Kinetics: The Reaction of I^- with H₂O₂ Using Initial Rate Methods

This experiment was developed by Profs LH Berka and NK Kildahl, WPI, based on a commercially available experiment.

1 lab period; work in pairs. Complete the Preparation before laboratory.

• To explore the kinetics of a slow redox reaction
• To determine the rate law
• To explore the effect of a catalyst on the reaction rate

Kinetics and Thermodynamics. The distinction between kinetics and thermodynamics is an important one. Kinetics is the study of the rate at which (and the path by which) a chemical reaction occurs. Thermodynamics deals with the extent to which the reaction occurs. There is no necessary connection between the rate, measured by a rate constant k (see below), and the extent, measured by an equilibrium constant K_eq. Thus, a reaction in which 99.9% of the reactant molecules are eventually consumed may take years to reach that stage. Conversely, it may take just milliseconds for 0.1% of a reactant to be converted to products, but with no further reaction over an extended period of time. The former is an example of a slow reaction (small k) that occurs essentially to completion (large K_eq); the latter is a fast reaction of limited extent.

The Differential Rate Law. For the general reaction (1), the rate is expressed as the change of concentration with respect to time, t, of any reactant or product.

Since A is consumed and F is produced by (1), their derivatives have opposite signs. Thus we introduce a minus sign in front of d[A]/dt to ensure a positive rate. Also, unless a = f, -d[A]/dt and d[F]/dt are different. In general, the rate of reaction is a function of the concentrations of reactants and products, as indicated in (3).

k is the rate constant of the reaction, and [A], [B], etc., are the molar concentrations of reactants and products in (1). The exponents m,n... appearing in (3) are usually either positive or negative integers or zero.

Because (3) is a differential equation, it is called the differential rate law for (1). The goal of a kinetics study is to determine the differential rate law for the reaction (i.e., the values of the exponents m, n, s, and v) and the numerical value of the rate constant k at each of several temperatures. Suppose that appropriate kinetics experiments carried out on (1) give m = l, n = 2, s = 0, and v = -1. The differential rate law for the reaction is then

The exponents m, n, etc. are called orders. We say that the reaction is first order in A, second order in B, and inverse first order in F. It is important to realize that there is no necessary connection between the reaction orders m, n, s, and v, in (3) and the stoichiometric coefficients a, b, d, f, in (l).

The Method of Initial Rates. A straightforward way to determine experimentally the specific form of the rate law (3) is to measure the initial rate, -(d[A]/dt)₀, as a function of each of the initial concentrations [A]₀, [B]₀, etc, and from the data deduce the reaction orders. Suppose that we found that -(d[A]/dt)₀ quadrupled when we doubled the [A]₀ while keeping the other concentrations constant. We would conclude that m = 2. Similarly, if -(d[A]/dt)₀ decreased by a factor of 2 when [B]₀ was cut in half while keeping other concentrations the same, we would conclude that n = 1.

This Experiment. We will study the initial rate at which hydrogen peroxide, H₂O₂, reacts with iodide ion, I- , and hydrogen ion, H⁺, to give triodide ion, I₃^-, and water in aqueous solution. By varying the initial concentrations of the participants, we will deduce the differential rate law for reaction (8).

To obtain the rate, we will measure the time required for a definite, known amount of H₂O₂ to be consumed. This can be accomplished by adding a known amount of S₂O₃^2- to the system. As reaction (8) proceeds, the I₃^- produced will react with S₂O₃^2- according to (9) and will be converted back to I^-.

As long as (9) occurs very rapidly relative to (8), [I^-] will not change until the S₂O₃^2- is gone. How do we signal that S₂O₃^2-is gone? I₃^- forms a blue complex with starch. So we put a little starch in the solution, too. As soon as S₂O₃^2- is used up, I₃^- will accumulate and immediately produce a blue color with the starch. If we know how much S₂O₃^2- we put in, we can calculate the amount of H₂O₂ used up in the time it took for the blue color to appear, because the combined stoichiometries of (8) and (9) tell us that 1 mole H₂O₂ is equivalent to 2 moles S₂O₃^2-. Once we know how much H₂O₂ was consumed, we can estimate the initial reaction rate as -D[H₂O₂]/Dt.

A sketch of our procedure follows.

1. Begin a run at t = 0 with known [I^-]₀, [H⁺]₀, [S₂O₃^2-]₀, and [H₂O₂ ]₀ such that [S₂O₃^2-]₀ << [H₂O₂]₀. The subscript "0" means initial concentration. The run begins when the final reagent, H₂O₂, is added.
2. Note the time, t₁, at which the blue flash appears. At this point,
   • S₂O₃^2- is used up
   • An amount of H₂O₂ = n₀(S₂O₃^2-)/2 has been used up in the time, t_l. Here n₀ = the moles of S₂O₃^2- initially present.

   Calculate the initial rate, -D[H₂O₂]/Dt, of the reaction.

The above procedure requires calculation of concentrations in the reaction mixtures. Each reaction mixture will be prepared by mixing aliquots of stock solutions of I^- and S₂O₃^2- with buffer solution and water to give a total volume of at least 17 and at most 19 mL. Reaction is then initiated by adding enough H₂O₂ stock solution to give a final total volume of 20.0 mL. If [A]_stock symbolizes the concentration of the stock solution of reagent A, V_A the volume of the aliquot of that stock solution used, and V_T the total volume of reaction mixture, the concentrations of reagents in the reaction mixture can be calculated using the following equations.

In these equations, [H₂O₂]₀ and [I-]₀ are the initial concentrations of these reactants in the reaction mixture (t = t₀ = 0).

Focus Questions

1. For each kinetics run, construct a table giving
   • The value of [I-]₀ (see Background section for method of calculation).
   • The value of [H₂O₂]₀.
   • The change in [H₂O₂] during the measured time.
   • The time required for the blue flash to appear.
2. What fraction of the total H₂O₂ reacts during the measured time interval in run 1?
3. By choosing appropriate runs, determine the order of the reaction in H₂O₂.
4. By choosing appropriate runs, determine the order of the reaction in I^-.
5. Calculate the rate constant for the reaction.

• six small beakers (100 or 150 mL)
• six 5-cm watch glasses
• graduated cylinders (5 and 25-mL)
• three 50-mL Erlenmeyer flasks
• thermometer
• two 2-mL graduated pipets
• two 5-mL graduated pipet
• small syringe pipet pump
• large syringe pipet pump
• 100-microliter pipettor
• stopwatch (optional)
• 50 mL 0.200 M KI
• 175 mL buffer solution
• 8 mL 0.2% starch solution
• 16 mL standard Na₂S₂O₃ ( 0.02 M)
• 25 mL 0.4 M H₂O₂

Note to instructor: Click here for recipes for preparation of solutions.

Hydrogen peroxide is an oxidizing agent; avoid contact with skin and clothing.

Record all data in your notebook. Obtain required glassware and a stopwatch. Clean the glassware thoroughly using brushes and Alconox detergent (pipets cannot be cleaned this way, but can be rinsed). Remove all soap with tap water, then rinse with distilled water and dry thoroughly.

Preliminary Activities

1. Label clean dry beakers, one for each required solution: (I^-, buffer, starch solution, S₂O₃^2-, distilled water, H₂O₂, (Table 1)).
2. Obtain appropriate amounts of the solutions (Table 1) and cover the supply beakers.
3. Record the exact concentrations of the I^-, Na₂S₂O₃, and peroxide solutions.

Kinetics Runs

Five kinetic runs will be carried out using the volumes of stock solutions in Table 1. The procedure for carrying out Run 1 will be outlined in detail. The others are carried out in a similar manner.

Run 1:

1. To a clean, dry 50-mL Erlenmeyer flask, add the quantities of reagents 1-5 indicated under run 1 in Table 1, and mix.
2. With a graduated pipet, add 1.0-mL of H₂O₂ stock solution to the flask, swirl, and simultaneously begin timing. RAPID, EFFICIENT MIXING IS ESSENTIAL.
3. At the blue flash
   • record the elapsed time to the nearest second (t = t_l).
   • Record the temperature.
   • IMMEDIATELY discard the reaction solution down the drain and rinse the reaction flask with tap water. If you allow the solution to remain in the flask, a lot of I₂ will form, making the flask difficult to clean.

Carry out Run 1 in triplicate. Then carry out runs 2-5 in similar fashion, also in triplicate.

If time allows, design a procedure to determine the order in H⁺.

If time allows, design a procedure to determine the temperature dependence of the rate.

Clean-up. When you have finished all of your work:

• Dispose of left over solutions and reaction mixtures as indicated below.
• Clean Pasteur pipets, graduated pipets, vials, and other glassware by recommended procedures, shake off excess water, and place in the drying oven.
• Clean and return all borrowed equipment to the instructor before leaving lab.
• Clean up your work area before leaving lab.

Disposal Methods

Discard left-over solutions in the sink; flush down the drain with water.