml Volume of Iodine\ml Time\s
10 10 2 57
8 10 2 66
6 10 2 99
4 10 2 139
2 10 2 287
10 8 2 71
10 6 2 94
10 4 2 134
10 2 2 271
10 6 1 49
10 6 4 190

Table 1.2: Concentration Values
Calculated Values of Concentration
Time, s 57 66 99 139 287 71 94 134 271 49 190
Acetone 0.667 0.533 0.400 0.267 0.133 0.667 0.667 0.667 0.66
7
0.667 0.667
HCL 0.333 0.333 0.333 0.333 0.333 0.267 0.200 0.133 0.06
7
0.200 0.200
Iodine 3.333
x10-4
3.333
x10-4
3.333
x10-4
3.333
x10-4
3.333
x10-4
3.333
x10-4
3.333
x10-4
3.333
x10-4
3.33
3
x10-4
1.667
x10-4
6.667
x10-4
Table 1.3: Reaction rates at various intervals
Time, s Reaction Rates
M/s Log J HCL Log (HCL) Acetone Log(Acetone) Iodine
Log
(Iodine)
57 8.771930 x 10-5
-4.0569 0.333 -0.477556 0.667 -0.175874166 3.333 x10-4 -3.47713
66 7.575758 x 10-5
-4.1206 0.333 -0.477556 0.533 -0.273272791 3.333 x10-4 -3.47713
99 5.050506 x 10-5
-4.2967 0.333 -0.477556 0.4 -0.397940009 3.333 x10-4 -3.47713
139 3.597122 x 10-5
-4.444 0.333 -0.477556 0.267 -0.573488739 3.333 x10-4 -3.47713
287 1.742160 x 10-5
-4.7589 0.333 -0.477556 0.133 -0.876148359 3.333 x10-4 -3.47713
71 7.042254 x 10-5
-4.1523 0.267 -0.573489 0.667 -0.175874166 3.333 x10-4 -3.47713
94 5.319149 x 10-5
-4.2742 0.2 -0.69897 0.667 -0.175874166 3.333 x10-4 -3.47713
134 3.731343 x 10-5
-4.4281 0.133 -0.876148 0.667 -0.175874166 3.333 x10-4 -3.47713
271 1.845018 x 10-5
-4.734 0.067 -1.173925 0.667 -0.175874166 3.333 x10-4 -3.47713
49 1.020408 x 10-5
-4.9912 0.2 -0.69897 0.667 -0.175874166 1.667 x10-4 -3.77806
190 2.631579 x 10-5
-4.5798 0.2 -0.69897 0.667 -0.175874166 6.667 x10-4 -3.17607

Graphical Results
Fig. 1.1: Plot of log (J) against log (Concentration of acetone)
-1 -0.9 -0.8 -0.7 -0.6 -0.5 -0.4 -0.3 -0.2 -0.1 0
-6
-5
-4
-3
-2
-1
0
f(x) = 0.309529392312935 x − 4.34538058781984
R² = 0.0649571489474092
Log J versus Log acetone
Fig 1.2: Plot of log (J) against log (Concentration of HCL)
-1.4 -1.2 -1 -0.8 -0.6 -0.4 -0.2 0
-6
-5
-4
-3
-2
-1
0
f(x) = 0.59363367388678 x − 4.0560918967157
R² = 0.298783823467274
Log J versus log HCL

Fig 1.3: Plot of log (J) against log (Concentration of Iodine)
-7 -6.5 -6 -5.5 -5 -4.5 -4 -3.5 -3
-6
-5
-4
-3
-2
-1
0
f(x) = NaN x + NaN
R² = 0 Log J versus Log iodine

DISCUSSION
The corrected concentration values of acetone and HCL (M 2) were obtained at the various time
intervals from the relation,
M 1 V 1=M 2 V 2
∴ M2= M 1 V 1
V 2
Where the variables V 1 and V 2 represent volume of the specific substance used and the total
volume of the solution respectively. The concentration M 1 for acetone and HCL used were 2M and
1M respectively. For instance, at the 57th second interval, the value of concentration of acetone is
given by;
¿ M1 V 1
V 2
=2 x 10
30 =0.666667 M
Other values are as recorded in table 1.2 of results.
The rates of reaction at various time intervals were then obtained from the formula;
Rate= Concentration of Iodine
Time taken
For instance, at 66th second, the rate is given as;
¿ 0.005
66 =7.57576 x 10−5 M /s
The rates at other time instants were obtained as illustrated above and results recorded in table
1.3. Based on the formula given on the lab manual, the logarithms of the rates of reaction (J)
were computed and that of HCL acid and acetone calculated too as therein results. Three separate
plots were then made to determine the order of reaction.

These plots are:
i). Graphical plot of log (J) against log (Concentration of HCL)
ii). Graphical plot of log (J) against log (Concentration of acetone)
iii). Graphical plot of log (J) against log (Iodine)
The graphical plot of log (J) against log (Concentration of HCL) yielded a linear graph passing
through log J axis and with a regression equation of
log J=−4.0561+0.5936 log [HCL]
At an R2 levels of 0.2988.
Comparing this equation with the one provided on the manual,
log J=log k' +α log [ A ]
The value of α=0.5936 ≈ 1 and log k' =−4.0561 Thus the reaction is of order one and the pseudo
rate constant is k' =10−4.0561=8.7882 x 10−5
Similarly, the graphical plot of log (J) versus log (Concentration of acetone) shown in figure 1.1
yielded a linear graph passing through log J axis and with a regression equation of
log J=−4.3454+0.3095 log [ Acetone ]
And at R2 levels of 0.065
Comparing this equation with the one provided on the manual,
log J=log k' +α log [ A ]
The value of α=0.3095 ≈ 1and log k' =−4.3454 .Thus the reaction is of order one and the pseudo
rate constant is k'=10−4.3454=4.5143 x 10−5.

Finally, the normalized graphical plot of log (J) versus log (Concentration of Iodine) shown in
figure 1.3 yielded a vertical line passing through log J axis at -8 and with a regression equation
of
log J=−8+0 log [Iodine]
And at R2 levels of 0
Comparing this equation with the one provided on the manual,
log J=log k' +α log [ A ]
The value of α =0and log k' =−8 .Thus the Iodine reaction is of order zero and the pseudo rate
constant is k'=10−8=−0.00000008. The overall reaction is however of order zero.
The accuracy of the data obtained is however questionable due to possible experimental sources
of errors including; fluctuations in the room temperatures which affect the rate of reaction
undesirably. However, this could be minimized by taking multiple readings at the different room
temperatures and averaging the results.
CONCLUSION
In a nutshell, the apparatus were set and the experiment carried out as therein procedure. The
obtained results were then analyzed and graphical plots obtained to characterize the reactions.
The plots obtained conformed to their expected theoretical profile. All of them were linear
graphs not passing through the origin. The regression equations of the plots gave the order of the
reaction and pseudo rate constant for HCL, acetone and iodine. It was established that the
reactions of HCl and acetone were both approximately of first order while that of iodine found to
be of order zero. The pseudo rate values were found to be far much less than zero. Kinetic
equations could be used to verify the graphical output parameters. In this respect, the
experimental objectives were justified.

REFERENCES
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543-551.
Foster, B. L. (2015). Principles of laboratory safety management in academia. Chemical Health
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ChemPhysChem, 17(16), 2567-2571.