4.
KINETICS OF IODINATION OF ACETONE
Aim
To determine the rate constant and activation energy of the reaction; iodination of acetone
Principle
The rate of a chemical reaction depends on several factors: the nature of the reaction, the
concentrations of the reactants, the temperature, and the presence of a possible catalyst. In
Part One: Determine the rate law for a reaction by changing some of the above variables and
measuring the rate of the reaction.
Part Two: Find out the relation between the rate constant and temperature to discover the
activation energy for this reaction.
In this experiment we will study the kinetics of the reaction between iodine and acetone:
The rate of this reaction is found to depend on the concentration of the hydrogen ion (acid,
HCl) as well as the concentrations of the reactants (acetone and iodine).
The rate law for this reaction is;
Rate = k[acetone]m[H+]n [I2]p
where k is the rate constant for the reaction and m, n, and p are the orders of the reaction with
respect to acetone, hydrogen ions (acid), and iodine, respectively.
The rate of the reaction can also be expressed as the change in the concentration of a reactant
divided by the time interval:
−𝒅[𝑰𝟐 ]
Rate =
𝒅𝒕
The iodination of acetone is easily investigated because iodine (I2) has a deep yellow/brown
colour. As the acetone is iodinated and the iodine converted to the iodide anion, this colour
will disappear, allowing the rate of the reaction to be easily monitored.
We can study the rate of this reaction by simply making I2 the limiting reactant in a large
excess of acetone and H+ ion. By measuring the time required for the initial concentration of
iodine (I2) to be used up completely, the rate of the reaction can be determined by the
equation;
−𝒅[𝑰𝟐 ] −([𝑰𝟐 ]𝒇𝒊𝒏𝒂𝒍 −[[𝑰𝟐 ]𝒊𝒏𝒊𝒕𝒊𝒂𝒍 ) −(𝟎−[𝑰𝟐 ]𝒊𝒏𝒊𝒕𝒊𝒂𝒍 ) [𝑰𝟐 ]𝒊𝒏𝒊𝒕𝒊𝒂𝒍
Rate = = = =
𝒅𝒕 (𝒕𝒇𝒊𝒏𝒂𝒍−𝒕𝒊𝒏𝒊𝒕𝒊𝒂𝒍 ) (𝒕𝒇𝒊𝒏𝒂𝒍−𝟎) 𝒕𝒇𝒊𝒏𝒂𝒍
[𝑰𝟐 ]
Rate =
𝒕𝒊𝒎𝒆
From the rate information we can determine the orders with respect to acetone (m), acid (n)
and iodine (p) by varying the amounts of reactants and measuring the effect on the rate. Once
the orders of reaction are known, we will be able to calculate the rate constant, k. In Part One
of this experiment, you will determine the rates of reactions, the orders of the reactants, and
finally the rate constant at room temperature.
In the second part of this experiment, you will study the rate of the reaction at different
temperatures to find its activation energy, Ea. The temperature at which the reaction occurs
influences the rate of the reaction. An increase in temperature increases the rate.
As with concentration, there is a quantitative relation between reaction rate and temperature,
but here the relation is somewhat more complicated. This relation is based on the idea that to
react, the reactant species must have a certain minimum amount of energy present (and the
correct geometry, if appropriate) at the time the reactants collide in the reaction step.
This amount of energy, which is typically furnished by the kinetic energy of the species
present, is called the activation energy (Ea, also known as the energy of activation) for the
reaction. The formula (Arrhenius equation) relating the rate constant k to absolute Kelvin
temperature T and Ea is:
−𝑬𝒂
𝒍𝒏𝒌 = + 𝒍𝒏𝑨
𝑹𝑻
In this equation, R is the gas constant (8.3145 J/mole K), and natural logarithms (ln) need
to be used (do not use base 10 logs!) The quantity A is referred to as the collision
frequency/Arrhenius factor and A can be used to determine the fraction of molecules
present with sufficient energy and geometry to become products at a given instant in time.
By measuring k at different temperatures, we can graphically determine the activation energy
for a reaction. In Part Two of this experiment, you will determine the effect of temperature on
rate and calculate the activation energy for the reaction.
Procedure:
Measure out the appropriate quantities of 1.0 M HCl, 4.0 M acetone and water using a 10 mL
graduated cylinder and place them in a conical flask.
Now measure out the appropriate amount of 0.0050 M iodine in a 10.00 mL graduated
cylinder. Start a timer (stopwatch) as you add the iodine to the 250 mL flask with the other
chemicals. Swirl the stoppered flask (which helps to prevent acetone evaporation) until the
yellow colour disappears, then halt the timer. It may help to place the flask on a white piece
of paper to help to understand when the colour disappears. Record the time elapsed in
seconds. Repeat this reaction mixture until two trials are within 20 seconds of each other.
Repeat this process for each of the four trials listed in the table below.
Trial 1
HCl (mL) Acetone (mL) Water (mL) Iodine (mL) Total volume (mL)
5 5 10 5 25
Time in seconds for yellow colour to disappear (I):
Time in seconds for yellow colour to disappear (II):
Average time:
Trial 2:
HCl (mL) Acetone (mL) Water (mL) Iodine (mL) Total volume (mL)
5 5 5 10 25
Time in seconds for yellow colour to disappear (I):
Time in seconds for yellow colour to disappear (II):
Average time:
Trial 3:
HCl (mL) Acetone (mL) Water (mL) Iodine (mL) Total volume (mL)
5 10 5 5 25
Time in seconds for yellow colour to disappear (I):
Time in seconds for yellow colour to disappear (II):
Average time:
Trial 4:
HCl (mL) Acetone (mL) Water (mL) Iodine (mL) Total volume (mL)
10 5 5 5 25
Time in seconds for yellow colour to disappear (I):
Time in seconds for yellow colour to disappear (II):
Average time:
Effect of temperature on reaction rate:
As before, add all of the chemicals except iodine to the conical flask (keep the mouth of the
flask closed). Using water bath, get the solution to a desired temperature before adding the
iodine. Record the temperature, then add the iodine to the flask, starting a stopwatch and
measuring how long the reaction takes to turn the solution clear. Time should be recorded in
seconds. The iodine does not need to be at the same temperature as the solution in the
Erlenmeyer flask.
Record one trial at room temperature and other three at higher temperatures (40 °C, 50
°C, 60 °C)
HCl (mL) Acetone (mL) Water (mL) Iodine (mL) Total volume (mL)
5 5 10 5 25
Temperature: ----------°C Time: ----------s
Temperature: ----------°C Time: ----------s
Temperature: ----------°C Time: ----------s
Temperature: ----------°C Time: ----------s
Determine the rate and rate constants:
We need to find out the concentrations of HCl, acetone and water in order to calculate rate
and rate constants.
The concentrations of bulk reactants are 1.0 M HCl, 4.0 M acetone, and 0.0050 M
iodine which were placed (with water) in conical flask. The final volume was always 25.00
mL.
Mixing chemicals dilutes the concentrations from the "bulk" value to a smaller value.
We can determine these diluted values using: M1V1 = M2V2
For example;
Let M1 = initial (undiluted) concentration of iodine (0.0050 M), V1 = 5.00 mL (of undiluted
iodine added to the mixture), and V2 = 25.00 mL (the total volume of the diluted solution
once HCl, acetone and water are added). Solving for M2, the concentration of iodine in the
diluted solution, one gets: M2 = 0.0050 M * 5.00 mL / 25.00 mL = 0.0010 M, which is the
concentration of iodine used in the reaction in trial 1.
Similarly, calculate the concentrations of HCl, acetone and iodine for all the four trials.
Conc. of HCl (M) Conc. of acetone (M) Conc. of iodine (M)
Trial 1
Trial 2
Trial 3
Trial 4
If you know concentration of iodine ([I2]), rate can be easily calculated using;
[𝑰𝟐 ]
Rate =
𝒕𝒊𝒎𝒆
Calculate the rate for all four trails
[I2] (M) Average time (s) Rate (M/s)
Trial 1
Trial 2
Trial 3
Trial 4
In order to calculate rate constant, we need to know the values of m, n and p. It was found
experimentally that, m=2, n=2 and p=0.
Therefore, rate constant can be calculated for all trials using;
Rate = k[acetone]m[H+]n [I2]p
Acetone (M) HCl (M) Iodine (M) Rate Rate
constant, k
(molL-1s-1)
Trial 1
Trial 2
Trial 3
Trial 4
Average k =
Repeat the same procedure to calculate k at different temperatures (convert temperature from
°C to K)
Temperature Acetone HCl (M) Iodine (M) Rate Rate
(K) (M) constant, k
(molL-1s-1)
Trial 1
Trial 2
Trail 3
Trial 4
To find out the activation energy, plot a graph between lnk and 1/T. Activation energy and
Arrhenius parameter can be calculated from the slope and y intercept respectively.
Slope = -Ea/R
Ea = -Slope×R
ln A = y intercept
A = e y intercept
Result:
Rate constant at room temperature =
Rate constant at ------°C =
Rate constant at ------°C =
Rate constant at ------°C =
Activation energy =
Arrhenius factor =
