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Iodine Clock Reaction: Inquiry-Based Lab Investigation
INTRODUCTION
The use of chemical kinetics to solve reaction mechanisms is predominantly based on the concentration dependence of the
rate of reactions. The rate of any chemical reaction will depend on concentration simply because, if the concentration of a
reactant is zero, there is nothing there to react and hence the rate would be zero. The whole theory of chemical kinetics is
based on probability. Consider the simplest possible reaction, A  B, in which reactant A spontaneously changes into
product B. (Radioactive decay is a typical example.)
The rate of this reaction, is defined as the rate of change of the concentration of the reactant A or of the product B. For this
simple reaction rate is related mathematically to the concentration of A, [A], by rate = k[A]. [A] is simply a measure of
the probability that a molecule of A is present in the volume of solution that you’re measuring. As the concentration
increases so does the probability that some A is present and hence the rate of reaction increases. The constant k, termed
the rate constant, is a constant of proportionality and is equal to the rate that the reaction would have if the concentration
of A were 1 M. The value of k depends on the chemical nature of A and on the experimental conditions, e.g. temperature
and pressure.
Now let us consider a slightly more complicated reaction, A + B  C. For this reaction, a molecule of A and a molecule
of B must meet or collide with one another to react and form C. Therefore, again based on probability, the rate of the
reaction must be proportional to the probability of finding a molecule of A and a molecule of B in the same unit volume at
the same time, which is given by the probability of finding a molecule of A multiplied by the probability of finding a
molecule of B (just like the probability of two heads when tossing coins is ½ x ½). Using concentration again as a
measure of the probability of finding a molecule, the rate of this reaction is given by rate = k[A][B]. As in the previous
example k is the rate the reaction would have if A and B were both present at a concentration of 1 M. Following the same
logic, it is simple to show that for a reaction 2A  B (i.e. A + A  B), the rate would be given by rate = k[A]2.
Many reactions, however, proceed in a number of small steps. Let us consider the hypothetical composite reaction 2A +
B  D. Of the many possible ways this reaction might occur, two are given below:
Mechanism (1) A + A  C, followed rapidly by C + B  D,
Mechanism (2) A + B  C, followed rapidly by C + A  D
To determine which of these two possible mechanisms is operating, one could carry out kinetic experiments on the
dependence of the rate of formation of D on the concentrations of the reactants A and B. Because in each mechanism the
first step is much slower than the second step, it will be a “bottleneck” and determine the rate of formation of D. Under
these conditions the rate of formation of D, for the two mechanisms would be given by:
rate of formation of D via mechanism (1): v = k[A]2, or
rate of formation of D via mechanism (2): v = k[A][B].
Therefore, measurements of the rate of the reaction at different concentrations of A and B would allow one to distinguish
easily between the two mechanisms. If the rate depended on the square of the concentration of A and not at all on the
concentration of B, this would support mechanism (1). If the rate of the reaction were directly proportional to the
concentrations of both A and B, this would support mechanism (2). If the reaction rate does not conform to either of these
relationships, then some other mechanism must apply. This is a simple example of how kinetics research can be used to
determine how a reaction occurs.
BACKGROUND
The iodine clock reaction was discovered by the Swiss chemist Hans Heinrich Landolt in 1886. There are a number of
variations of it, but all of them involve the mixing of two colorless solutions. Initially there is no visible reaction, but after
a certain period of time, the mixed solution suddenly turns dark blue. In the first part of this experiment you will study the
iodate variation of the iodine clock reaction, which involves the reaction between iodate ions, IO3–, hydrogen sulfite ions,
HSO3–, and hydrogen ions, H+. The stoichiometry of the overall reaction is as follows:
2IO3–(aq) + 5HSO3–(aq) + 2H+(aq)  I2(aq) + 5HSO4–(aq) + H2O(l) (1)
If the solution contains starch, it will react with I2, the elemental iodine produced by the reaction, and form a blue starchiodine complex. The color however does not appear immediately, because as soon as any I2 is formed, it immediately
reacts with any HSO3– still present and is converted into colorless I–:
I2(aq) + HSO3–(aq) + H2O(l)  2I–(aq) + HSO4–(aq) + 2H+(aq) (2)
Thus, reaction (2) efficiently removes any I2 produced by reaction (1). Only after all of the HSO3– has been consumed by
reactions (1) and (2) will the I2 concentration increase and its reaction with the starch begins. This not only explains the
sudden appearance of the blue color, but it also allows us to determine the rate of the reaction. A measure of the rate of a
reactant’s consumption can be made by dividing the initial concentration of reactant of interest by t, the time it takes for
the blue color to appear, [Reactant]o /t.
This question demonstrates a fundamental principle of kinetic research. Solving reaction mechanisms always involves
making a theoretical prediction and then doing experiments to test your prediction. If the prediction turns out to be wrong,
then you must modify your mechanism and make a new prediction, until you have a mechanism that is consistent with the
experimental result. A word of warning, however; consistency doesn’t necessarily mean a mechanism is right, simply that
it agrees with the data currently available. Further experiments could still prove it to be wrong.
CHEMICALS PROVIDED
0.006M HCl
Concentrated HCl
0.2M KIO3
0.2M NaHSO3
Solid Na2S2O5 (reacts with water to create sodium bisulfite—assume full dissociation)
Distilled Water
0.05M CuCl2
PROCEDURE
PART 1: Considering a Reaction Mechanism
Design and carry out an experiment to determine if the iodine clock reaction occurs in a single step. To answer this
question, you first need to consider the following question: If Reaction (1) were a single-step reaction (elementary
reaction), how would you expect the overall rate to depend on the concentrations of reactions: IO3–, HSO3– and H+?
Hints
*You will need to use the pseudo-first order lab technique.
*If you vary two or more things simultaneously, it is impossible to know what change is causing any of the differences
that you observe.
*By combining different solutions together, you effectively dilute them from their original concentrations.
*Integrated rate laws are probably not necessary
Suggested Base-Line Procedure
Use two clean beakers to prepare initial solutions of KIO3 and NaHSO3. Add approximately 40mL of the stock solution
of HCl to each beaker. Add 2.0 mL of KIO3 to the first beaker and add about 2.0mL of NaHSO3 to the second beaker.
The concentrations of the two solutions will change due to the addition of the great volume of HCl. However the
concentration of HCl should not be dramatically impacted by the addition of the two solutions. Add a scoopula tip of the
starch to the beaker containing the potassium iodate and swirl to dissolve fully. With a timer ready, pour the contents of
the KIO3 beaker into the NaHSO3 beaker. The timer should start the second the two solutions meet. Mix with a stirring
rod, by swirling, or by pouring the two mixtures back and forth until the blue-black color appears. Stop the timer when it
appears. The rate can be determined by the initial concentration of the reactant of interest, adjustments necessary based
upon dilution divided by the time it took for the color to appear. All products can be put down the drain with plenty of
water. Repeated trials require thorough washing of the beakers!
DATA ANALYSIS
Create necessary graphs to determine order of each reactants.
Determine whether it is plausible to propose a one-step mechanism for this reaction.
Calculate the value of k.
PART 2: Determining the Activation Energy
Design and carry out a procedure to determine the activation energy required for the un-catalyzed reaction to occur. The
only way that we know of to determine the activation energy is to prepare an Arrhenius plot. An Arrhenius plot has an xaxis of 1/T, where T is temperature in Kelvin, and a y-axis of ln k, where k is the rate constant, and the slope of that line is
–Ea (activation energy)/ R (gas-law constant: 8.314 J/mol-K).
Hints
*Concentrations do not need to vary, but temperatures do.
*At least three different trials should be run—it might make sense to do room temperature, above room temperature and
below room temperature.
DATA ANALYSIS
Create the necessary graph to assess activation energy.
What is the activation energy of the reaction?
PART 3: Adding a Catalyst
Design and carry out a procedure to determine the affect of a catalyst, Cu2+ on the rate of the reaction. A catalyst is a
substance that increases the rate of reaction (both in the forward and reverse directions) without being consumed in the
reaction. Catalysts often allow the reactants to react using a reaction pathway (or mechanism) that is lower in activation
energy than without the presence of the catalyst. Though it is not necessarily true, we will assume that the exponents “x”,
“y”, and “z” of the given reactants don’t change with the addition of the catalyst. Copper ion is not even that good of a
catalyst. It lowers the activation energy, but not nearly as much as most biological enzymes lower the activation energy.
Because the decrease in activation energy with the catalyst is small, it is hard to measure accurately even though the
difference in rate with the catalyst is easy to see.
Hint
*Carry out multiple trials—only the catalyst concentration should vary!
DATA ANALYSIS
How does the catalyst affect the rate?
Does the concentration of the catalyst impact the rate?
Could you determine how the activation energy changes?
PART 4: Setting the Clock
Design and carry out a procedure that sets the reaction clock for exactly 30 seconds. Be ready to duplicate your results!
Lab Schedule
Wednesday, November 4th: D-Day, 80 minutes in lab
Thursday, November 5th: E-Day, Kinetics Test
Friday, November 6th: A-Day, 60 minutes in lab
Monday, November 9th: B-Day, 40 minutes lab
Thursday, November 11th: D-Day, 60 minutes in lab, 20 minutes group presentations & setting the clock challenge