Chemistry 212 Lab
Spring 2012
Chemical Kinetics, Part II
The "Iodine Clock" Reaction
Purpose:
A second study in kinetics, involves the method of initial rates, the effects of temperature on rates of
reaction, and the affect of catalysis on rates of reaction. The use of a “timer” or secondary reaction, to study
another, is introduced. Methods and measurements are semi-quantitative.
Introduction:
The "Iodine Clock" experiment is one of the oldest, most tested chemical experiments in all of
laboratory science. It is as the name also suggests: by a careful adjustment of the “ingredients”, the iodine clock
reaction can be made to “finish” as regularly as a clock, and equally as predictable, in the sense of having a
perfectly timed (and thus, measurable in time) reaction sequence. The problems involved with achieving this
level of exact timing: to do so requires very high precision glassware and very careful attention to detail. It is
simply not practical to have the means to do this for second semester undergraduates.
That said, the iodine clock can be modified, as it is in this procedure, to give semi-quantitative
information related to reaction order, and also lends itself well to multiple experiments in which the
concentrations of starting materials can be varied independently of one another. But the reader is cautioned that
absolutely perfect, whole number values for reaction order cannot be achieved by the procedures that follow;
but the answers will be “close enough” to give, ultimately, the complete rate law expression for the iodine clock
sequence of reactions.
Chemical kinetics encompasses a vast field of study, and has three primary goals:
1). To determine the rates of chemical reactions. How long does a chemical reaction take before it can be
considered "complete".
2). To study the factors that seem to most influence the rate. These are primarily the temperature at which a
given reaction is run, and the concentrations of the reacting species, but there are other considerations, and
3). To deduce the series of steps that occur in the conversion of reactants to products. It is the rare chemical
reaction that is not composed of at least two steps.
Those chemical reactions that investigators find particularly interesting, or that have industrial, biological, or
medical interests are studied quite extensively, and their reaction kinetics and by and large well understood. By
learning how chemical reactions occur, and the factors that control how fast a chemical reaction occurs, certain
important reactions (or, those particular reactions that provide us with useful materials) can then be optimized
so that the maximum amount of product can be made, from limited amounts of starting material, in the shortest
possible time.
The rate of any chemical reaction, at the most basic level, relies on only three factors:
1). Reactant molecules must collide, come into direct contact with one another. No collisions = No reaction.
2). These collisions must be sufficiently energetic to cause existing bonds to break, while not preventing new
and different chemical bonds to form
3). The reactant molecules must collide with a favorable orientation.
It is only the first two factors that we can usually control, by way of adjustments to the concentrations of
reactant molecules, and the proper application (or withholding) of heat from a chemical reaction. In nontechnical language, we can only control 1). how many reactant molecules are present in a reaction system, and
2). the temperature at which the reaction is run, by raising or lowering it. In special cases, we can also control
the specific orientation at which reactant molecules collide, but that will not be a variable under study in today's
experiment.
General discussion of rates, and the rate law for a chemical equation
The following is a general discussion of rates, mechanisms, theory, and experimental conditions. The
purpose is to give the reader a greater sense of where the seemingly arbitrary form of a general rate law
expression comes from. There will also be more specific discussions rates, and the rate law specifically for the
iodine clock reaction system.
The rate of any chemical reaction depends, more or less, upon the three factors mentioned previously.
Rates are therefore products of multiple factors, implying that they are comprised of various quantities
multiplied together to give a mathematical expression for the rate of a chemical reaction. Any general rate
expression looks something like this, in its most qualitative form:
Rate = (total number of collisions) x (fraction of sufficiently x (fraction of collisions of
energetic collisions)
proper orientation)
Term #1
Term #2
Term #3
In any given reaction, we can do little (nothing, actually) to control the magnitude of Term 3, the fraction of
reactant collisions going on that have the proper orientation to bring about the transition of reactants to
products. For our purposes, we may assume this term to be a constant value at a constant, stable temperature.
Let us call this constant korientation., signifying that this specific and constant number of collisions with the
correct orientation will be occurring when we carry out this chemical reaction at a specified temperature.
Nor can we do much about Term 2 at a given, stable temperature, which is the fraction of collisions occurring
that have a sufficient energy for the reactant-to-product transformation. Since we do not generally change the
reaction temperature while the reaction is occurring, we may also consider this term to be a constant. Let us
call this constant kenergy., signifying collision energy. How would a rate expression look after these
assumptions/simplifications are made:
Rate = (total number of collisions) x kenergy x korientation
This is a much easier equation to handle, as now we only need to know what factors will influence the total
number of collisions to determine the rate. As usual, it's almost that simple. For our purposes, we can get by
with a simplification that very nearly tells us everything we'd like to know. It is this: the total number of
collisions is a function of concentration. The more molecules you have in a given reaction system, the more
collisions you will have. There are two ways to increase the concentration of reactant molecules 1). you can
simply add more of them, or 2). you can confine the same number of molecules to a smaller volume. Adjust
molarity, in other words. This is all you are doing: influencing the rate by changing the molarity: moles (actual
amount) per liter (volume). Our new and improved rate expression is this:
Rate  k [A]n
where “[A]” is the concentration of a given reactant molecule or atom
raised to an exponential value, usually, 0, 1 or 2
This expression tells us that the rate is equal to a constant (called, conveniently, the rate constant) multiplied
by the concentration of some reactant molecule, A. Remember, k is temperature-dependent--by virtue of having
combined the non-temperature dependent korientation. with the temperature dependent kenergy. Thus, if you change
the reaction temperature, you will also change the rate. And, if you change the reactant concentrations, you will
change the rate. Those are, in fact, the two variables primarily under study in today's experiment.
The "Iodine Clock", and today's experiment
You are probably unaware that there are many variations on the "Iodine Clock" reaction. All of them,
however, rely upon this particular reaction to study concentration and temperature effects:
S2O82-(aq) + 2I-(aq) ----------> 2SO42-(aq) + I2(aq) (I)
Today's experiment has a slightly more complicated rate law than that listed above, as it depends on both the
concentration of persulfate ion (S2O82 and iodide ion (I). The rate law you will use has one further
complication: the inclusion of the orders of each reactant. The order is a measure of how changing the
concentration of a given reactant changes the overall rate of the reaction. The rate law you will complete has
this form:
Rate = k [S2O82-]x [I-]y
where x and y are the orders of persulfate and iodide, respectively. Evaluating x and y requires
experimentation: note that for iodide we do not automatically use the stoichiometric factor of 2. It is important
to realize that reaction stoichiometry and reaction rates are not in any way related. You will determine the
values for x and y by comparing reaction rates of different, independent reactions. It is easy to explain what
you will do, and what you will see, but it's not so easy to explain why you see it (particularly, in the sequence
that you see things happening)
Orders of reaction rates
One of the objectives of this experiment is to determine the value for the orders of the reactants. You
will do this by varying the amounts of the two reactants, first holding one concentration constant, then the other,
and then doing another, single experiment as a "control". “Experiment #1” is the control, and its rate is
measured by a method you will see shortly. Experiment #2 tests the dependence of the rate on the concentration
of the first reactant, while holding the second reactant's concentration constant. The rate law expressions for
experiments #1 and #2 are thus:
Rate1 = k [S2O82-]x [I-]y
Rate2 = k (2[S2O82-])x [I-]y
The rate in experiment #2 will be related to the rate in experiment #1 only by a factor of 2 to the reactant's order
exponent, Rate2 = 2x times Rate1. Similarly, the rate law for the third experiment varies the other reactant,
making the rate in Experiment #3 dependent by 2 to the second reactant's order exponent:
Rate3 = k [S2O82-]x (2[I-])y
in this case, Rate3 = 2y times Rate1. This is why one must occasionally perform multiple reactions to find
kinetic data: to determine the complete rate law with the proper order exponents. You cannot determine the
order information directly from the reaction stoichiometry, and unlike the previous experiment, there is no way
to track the progress of reaction in real time, as it is happening. Most of the time, values for reactant orders are
whole numbers, or close to them, indicating (usually) a doubling (21) or quadrupling (22) of reaction rates as a
function of doubling concentration. You should obtain small whole numbers for your x and y in this
experiment, or be able to sensibly round your experimental values to the nearest whole number.
Mathematical Expression of Rate for the “Iodine Clock”
In all cases, reaction (I) above is what you are trying to study. The moment I2 appears in solution, they will go
instantly dark due to the presence of starch. You will notice in the very first reaction you set up and run that
while persulfate and iodide are in the same solution, the reaction does not instantly turn dark…
The simple reason is that there is a second reaction occurring that delays this instant darkening by starch
combining with the I2 being formed, one that consumes I2 and essentially “delays” the onset of the dark blue
color. The technical term for such a “delay mechanism” is a “coupling reaction.”
In times past, this handout has attempted to explain the coupling reaction, the reason for the delay, and
generally left students baffled and highly confused as to what ends up being the “Mathematical Expression of
Rate” asked of them in the data tables. WITHOUT the technical details it is THIS, for any given experimental
run:
 [S2O82]
Rate =  ------------t
There is a more complicated version of how this is ultimately derived, and thiosufate ion IS involved with that;
but without the details the experiment becomes a simple matter of mixing the quantities together according to
the table in the Experimental Section.
Summary of what to measure, and how to treat your data
1). You must prepare the experiments in the exact manner prescribed, and watch for the sudden appearance of
the blue starch/iodine color, as a function of time, as carefully as possible.
2). You then must calculate the rates of each reaction, using the mathematical rate expression given in this
text. That negative sign is not a typo - think about how the concentration of persulfate has changed over time.
A final-minus-initial determination of change in concentration of persulfate yields a negative number, and thus
the negative sign out in front of “mathematical expression of rate” is a fixer - so that the reaction rate is not a
negative value.
3). Calculate the following quantities in the following order:
Rates for experiments 1, 2, 3, 4, and 5
Determine the value for the orders x and y from the rate data of experiments 1, 2, and 3 only.
Determine the rate constant, k, for experiments 1, 2, 3, 4 and 5. You will need concentrations for I- and for
S2O82 in moles per liter (remember also that you are actually performing dilutions) and the calculated rates
before determining each value of k. You should notice something about the values for k. Be sure to mention
this in your lab experimental conclusion, as well as any other interesting things you notice about concentrations
and rates.
All solutions in this experiment are slightly irritating - to the extent that you'll need to wear gloves. Just
remember that if you spill chemical (ANY chemical) on yourself, you should rinse it off at once. The solutions
in every case can be disposed of down the sink with running water when the experiments have ended.
Experimental Procedure (all solutions may be rinsed down the sink at the end of each trial)
Take a reading of the temperature in the lab to the nearest 0.1° C.
The first 3 experiments use the following amounts of the following reagents:
Experiment
0.20 M
NaI
0.20 M
NaCl
0.010 M
2%
Starch
0.20 M
0.20 M
Na2S2O3
K2SO4
K2S2O8
1
10 drops
10 drops
10 drops
5 drops
10 drops
10 drops
2
10 drops
10 drops
10 drops
5 drops
0 drops
20 drops
3
20 drops
0 drops
10 drops
5 drops
10 drops
10 drops
The best way to do this is to add all five of the first six reagents directly to a small/medium test tube. Add the
sixth reagent from a second small test tube (the specified amount of K2S2O8 per experiment) only when you are
ready to begin timing the reaction that follows to the nearest second. You may either use a wrist watch or the
second hand of a room clock, if one is available. Further experimental notes:
1). Be sure to rinse your test tubes with distilled water before adding your reactants.
2). There will be one set of dropper bottles per bench. Be nice and share it.
3). Begin timing as you begin to deliver the persulfate solution to the solution of the other five reagents. Place
a gloved finger on top of the test tube and gently shake the solution for a good 10 seconds to ensure good
mixing.
4). The blue color you are watching for will appear suddenly. Be sure to record the exact length of time (to the
nearst second) it takes for each trial to go blue.
5). Repeat with experiments 2 and 3, washing and rinsing your test tubes between experiments.
Testing experiment #1 for temperature dependence, and the influence of the "Mystery Reagent"
Experiment #4:
Prepare for another trial of Experiment #1 (to be performed at room temperature) by adding the first five
reagents as before, but with one drop (and one drop only) of the "mystery reagent" from the dropper bottle.
Add the K2S2O8 solution as before, just as you are ready to begin timing. Record the exact time elapsed before
you see the blue color. What IS that stuff?? More to the point, what is called a substance added to reaction that
is neither a reactant, nor a product, and has a profound affect on the reaction rate?
Experiment #5:
Prepare again for another trial of Experiment #1 (no mystery reagent this time) by adding the
combination of the first five reagents to the test tube. In a second, small test tube, pipet the required amount of
K2S2O8 solution. Place both solutions in the provided water bath for at least 30 seconds, and take the
temperature of the five-reagent mixture. You can assume that both the large and small test tubes are at this
temperature. Mix in the K2S2O8 solution and watch for the blue color to appear as before. Don’t look away
from this trial for too long, or you might miss it.
Data Table: Chemical Kinetics, Part II-- The "Iodine Clock" Reaction
Temperature in the lab
_______________ ° C
Experiment #1
Calculated concentration of S2O82-
_______________ mol/L
Calculated concentration of I-
_______________ mol/L
Total elapsed time (t)
_______________ sec
Rate of reaction
(from mathematical rate expression)
_______________ M/sec
Rate constant, k
_______________ sec-1
Experiment #2
Calculated concentration of S2O82-
_______________ mol/L
Calculated concentration of I-
_______________ mol/L
Total elapsed time (-t)
_______________ sec
Rate of reaction
(from mathematical rate expression)
_______________ M/sec
Rate constant, k
_______________ sec-1
Experiment #3
Calculated concentration of S2O82-
_______________ mol/L
Calculated concentration of I-
_______________ mol/L
Total elapsed time (t)
_______________ sec
Rate of reaction
(from mathematical rate expression)
_______________ M/sec
Rate constant, k
_______________ sec-1
Calculated reaction orders (show work in lab report): x: _____ y: _____
Complete rate law (with orders):
______________________________
Experiment #4
Calculated concentration of S2O82-
_______________ mol/L
Calculated concentration of I-
_______________ mol/L
Total elapsed time (t)
_______________ sec
Rate of reaction
(from mathematical rate expression)
_______________ M/sec
Rate constant, k
_______________ sec-1
What is the effect on the rate of reaction caused by the one drop of "mystery reagent?
What is a more common term for these "mystery reagents", those that increase the rate of reaction without being
a part of the rate law expression?
Experiment #5
Temperature at which this reaction was run:
__________ °C
Calculated concentration of S2O82-
_______________ mol/L
Calculated concentration of I-
_______________ mol/L
Total elapsed time (t)
_______________ sec
Rate of reaction
(from mathematical rate expression)
_______________ M/sec
Rate constant, k
_______________ sec-1