CHEMISTRY 150
Chemistry for Engineers
IODINATION OF ACETONE
DEPARTMENT OF CHEMISTRY
UNIVERSITY OF KANSAS
Iodination of Acetone
Introduction
Reaction rate is a fundamental property of a chemical transformation. Changes in reaction conditions (e.g., reactant concentration(s), temperature, pH) can influence the rate with which reactants are consumed and products are formed in a given process. Studying the reaction kinetics,
or the rates of reactions, is often essential for fully understanding the mechanism of a chemical
transformation.
A rate law is an expression that relates the reaction rate to the reactant concentrations. For a
general reaction of the form,
A + B → C + D
(1)
the reaction rate, v, is the time rate of disappearance of A (or B) which must, by stoichiometry,
equal the time rate of appearance of C (or D). One way to think about the rate is to consider a
small time interval Δt. If the change in concentration of A during Δt is denoted Δ[A], then the
rate is v = −Δ[A]/Δt = Δ[C]/Δt. Note that the rate has units of concentration/time, e.g., M/s. For
this generic case, the rate law is of the general form
rate = v = k[A]x[B]y.
(2)
In this expression, [A] and [B] are the concentrations of the reactants, x and y define how
strongly the rate depends on these reactant concentrations, and k is the rate constant for the process at a given temperature. The exponents in a rate law (defined here as x and y) are frequently
positive integers. However, they can be fractions, zero, or even negative. These exponents are
called the reaction orders. In this case, the reaction is order “x” (or xth-order) with respect to reactant A and order “y” (or yth-order) with respect to reactant B. Thus, if it was determined that x
= 1 and y = 2, the reaction would be first order in A and second order in B. The overall reaction
order is given by the sum of the exponents. In this example, the reaction would be third order
overall as x + y = 3. As indicated in Eq. (2), rate laws are also defined by the rate constant, k. (It
might be a bit confusing, but it is important to understand the difference between the reaction
rate, v, and the reaction rate constant, k.) For two reactions described by the same rate law, the
reaction with the larger rate constant proceeds more quickly. Rate constants, like the equilibrium
constant, depend on temperature. Rate laws are generally phenomenological, in that they must
be determined empirically and even the same reaction can have different rate laws under different conditions. In this experiment you will determine the rate law for a reaction.
Consider the definition of the reaction rate given above. If the concentration of A is measured as
a function of time as the reaction proceeds, one could write the result as [A](t). Plotting [A](t)
versus t would yield a decaying curve (since [A] decreases with time as the reaction proceeds).
At any time t the slope of this curve is the instantaneous reaction rate v(t). Obviously it would be
convenient to know the function [A](t) for a given rate law. In many cases, this can be obtained
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by “solving” the rate equation. The result is called an integrated rate law, and gives [A](t),
[B](t), [C](t), etc.
The integrated rate law is often the pathway to determining the phenomenological rate law, that
is, the exponents x and y in Eq. (2). For example, if the concentration of a reactant or product
can be measured for different initial values of the reactants, the result can be compared to what
would be expected based on the integrated rate law. That is, the measured [A](t) can be compared with the expected [A](t) from different possible rate laws. By making this comparison for
different initial concentrations of A and B the rate law can be unambiguously determined.
In this laboratory experiment, you will gain first-hand experience acquiring kinetic data. You
will measure the time it takes for a chemical reaction to occur, and use that information to calculate the overall rate of the reaction. In this experiment, you will study the acid-catalyzed reaction
between acetone (CH3COCH3) and iodine (I2).
CH3COCH3 + I2 (aq, brown)
H 3O +
→
CH3COCH2I + HI
This is a convenient chemical reaction for us to study because I2 is a colored species, and
changes in I2 concentration can be observed with the eye or easily monitored with spectrophotometers like those in your laboratory. A catalyst will be used to ensure that the reaction proceeds fairly quickly. In this case, excess hydronium ions (H3O+) serve as the catalyst and are introduced by adding HCl to the reaction mixture. Recall that hydrochloric acid is a strong acid
that ionizes completely in water to form the hydronium ion (H3O+) and the chloride ion (Cl-).
HCl (aq) + H2O (l) → H3O+ (aq) + Cl- (aq)
In the chemical equation showing the reaction between acetone and iodine, H3O+ is written directly above the reaction arrow, as it is not a reactant, per se, but it is an important participant.
The rate law for the iodination of acetone then has the following general form:
rate = v = k[CH3COCH3]x[H3O+]y[I2]z,
(3)
where [CH3COCH3], [H3O+], and [I2] refer to the concentrations of acetone, hydronium ion, and
iodine, respectively. As discussed previously, k is the reaction rate constant and x, y, and z represent the order of the reaction with respect to CH3COCH3, H3O+, and I2, respectively. It turns out,
however, that this reaction is zero order with respect to iodine (i.e., z = 0) under the conditions
we will use here. This means that Eq. (3) can be simplified to
rate = v = k[CH3COCH3]x[H3O+]y,
(4)
since [I2]0 = 1. Equation (4) shows that the rate of the reaction does not depend on the I2 concentration. Keep in mind, however, that I2 still must be present for the reaction to occur, even
though the I2 concentration does not affect how fast the reaction occurs for your experiments.
The goal of this laboratory experiment is to determine the rate law for the iodination of acetone.
You and your group members will conduct a series of simple experiments to determine not only
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x and y but also k, the rate constant for the iodination of acetone at the lab room temperature.
How can we determine so many different values using just one experiment? The answer is that
we will do a series of closely related experiments, taking advantage of the variety of concentrations of reaction participants available to you. Specifically, you will vary the concentration of
acetone in some experiments, and you will vary the concentration of HCl (that is, [H3O+] since
HCl in water can be assumed to fully ionize to Cl- and H3O+), in others. In each of these individual trials, your goal will be to assess the rate at which I2 disappears. Since the reaction rate
involves both concentration and time, the two pieces of “raw data” needed, at minimum, are the
starting concentration of I2, along with the amount of time required for it to disappear. We will
call the reaction starting time t0 and we will refer to the ending time (that is, when all of the I2 is
gone and the solution has turned colorless) simply as t. Using the reaction and conditions described here, you can assume that both x and y are equal to some small whole number.
Pre-lab
Safety: Goggles must be worn at all times. Most chemicals can be toxic and hazardous if
splashed on clothing, exposed skin or in the eyes. At the very least, some of the compounds used
in this laboratory can permanently stain your clothes. If chemicals splash on skin or clothes, remove the affected clothing and flush the affected areas thoroughly with cold water.
Iodine is toxic and stains skin and clothing. Note: A small skin stain is not dangerous. Iodine
solutions were once used to treat minor cuts.
Acetone is highly flammable. There should be absolutely no open flames in the laboratory. Acetone and acetone solution bottles should be closed when not in use.
HCl and other acids are hazardous if splashed on clothing, on exposed skin, or in the eyes. Prolonged exposure of the skin to even fairly dilute solutions of acid can cause burns. If acids or
bases splash on your skin or clothes, remove the affected clothing and flush the affected areas
thoroughly with cold water.
Pre-lab Assignment: Please write out the following in your lab notebook. This assignment must
be completed before the beginning of lab. You will not be allowed to start the experiment until
this assignment has been completed and accepted by your TA.
1) Briefly describe the objectives of this experiment.
2) Write out the experimental procedure in your lab notebook according to the “Guidelines for
Keeping a Laboratory Notebook” handout.
In addition to these pre-lab requirements, a short quiz will be given at the beginning of lab based
on the material in this lab write-up.
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Procedure
Part 1 - A First Timing of the Reaction
In Part 1 of this experiment, you will practice preparing a reaction mixture and carrying out a
timed trial for the iodination of acetone. Please follow the steps given below.
1. Turn on the Ocean Optics spectrometer and, after a brief warm-up, calibrate the instrument.
2. Prepare an acetone solution by placing 10.0 mL of pure acetone in a 50 mL volumetric flask,
and diluting to the mark with distilled water. Calculate the molar concentration of this solution. The density and molar mass of acetone are 0.791 g/mL and 58.078 g/mol, respectively.
3. Place 1.9 mL of distilled water in a cuvette.
4. Add 0.8 mL of 1.0 M HCl.
5. Add 0.8 mL of the acetone solution prepared in step 2.
6. Ready your stopwatch/clock to time your reaction.
7. Add 0.5 mL of 0.005 M I2 solution.
8. Begin timing the reaction immediately upon addition of the I2. Place a cap on the cuvette and
invert the cuvette several times to mix the contents.
9. Place the cuvette in the calibrated spectrometer. Monitor the absorbance at 400 nm; record
these readings at regular intervals until the absorbance reaches a nominal zero value (i.e.,
negligible absorbance indicating the absence of I2).
Your group should repeat this procedure until you are convinced you can complete it reproducibly at which point you can move on to Part 2 where you will plan your own runs to investigate
the effect of different concentrations.
Part 2 - Studying the Effect of [CH3COCH3] and [H3O+]
As you plan for subsequent timed trials – that is, the ones in which the concentrations of acetone
and acid are varied, with the aim of determining the exponents, x and y, in the rate law – there
are several experimental constraints that must be borne in mind. First, our spectrometer, like any
spectrometer, has a limited working range for absorbance measurements. As such, it is recommended that you do not adjust the concentration of I2 from one trial to the next. Instead, add 0.5
mL of 0.005 M I2 for every trial (see Table 1). This is not a practical inconvenience, since we
have already stated that [I2] does not impact the rate of the chemical reaction under the conditions selected. Second, the reagent solutions are limited to your prepared acetone solution and
the provided stock solutions. As such, your group will have to adjust the volume of a given reagent added to a cuvette if you wish to vary its concentration. This approach works well because
each cuvette holds 4.0 mL of liquid; any increase or decrease in the quantity of acetone or HCl
added should be accompanied by a corresponding change in the volume of added water, so that a
net volume of 4.0 mL is employed for every trial. (This is also a practical concern; the spec-
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trometer is not designed for use with partially filled cuvettes.) For example, if your group
wanted to double the concentration of acetone relative to the “starting recipe” given above, the
following combination of reagents would be recommended:
1.1 mL distilled water
1.6 mL acetone solution
0.8 mL 1.0 M HCl (acid)
0.5 mL 0.005 M I2 solution
The “starting recipe” from Part 1 and the reaction mixture shown above are listed as Trial 1 and
Trial 2, respectively, in Table 1. Using this scheme, your group can systematically double, triple,
or even halve the concentration of acetone or H3O+ for a given trial. Once your group has devised a plan, you must have it approved by your TA before you begin your experimental
measurements.
Construct a table in your lab notebook similar to Table 1 shown below. Data corresponding to
each trial should be recorded in this table. Keep in mind that the number of rows in your table
will depend on the number of trials you carry out.
Table 1. Volumes of Reagents Used in the Timed Trials and the Reaction Time
Trial #
Volume of
Volume of
Volume of
0.005 M I2 ____ M CH3COCH3 1.0 M HCl
Volume of
Distilled Water
Total
Volume
1
0.5 mL
0.8 mL
0.8 mL
1.9 mL
4.0 mL
2
0.5 mL
1.6 mL
0.8 mL
1.1 mL
4.0 mL
3
0.5 mL
4.0 mL
4
0.5 mL
4.0 mL
5
0.5 mL
4.0 mL
6
0.5 mL
4.0 mL
Reaction
Time, (s)
Part 3 - Interpreting your Data
Construct a second table in your lab notebook similar to Table 2 shown below. The number of rows
will depend on the number of trials you carried out.
Table 2. Initial Reactant Concentrations and the Reaction Rate for each Trial
Trial #
1
[I2]i
(M)
[CH3COCH3]i
(M)
[HCl]i
(M)
2
3
4
5
6
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Reaction Rate, ν
(M/s)
Use the data from Table 1 to help you determine the “initial” concentrations of I2, CH3COCH3,
and HCl for each trial. Record these concentrations in Table 2. Note: These “initial” concentrations are NOT the reagent concentrations listed in the Table 1 column headings.
The rate of disappearance of I2 is calculated from your experimental data using the equation
rate = v = -Δ[I2]/Δt = -([I2]f - [I2]i)/Δt.
This equation can be simplified to
rate = v = [I2]i/tr,
where [I2]i is the initial concentration of I2 and tr is the reaction time: the time elapsed from the
start of the reaction (i.e., upon adding the I2) until the absorbance reaches a negligible value (indicating the complete consumption of I2). This is an approximate approach in which it is assumed that tr is sufficiently short so that the two equations given above are equivalent. For each
trial, use [I2]i from Table 2 and the reaction time from Table 1 to determine the associated reaction rate. These reaction rates should be recorded in the last column of Table 2.
If you have carried out a sufficient set of trials in Part 2, you will be able to use the data recorded
in your table to determine the values of x and y. First, examine all trials in which the concentration of the acid (HCl) was held constant while the concentration of acetone (CH3COCH3) was
varied. The change in rate, if any, observed in these trials where [CH3COCH3] is systematically
varied should exhibit a simple pattern. If the concentration of acetone does not have any impact
on the reaction rate, it can be dropped from the rate law entirely since it indicates that x=0 in Eq.
(4), that is, the reaction is zeroth-order in CH3COCH3. If [CH3COCH3] does impact the reaction
rate, you should carefully examine how, on a quantitative level. If doubling the acetone concentration doubles the rate, halving the acetone concentration halves the reaction rate, etc., then the
reaction is said to be first order with respect to CH3COCH3. In such a case, the exponent x for
acetone in the rate law would simply be equal to 1, and you have just elucidated an important
part of the rate law! If, on the other hand, doubling the concentration of acetone quadruples the
reaction rate, you’ve just identified the mathematical phenomenon of a second order reactant,
and the exponent associated with [CH3COCH3] should be noted as 2. You can use this same approach to determine the simple whole-number exponent for H3O+. In lab, your TA will discuss a
more formal approach to analyzing your experimental data and determining x and y.
Once you have determined x and y, you have sufficient information to determine the value of the
rate constant, k. Calculate a value of k using the data you obtained in each trial. Use these values to determine an average value for the rate constant (kave).
Report
Your lab report should be a formal, individual report prepared according to the “Guidelines for
Laboratory Reports” you have been given. In addition to the categories discussed in these guidelines you should provide answers to all the questions posed in this laboratory experiment writeup.
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Glossary
integrated rate law
an expression for the time-dependent concentration
of reactants and products in a reaction obtained by
“solving” the rate law (see below); comparison of
concentrations as a function of time with measured
concentrations versus time is one way to determine
the rate law
rate constant
the constant of proportionality between the reaction
rate and the reactant concentrations to their reaction
orders; for analogous rate laws (i.e., same reaction
orders) the rate constant measures the speed of the
reaction; the rate constant depends only on temperature
rate law
a generally empirical relationship between the rate
of a reaction (see below) and the concentrations of
the reactants, e.g., rate = v = k[A]x[B]y; a rate law
is defined by the reaction order of each reactant,
e.g., x and y above, and the rate constant, k
rate of reaction
the speed with which a reaction proceeds expressed
in terms of concentration of reactants consumed per
unit time or concentration of products produced per
unit time; if A is a reactant and C is a product, then
the rate can be written as v =-Δ[A]/Δt = Δ[C]/Δt,
for the limit Δt→0; the rate generally depends on
time, changing as the reaction proceeds
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