Kinetics

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A KINETIC STUDY OF AN IODINE CLOCK REACTION
PURPOSE
To investigate the kinetics of the reaction that occurs between iodide and persulfate ion. You will: (1) determine the rate law, (2) determine the numerical value of the rate constant at room temperature, (3) explore the effect of temperature on the reaction and determine the activation energy (Ea), and (4) investigate catalytic activity of selected metal ions on the reaction. INTRODUCTION

Reaction times vary from picoseconds (10-12 seconds) to years. It is an experimental challenge to design methods of studying reactions over such a wide range of rates. Reaction rates are similar to other rates: it is a change of something per unit time (like velocity - miles per hour, flow rate liters per second, etc.). In reactions the rate is measured by changes in concentrations per unit time. For example, in the following reaction (all species are in aqueous solution): S2O82- + 3 I- → 2 SO42- + I3-

(1)

the reaction rate could be determined by following a decrease in the concentrations of reactants (S2O82- or I- ) or by the increase in the concentration of the products (SO42- or I3-). However, we need to recognize that these rates of change would all produce different numbers.

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2Δ S2O2Δ I8
3
1Δ I
1 Δ SO4
That is −
= rate of reaction.
=−
=
=
Δt
3 Δt
2
Δt
Δt

(2)

We will investigate this reaction in the laboratory using the method of initial rates (consult your text if you are not familiar with this terminology) by measuring the amount of time required to produce a known amount of I3-. The presence of I3- can be easily detected using a starch indicator. In order to understand how we will accomplish the measurement of the rate of production of I3- it is important to understand some details of how this reaction proceeds. First recognize that this is an oxidation-reduction reaction with the half-reactions: reduction:

oxidation:

S2O82- + 2 e- → 2 SO423 I- → 2 e- + I3-

(3)
(4)

It will also become important to recognize that the overall reaction does not occur in a single step, but in a series of steps (or in a mechanism). This reaction can be broken down into at least two simple steps (rather than one overall reaction). One example (correct? or incorrect?) may be: slow

fast
fast

S2O82- + 2 I- → 2 SO42- + I2
I2 + I- → I3I3- + starch → blue-black starch complex

(5)
(6)
(7)

Thus, as the first reaction produces iodine, I2, it is very rapidly converted into a blue-black complex which can be easily seen. The problem with this is that the color change will occur 57

very quickly at the very beginning of the reaction. In order to produce reaction times that are easily measured, we will intervene in the normal sequence of reactions listed above in order to postpone the production of the blue-black color until a known quantity of I3- has formed. In order to do this we will use another redox reaction:

very fast

I2 + 2 S2O32- → S4O62- + 2 I-

(8)

This reaction prevents the iodine, I2, from reacting to form the blue-black starch complex. Into each of our reaction mixtures, we will place the same known quantity of thiosulfate, S2O32-, and only after all of it is consumed will the iodine, I2, be free to react further to form the blue-black starch complex. In this fashion, we will know how much iodine, I2, has been formed when the color appears. By measuring the time this takes, the rate of reaction can be found. Experimental Determination of Rate Laws

The rate of reaction is going to be determined indirectly by knowing how much thiosulfate, S2O32-, has been consumed. Reaction (8) allows us to relate this to the amount of I2 produced in reaction (5) and consumed in reaction (8):

rate of reaction = rate of consumption of I2 =

(10)

Thus by knowing the initial concentration of thiosulfate in each reaction mixture, the rate of reaction can be found by measuring the amount of time for the blue-black color to...
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