University of Puget Sound
Department of Chemistry
Chem 111
Spring, 2010
RATES OF REACTIONS/A CLOCK REACTION
LEARNING GOALS AND ASSESSMENTS
1. Understand how the Clock Reaction works.
a. Write the chemical equation for the reaction studied in this experiment.
b. Write chemical equations corresponding to the “clock” part of the experiment.
c. Explain why thiosulfate does not appear in the rate law, Eq. [4], for this reaction.
d. Predict whether Δt will be bigger or smaller when the amount of thiosulfate is increased.
e. Calculate reagent concentrations from given stock concentrations.
2. Understand general principles regarding the rate laws as deduced from initial rate data.
f. Given a rate law, predict relative changes in reaction when initial concentrations of
reagents are doubled, halved, etc.
g. Given a rate and a set of concentrations, calculate a rate constant.
INTRODUCTION
A general chemical reaction is given in equation [1]. The reaction will be used to illustrate some
ideas regarding chemical kinetics.
A
+
B
C
[1]
There are two ways of expressing the rate of a chemical reaction. The rate can be expressed in
terms of the change in concentration of the reactants or products with respect to time,
rate =
-∆[A] -∆[B] ∆[C]
=
=
∆t
∆t
∆t
[2a]
where ∆[A] represents a change in the concentration of reactant A, and so forth. Ordinarily, the rate
varies as a reaction proceeds. A given rate measurement could be the initial rate, or the average rate
over a longer time interval.
The rate can also be expressed in terms of concentrations of reactants,
rate = k[A]m[B]n
[2b]
where k is known as the rate constant, characteristic of a given reaction run under defined
conditions. The exponent m is the order of the reaction with respect to A, and n is the order of the
reaction with respect to B. The sum of m and n define the order of the reaction as a whole. The
order of a reaction is an experimentally determined value which cannot be deduced from the overall
balanced equation for the reaction.
In this experiment we will be studying the chemical kinetics of the following reaction between
iodide ion, I-(aq), and bromate ion, BrO3-(aq), under acidic conditions:
1
6 I-(aq)
+ BrO3-(aq)
+ 6 H+(aq)
3 I2(aq) + Br-(aq)
+ 3 H2O
[3]
This reaction proceeds reasonably slowly at room temperature, its rate depending on the
concentrations of the I-, BrO3-, and H+ ions. For this reaction the rate expressed in terms of
concentration of reactants, or rate law, takes the form
Rate =
-∆[BrO3-]
= k[I-]m[BrO3-]n[H+]p
∆t
[4]
The main goal of this experiment is for you to determine the values of m, n, and p, the
reaction orders, and the value of the rate constant, k, for this reaction.
Our method for measuring the rate of the reaction involves what is called a “clock” reaction.
How does the clock work? We add two extra reagents to the reaction vessel, starch and thiosulfate
(S2O32-(aq)). At first, the I2 produced by reaction [3] is gobbled up by the thiosulfate as fast as it is
made,
I2(aq)
+
2 S2O32-(aq)
2 I-(aq)
+
S4O62-(aq)
[5]
The products of this reaction are colorless – you can’t see them. However, once the thiosulfate is
used up, the concentration of I2 is free to react with starch, forming a starch complex you can see
(it’s deep blue).
I2(aq) + starch
→
blue starch/I2 complex
[6]
So the clock begins ticking when the reagents are mixed, and stops ticking when the solution turns
color, meaning the thiosulfate has been used up.
How does the formation of the blue starch/I2 complex give information on the amount of BrO3ion reacted? Working backwards from reaction [5], 2 moles of thiosulfate are consumed for every
mole of I2. According to equation [3], three moles of I2 are produced for every one mole of BrO3-.
So, for every mole of BrO3- that reacts, 6 moles of thiosulfate reacts. Therefore, when the blue color
forms, the amount of BrO3- that has reacted (i.e., ΔBrO3-) is equal to 1/6 the initial concentration of
thiosulfate. Fortunately, you only need to calculate ΔBrO3- once because each experiment has the
same initial concentration of thiosulfate. To find the rate you need to divide ΔBrO3- by ∆t, the time
it took to turn blue (see Eq. [2a]), which of course will be different for each experiment. Since the
amount of thiosulfate is small compared to the amount of BrO3- , this rate can be considered an
initial rate for the reaction.
In this experiment you will observe the rate of reaction [3] for various initial concentrations of
[BrO3-], [H+] and [I-]. By changing the concentration of one reactant at a time, you can determine
the order of the reaction with respect to that reactant (Eq. 4). Once you know that, you can
determine the rate constant too.
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PROCEDURE
DEPENDENCE OF REACTION RATE ON CONCENTRATION.
Since there are several reagents to mix, and since you don't want the reaction to start until you are ready,
you need to put some of the reagents into one flask and the rest into another, selecting them so that no
reaction occurs until the contents of the two flasks are mixed. Using the burettes that are set up to
measure volumes, measure 10.0 mL of 0.010 M KI, 10.0 mL of 0.001 M Na2S2O3, and 10.0 mL of
distilled water into a 250-mL Erlenmeyer flask (Reaction Flask I). Then measure 10.0 mL of 0.040 M
KBrO3 and 10.0 mL of 0.10 M HCl into a 125-mL Erlenmeyer flask (Reaction Flask II). To Flask II add
3-4 drops of starch solution. When the flasks are mixed the total reaction volume will be 50.0 mL.
Pour the contents of Reaction Flask II into Reaction Flask I and swirl the final solution to mix it
thoroughly. Note the time at which the solutions were mixed. Continue swirling the solution, and note
when the blue color appears (it should 1-2 minutes). You should also record the temperature of the
solution, to 0.1°C. Repeat the reaction using the same volumes of each reactant until the times you obtain
are reproducible (±5 sec.)
The reaction flasks should be rinsed with distilled water between runs. It is important to keep the
temperature the same in all the runs. Repeat any reactions that did not appear to proceed properly.
Measure the temperature after each reaction to assure that the temperature was the same for each
reaction.
To determine the order of the reaction for each reactant, you want to measure the time it takes for color
to appear for two different sets of solutions. For example, you can repeat the reaction described above
but double the volume, hence doubling the concentration, of the 0.010 M KI solution. To maintain the
same total volume for the reaction solutions, eliminate the 10.0 mL of distilled H2O. (NOTE: The same
total volume is required for each reaction mixture in order to maintain a constant concentration of all
other reactants.) You will now have data that will allow you to determine the order of the reaction with
respect to KI.
Continue your investigation of the reaction rate by varying the concentration of one reactant at a time, in
a logical manner. Once you determine how the rate of the reaction varies with each reactant you will be
able to determine the order of the reaction with respect to each reactant.
All reagents can be safely disposed down the sink.
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WHAT TO DO
Before Lab: Read this handout.
Set up your lab notebook appropriately for this experiment including a title page which
includes the title, your name and your partner's name (to be determined when you get to lab), the
date, TA's name (to be determined when you get to lab), the purpose of the experiment and
disposal instructions for all reagents.
You should lay out two tables before coming to lab -- they don’t have to be complete,
although you’ll finish the lab earlier if you do.
•
In Table 1, summarize the volumes of each reagent for all the reactions you plan to run.
Your table should include the volumes for the two experiments suggested below, plus
any additional experiments you plan to do (however many you think you’ll need).
Remember, your goal here is to do enough experiments so that you can deduce the rate
law for your reaction.
TABLE I. VOLUMES OF REAGENTS USED FOR RATE LAW DETERMINATION
Flask I
Exp.
#
1
2
…
•
Flask II
mL
0.010 M KI
mL
0.001 M Na2S2O3
mL
H 2O
mL
0.040 M KBrO3
mL
0.10 M HC1
Starch
mL
10.0
20.0
10.0
10.0
10.0
0.0
10.0
10.0
10.0
10.0
3 drops
3 drops
In Table 2, summarize the initial concentrations of each reagent present in each of your
rate experiments in the reaction mixture (i.e., after you combine flasks I and II). For
example, for experiment #1 described above, the concentration of [KI] in the final
reaction vessel would be determined in the following way.
[KI] =
(mL used)(0.010 M)
(total mL)
=
(10.0 mL)(0.010 M)
(50.0 mL)
[KI] = 0.0020 M
TABLE 2. INITIAL CONCENTRATIONS IN MOL/L OF REAGENTS USED FOR RATE LAW DETERMINATION
•
Exp. #
[KI]
1
2
…
0.0020 M
[Na2S2O3]
[KBrO3]
[HCl]
Finally, using the information from Table 2 for [Na2S2O3], calculate ∆[BrO3-]. Here’s
how: first calculate the concentration of S2O32- initially present, then calculate the
corresponding change in concentration of BrO3- when the color changes to blue.
Remember, for every mole of S2O32- that gets used up, only 1/6 of a mole of BrO3reacts.
4
During the Lab: With your partner, complete Tables 1 and 2, and turn in a copy of all
relevant pages from your notebook. Then you may start the experiment. Record the time,
temperature and any relevant observations for each of your experiments.
At Home: Prepare a table (Table 3) that includes the times, temperatures, and rates for each
reaction. Then determine the reaction order for each reactant and the overall reaction order.
Once you’ve obtained these values, you should calculate a rate constant for each run. Since
you’ll get a different value for each experiment, you should average these values. Be sure to
write the units for the rate constant.
To Be Turned In at the Beginning of the Next Lab: A typewritten summary page that
includes Tables 1-3 and a brief description of your results (reaction orders and average rate
constant). Refer to the "Guidelines" for the format of your tables!
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