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Laboratory ManualREACTION KINETICS
EXPERIMENT 9F
Reaction Kinetics
This experiment is done in pairs.
Useful background reading (this is not compulsory but may be helpful):
“Chemistry Core Concepts”. 1st Edition. John Wiley & Sons Australia Ltd.
Section 13.1, 13.2, 13.3 (pages 564-565), 13.4, 13.5 (pages 575-577); Worked examples 13.1,
13.2; Practice exercises 13.1, 13.2
“Chemistry Core Concepts”. 2nd Edition. John Wiley & Sons Australia Ltd.
Section 13.1, 13.2, 13.3 (pages 759-760), 13.4, 13.5 (pages 774-776); Worked examples 13.1,
13.2; Practice exercises 13.1, 13.2
What is the relevance of this prac?
This experiment relates to the reaction kinetics section of your lectures. You will see how
reactant concentration, temperature and addition of a catalyst affect the rate of a reaction.
Learning Objectives (remember these are different to the scientific
objectives):
On completion of this practical, you will have:

Learned how an increase in reactant concentration affects reaction rate.

Learned how an increase or decrease in temperature affects reaction rate.

Learned how the addition of a catalyst affects reaction rate.

Learned how to use experimentally-derived data to deduce a rate law for a reaction.
In this experiment you will study factors affecting the rate of the oxidation of iodide ions,
I–(aq), with persulfate ions, S2O2–
8 (aq):
–
S2O2–
8 (aq) + 2I (aq)
→
2SO2–
4 (aq) + I2(aq)
Peroxydisulfate (persulfate – note that the two central oxygen atoms are peroxide-like and
hence have an oxidation state of -1 instead of the usual -2):
O
O
O S O O S O
O
O
You will thereby have the opportunity to develop some understanding of this reaction and
also the principles of reaction kinetics.
EXPERIMENT 9F—1
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REACTION KINETICS
Introduction (see pages 16-18 of this prac script for further background information)
Why should the rate of a reaction be important? Indeed, what is the rate of a chemical
reaction? To answer the last question first: the rate of a chemical reaction is a measure of
how fast products are being formed by that chemical reaction. Conversely, it is also a
measure of how fast reactants are being consumed. This will be discussed in more detail
later on.
The rates of chemical reactions are of great importance in industrial and biological
processes. The economic viability of many industrial processes is largely affected by the rate
at which the products can be formed. Of more vital importance, in the true sense of the
word, every chemical reaction taking place in your body is occurring at a rate carefully
controlled by the most complex of catalysts—enzymes. Life would be impossible without
the rates of countless, complicated chemical processes being controlled with exceptional
precision by exquisitely formed enzyme catalysts.
The thermodynamic favourability of a process (ie the work needed to complete that process
or the energy released by that process) can be determined by applying the Laws of
Thermodynamics and simply measuring the initial and final energy of the system. This tells
us nothing about the rate at which the process can occur. Indeed, many highly favourable
processes will not proceed at a measurable rate until something ‘triggers’ or catalyses them
eg: A litre of petrol will burn vigourously in air to produce a great deal of energy. However,
no obvious reaction will be observed until the vapour above the petrol is ignited
(carefully!).1
Reaction kinetics is the study of the rate and mechanism of chemical reactions. Our
understanding of chemistry is greatly increased by studying the details of chemical
processes.
MECHANISM
Refer to Brown et al, “Chemistry: The Central Science”, 2nd Edition, section 12.6.
As chemists, we often write equations describing chemical reactions. Consider the reaction
you will study in this experiment:
–
S2O2–
8 (aq) + 2I (aq)
→
2SO2–
4 (aq) + I2(aq)
When we write such an equation we are making a statement about the starting materials
and the end products of the reaction. A chemical equation does not imply that the process
takes place in one step. To state that this is the reaction mechanism is to imply that the
three reactant molecules will collide absolutely simultaneously and then instantly be
converted to products. This is extremely unlikely.
When studying reaction kinetics an attempt is made to describe a chemical reaction as a
series of single steps or elementary steps. The elementary steps must satisfy two
requirements:
1.
The sum of the elementary steps must give the overall balanced equation for the
reaction.
1It is not recommended that you try this experiment!
EXPERIMENT 9F—2
Laboratory Manual
2.
REACTION KINETICS
The rate determining step (RDS), which is the slowest step in the sequence of steps
leading to product formation, should predict the same rate law as is determined
experimentally.
If it is possible to write a series of chemical equations that describe each of the elementary
steps in a chemical process then that series of steps is a reaction mechanism.2
Thus, the sequence of steps in the study of a reaction mechanism are:
measuring the rate
of a reaction
formulating the rate
law
postulating a
reasonable reaction
mechanism
In the experiment you will do today, you will measure the rate of the reaction of the
oxidation of iodide ions with persulfate ions under various conditions and use your results
to formulate an experimentally determined rate law.
Several possible hypothetical mechanisms for the reaction being studied and the rate law or
rate expression which can be derived from them will be provided. Your experimental rate
law will either support or contradict the mechanisms provided. The mechanism which
supports your experimentally derived rate law will be the likely mechanism for the reaction.
For example, one of the proposed mechanisms for the oxidation of iodide ions with
persulfate ions is:
I–(aq)+ I–(aq)
slow

RDS
*[I….I]2–(aq)
fast
*[I….I]2–(aq) + S2O2–

I2(aq) + 2SO2–
8 (aq)
4 (aq)
These two equations are elementary steps in the overall chemical reaction. They state that
two iodide molecules collide to form an intermediate di-iodide species which lasts long
enough to react, at some time later, with one molecule of persulfate to form the final
products. This is much more plausible than a simultaneous and instantaneous reaction,
however it is only one of several possible mechanisms which you will consider.
RATE DETERMINING STEP AND RATE LAW
Refer to Brown et al, “Chemistry: The Central Science”, 2nd edition, section 12.6.
As already stated it is often the case that one of the elementary steps in a reaction
mechanism is much slower than the rest. If this is so, the overall reaction rate is equal to the
rate of this slowest step and this step is called the rate determining step (RDS).
Some reactants may be not be involved in the RDS.
Consider the hypothetical mechanism for the oxidation of iodide ions with persulfate ions
above. The first step is the rate determining step so the rate of the reaction can be
determined from the first step alone:
2 It is unfortunate that the symbols used to describe elementary steps in a reaction mechanism are identical to
those used in normal chemical equations. You need to learn to take note of the context of the equations.
EXPERIMENT 9F—3
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REACTION KINETICS
rate = k [I–] [I–]
 rate = k [I–]2
This is the rate law or rate expression for the reaction.
If this rate law is the same as the experimentally determined rate law formulated from your
experimental results, the mechanism is the likely one for the reaction. If the rate law
contradicts your results, other mechanisms must be considered.
RATE
The rate of a chemical reaction is defined as being the absolute value of the rate at which
the concentration of reactants or products are changing — divided by the stoichiometric
ratio if necessary eg:
aA + bB
→
cC + dD
The rate of disappearance of A, eg. is (by definition):
Rate
=
–
d[A]
1

dt
a
(ie the rate of disappearance of A)
RATE LAWS AND ELEMENTARY STEPS
Refer to Brown et al, “Chemistry: The Central Science”, 2nd edition, sections 12.4 and 12.6.
Once the mechanism of a reaction is determined, the observed rate of the reaction may be
rationalised and, if necessary, the reaction conditions can be adjusted to control the rate of
the process.
While it is not possible generally to predict the rate of a chemical reaction from the
equation that describes it, it is possible to do so for an elementary step in a reaction
mechanism.
For two atoms, ions or molecules to react they must first collide. The rate of the reaction is
therefore proportional to the rate of collision. The rate of collision is proportional to the
concentration of the species involved in the elementary step.
The rate of an elementary step is proportional to the concentration of reactants involved
in that step.
Consider three simple examples of elementary steps:
Example One
Example Two
A → B
C+D → E
Rate  [A]
Rate  [C] × [D]
Example Three
2F → G
Rate  [F]2
Example One
In Example One above, the process is only affected by the concentration of species A and
this process is described as a First-Order Process.
Rate
EXPERIMENT 9F—4
=
k × [A]
Laboratory Manual
REACTION KINETICS
k
=
rate constant
For a First-Order process the concentration of the reactant and products approach their
equilibrium values exponentially.3
Diagram 9.1
First-Order Process
An important property of First-Order processes is that the time taken for the concentration
of reactants to halve (from any arbitrary starting concentration) is always constant at
constant temperature. This period of time is characteristic of the rate constant and is called
the half-life of the reaction (t½ ).
Example Two
In Example Two, the rate of the reaction is affected by the concentration of the two species
involved. Such a process is called a Second-Order Process.
Rate
=
k × [C] × [D]
In contrast to example one, the half-life of a Second-Order process increases as the reaction
proceeds.
3Because the rate of a reaction varies as the concentration of at least one of the reactants varies, it is not convenient to
compare the rates of different reactions directly. However, the rate constant, which is characteristic of a particular reaction
under certain conditions, does not vary with concentration and so is a useful measure for comparing the rates of different
reactions.
EXPERIMENT 9F—5
Laboratory Manual
Diagram 9.2
REACTION KINETICS
Second-Order Process
Example Three
In Example Three on page 5 the rate of the reaction is only affected by the concentration of
F but in this case the rate will be proportional to the square of the concentration of F. Such
a process is also called a Second-Order Process.
The mathematics of second-order processes is much more complicated than that of firstorder ones and will not be examined here.
EFFECT OF TEMPERATURE ON RATE
Refer to Brown et al, “Chemistry: The Central Science”, 2nd edition, section 12.5.
Earlier it was argued that the rate of an elementary step is proportional to the
concentration of species involved in that step because the rate is proportional to the rate of
collisions between the species involved. Increasing the concentration of species is not the
only way to increase the rate of reaction. The rate of collision can also be increased by
increasing the temperature of the reaction.
Not all collisions will result in the desired reaction to occur. Collision theory states that
colliding molecules must possess sufficient energy between them so that a certain minimum
activation energy (Ea) is exceeded.4 As the temperature of a reaction mixture is raised, the
thermal energy of the reacting species is raised and so the number of ‘successful’ collisions
is increased dramatically.
The Arrhenius Equation describes the relationship between rate constant and temperature:
4Brown et al,
k
=
A × e–Ea/RT
k
=
rate constant
Ea
=
activation energy (Joules)
R
=
gas constant (8.314 J K-1 mol-1)
T
=
temperature (Kelvin)
A
=
collision frequency
“Chemistry: The Central Science”, 2nd edition, section 12.5.
EXPERIMENT 9F—6
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REACTION KINETICS
CATALYSTS
Refer to Brown et al, “Chemistry: The Central Science”, 2nd edition, section 12.7.
In the previous section it was stated that in order for collisions between reactants to be
‘successful’ the molecules had to possess sufficient energy to overcome an activation energy
characteristic of the particular reaction. A consequence of the Arrhenius Equation is that as
the temperature increases so too does the rate constant.
The rate constant could also be increased if the activation energy can be reduced. Catalysts
are species that are not consumed in a reaction but provide an alternative pathway for a
reaction. This alternative pathway has a lower activation energy and therefore the reaction
will proceed faster.
THIS EXPERIMENT
It is often possible to plot the changing concentration of reactants or products of a chemical
reaction as a function of time. The rate of the reaction at any time during the measurement
period can then be determined by measuring the slope of the tangent to the curve at that
point.
The experiment that you will perform is designed so that it is reasonable to assume that the
rate is constant throughout the measurement period.5 Therefore, it is reasonable to assume
that the average rate of the reaction is equal to the initial rate of the reaction.
Under these conditions the time taken to produce a certain quantity of product can be used
to calculate the initial rate of the reaction.
Rate
[Product]
t
=
d[Product]
[dt]
≈

1
Dt
(if Δ[Product] is constant)
(for small Δt)
The following diagram indicates the following important points regarding the experiment
you will perform:

The rate of a reaction is approximately constant if the observation period is much less
than the half-life of the reaction.

The time taken to form a certain, fixed quantity of product (Δt) is inversely proportional
to the rate constant (or rate).
5For this approximation to be reasonable,
Δt, the period of observation, must be much less than the half-life of
the reaction ie: Δt