Lab 12 - Rate Properties of an Iodide Oxidation Reaction

Goal and Overview

The rate law for the reduction reaction of peroxodisulfate (PODS) by iodide:

S₂O₈²⁻(aq) + 2 I⁻(aq) → I₂(aq) + 2 SO₄²⁻(aq)

will be determined. The orders of reaction with respect to PODS and to iodide will be found by measuring rates for various concentrations of the reactants. Rates will also be measured at three additional temperatures, such that the Arrhenius equation can be used to calculate the activation energy and the pre-exponential factor for the reaction.

Objectives and Science Skills

• •
  Understand and explain common factors that influence the rates of chemical reactions.
• •
  Apply linear fitting methods to find relationships between dependent and independent variables.
• •
  Quantitatively and qualitatively evaluate experimental data to determine the reaction orders, rate constant, Arrhenius factor, and activation energy of an iodine clock reaction.
• •
  Identify and discuss factors or effects that may contribute to deviations between theoretical and experimental results and formulate optimization strategies.

Suggested review and external reading

• •
  relevant reference and textbook information on kinetics and dilutions

Background

Thermodynamics describes the equilibrium phases and concentrations of reactants and products under a given set of conditions. It gives no information on how quickly the equilibrium concentrations are reached. Chemical kinetics is the study of chemical reaction rates. Understanding how fast a reaction occurs, along with the factors that affect the rate, often helps to determine exactly how a reaction is occurring (the mechanism of the reaction). The speed with which a reaction occurs is described by the reaction rate. For a simple reaction like

a + b → c + d, 

the reaction rate can be expressed as concentration change with respect to change in time. Since the stoichiometry of the reaction above involves coefficients of 1 only, there is a simple relationship between the rates of formation of each product and the rates of consumption of each reactant. The negative signs on the reactant terms ensure that the calculated rate is positive for the forward reaction (the reactant concentration change is negative when the reaction proceeds from left to right). Reaction rates often depend on reactant concentrations. In a number of simple cases, the rate is proportional to a reactant's concentration, raised to some power. The overall dependence is the product of these terms. Here, the reaction is m^th order in a and n^th order in b; its overall order is m + n. The proportionality constant, k, which is called the reaction rate constant, depends on temperature and is given by the Arrhenius equation. Where A is a constant called the pre-exponential factor; E_a the activation energy for the reaction; R is the universal gas constant in units of J/K · mol; and, T is the absolute temperature in Kelvin. You will determine the rate law for the PODS/I^– reaction (m, n, A, and E_a). This enables the rate for any concentrations of PODS and I^– at any temperature to be calculated.

The Peroxodisulfate Iodide Clock Reaction (PODS/I^– reaction)

In this reaction, two iodide ions provide one electron each to reduce the peroxodisulfate ion, S₂O₈^2–, to form two stable sulfate ions and molecular iodine. The overall redox reaction is shown below. The rate law can be written as follows. The stoichiometric coefficients for iodide and for sulfate appear in the denominators for their rate expressions. I^– is consumed twice as fast as PODS; SO₄^2– is formed twice as fast as I₂. The reaction is m^th order in S₂O₈^2–, n^th order in I^–, and (m + n)^th order overall. The specific rate constant, k, for this reaction, as it is balanced, contains the rate's dependence on temperature. This reaction is quite slow, making its rate ideal to measure in a lab setting. It is also convenient to study because molecular iodine, I₂, is easy to detect. I₂ reacts with starch to form a deeply blue-colored complex. By adding starch to the reaction mixtures, a deep blue color is observed as soon as iodine is produced by the reaction. However, the blue color appears almost immediately and it gets extremely intense very quickly. In order to control how quickly the blue color appears, a second reaction is used. A small amount of thiosulfate ions, S₂O₃^2– is added to the initial mixture. The S₂O₃^2– ions react immediately with any I₂ present to regenerate two I^– ions. This prevents the solution from turning blue until all of the S₂O₃^2– is consumed. Once

[S₂O₃²⁻]

goes to zero, the I₂ produced will no longer be removed and the solution turns blue. By using a known quantity of S₂O₃^2–, the rate can be calculated in terms of

[S₂O₃²⁻].

Because the coefficient of S₂O₃^2– in the I₂ scavenging reaction is 2 (while the coefficient for I₂ is 1), there is a 2 in the denominator of this term. The change in S₂O₃^2– concentration is the negative of the [S₂O₃^2–]_initial because [S₂O₃^2–]_final = 0. The change in time is simply the measured time, t, from the beginning of the experiment because t_initial = 0. The initial S₂O₈^2– and I^– concentrations change very little during the reaction because the amount of S₂O₃^2– added is relatively small, so their concentrations can be calculated using the dilution equation. Use Eq. 7

Rate =

Δ[I2]

Δt

 = −

Δ[S2O32−]

2Δt

 =

[S2O32−]initial

2t

 = k[S₂O₈²⁻]^m[I⁻]ⁿ

 

 for your experiments. Each run uses the same initial concentration of S₂O₃^2–.

• •
  To determine m, use different concentrations of PODS.
• •
  To determine n, use different concentrations of I^–.
• •
  Once you have found m and n, determine the average value of the rate constant k at room temperature.

To determine A and E_a, data at different temperatures are required. By doing trials with constant [S₂O₈^2–] and [I^–] at several temperatures, you can determine k values under these conditions. The activation energy E_a and pre-exponential factor A will be determined graphically. Once m, n, E_a, and A are known, the reaction rate for any concentrations of PODS and I^– at any temperature can be calculated.

Procedure

Part 1a: Reaction Order with Respect to Peroxodisulfate Ion (PODS), m

There are five runs in this series. The following table gives the volumes of reagents that must be accurately pipetted into a clean, dry 150-mL beaker for each trial.

• 

Once you have mixed the first four reagents, initiate the experiment.

Part 1b: Reaction Order with Respect to Iodide Ion, n

There are five runs in this series. The following table gives the volumes of reagents that must be accurately pipetted into a clean, dry 150-mL beaker for each trial.

• 

Once you have mixed the first four reagents, initiate the experiment. Note: Put your results on the board and compare your results with those of others.

Part 1 Data Analysis: To Determine m and n From Your Measured Times (Parts 1a and 1b)

Part 2: Effect of Temperature on the Reaction Rate

Initial Part: Begin this immediately so that the solutions reach the temperature of the water baths.

You already have the data for run #2 of Part 1a (room temperature). Use the concentrations listed in Part 1a, trial #2 to find the rate at 0°C, 35°C, and 45°C.

• 

Final Part: Be sure to record the temperature of each bath.

Note: Put your results on the board and compare your results with those of others.

Part 2 Data Analysis: E_a and A Determination (temperature dependence of the rate)

Waste disposal: follow your TA's instructions. Nothing goes down the sink.

Results

Complete your lab summary or write a report (as instructed).

Abstract

Results

Sample Calculations

Discussion/Conclusions

Review

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