Kinetics: The Reaction of I^- with H₂O₂ (Version 2)

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 find out how the concentration of a reactant varies with time
• To define how reaction rate is to be defined in terms of this curve
• To determine how reaction rate is affected by the concentrations of reaction participants
• 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 Concentration-time Curve. Consider the generic reaction in (1).

As reaction proceeds from some starting point, A and B are consumed and D and F are produced by (1). In this experiment, we will be concerned with how rapidly this transformation takes place; how the concentration of one of the reactants, say A, varies with time; how rate is expressed mathematically in terms of the resulting concentration-time curve; how the rate is influenced by changing the amounts (concentrations) of reactants; how rate is influenced by changing the concentration of a substance that does not appear in the equation; and how rate is influenced by temperature.

A plot of the concentration of a reactant (or product) versus time is called a concentration-time curve. Concentration-time curves for all molecular processes have a characteristic shape. One of our objectives will be to determine the appearance of such a curve. We will find throughout further studies that the characteristic concentration-time profile comes up repeatedly in a variety of contexts.

How is Rate Expressed Quantitatively? With a concentration-time profile at hand, we will discuss and agree upon how such a curve can be used to express a reaction rate. From our discussion will emerge an understanding of how the reaction rate changes as reaction proceeds; and how it depends upon the concentration of the species being measured.

Dependence of Rate on Concentration of Participants. Having agreed on a mathematical treatment of rate, we will be free to explore the manner in which the concentration-time curve is affected by the concentrations of other participants, those whose concentrations are not being monitored. The result of these efforts will be an expression called a rate law.

Dependence on Temperature. Finally, we will explore the effect of temperature on the concentration-time profile.

This Experiment. We will study the rate at which hydrogen peroxide, H₂O₂, reacts with iodide ion, I-, to give I₃^-, and water in aqueous solution. The equation for the reaction is (2).

To accomplish this we will measure the concentration of hydrogen peroxide at various times during the course of reaction, keeping the concentrations of the other two reactants, I- and H+, constant. It is simple to keep H+ concentration constant by running the reaction in the presence of a buffer. We are left with two questions: how do we keep [I^-] constant; how do we monitor [H₂O₂]? Using some neat chemistry, we can do both.

Thiosulfate, S₂O₃^2-, reacts quantitatively with I₃^- according to (3).

If we have S₂O₃^2- present in the solution while (2) is proceeding, I₃^- produced by (2) will be converted back to I^- by (3). As long as (3) occurs rapidly relative to (2), [I^-] will not change until the S₂O₃^2- 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^- accumulates and produces 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 (2) and (3) tell us that 1 mole H₂O₂ is equivalent to 2 moles S₂O₃^2-.

A sketch of our procedure for determining the concentration-time profile follows.

1. Begin a run at t = 0 with known [H₂O₂ ]₀, [I^-]₀, [H⁺]₀, [S₂O₃^2-]₀ such that [S₂O₃^2-]₀ <<[H₂O₂]₀. The subscript "0" means initial concentration.
2. Note the time, t_l, at which the blue color 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.

   Calculate [H₂O₂] in the solution at time t_l. As soon as the blue color appears and we note the time,

3. Quickly add another small, known amount of S₂O₃^2-. What will happen to the color?
4. Again note the time, t₂, at which blue appears.
   • From the amount of S₂O₃^2- added at time t₁, calculate [H₂O₂] in the solution at time t₂.
5. Continue 3 and 4 to obtain [H₂O₂] remaining at each of several times during 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 and water to give a total volume of 19.0 mL. Reaction is then initiated by adding 1.0 mL of a H₂O₂ stock solution, to give a final total volume of 20.0mL. If we let [A]_stock symbolize the concentration of the stock solution of reagent A, V_A be the volume of the aliquot of that stock solution used, and V_T,avg be the average total volume of reaction mixture (an average is required because total reaction mixture volume changes slightly each time a small volume of thiosulfate solution is added), the concentrations of reagents in the reaction mixture can be calculated using the following equations.

In these equations,

Finally, since we add S₂O₃^2- stock solution to the reaction mixture in 1-mL increments,

Syringe Use. The proper use of the syringe in this experiment is given below. Rinsing should be carried out only at the start of a lab period. Filling will be necessary several times and will have to be carried out quickly in the middle of most of the kinetic runs. Therefore, you may wish to practice both filling and injecting using distilled water.

1. Rinsing - fully depress the plunger and draw up about 1 mL of solution into the syringe. Hold the syringe horizontally and withdraw the plunger to about one mm above the 1.00 mL calibration mark. Shake the syringe to rinse the inside of the barrel and bottom of the plunger with the solution. Depress the plunger to expel the solution into the sink. Rinse the syringe four times.

2. Filling - Fully depress the plunger, then draw about 0.05 mL air into the syringe. Place the tip of the syringe below the surface of the Na₂S₂O₃ solution in the supply beaker and draw slightly more than 1.00 mL of solution into the syringe. Withdraw the tip of the syringe from the solution and dislodge air bubbles by flicking the syringe with your finger. Depress the plunger until the lower part of the liquid meniscus is exactly at the 1.00 mL calibration mark. Touch the tip of the syringe to the inside of the beaker.

3. Injecting - To inject a 0.10-mL increment of Na₂S₂O₃ solution, place the tip of the syringe just inside the mouth of the receiver flask, and depress the plunger until the lower part of the meniscus is exactly even with the appropriate calibration mark. Touch the tip of the syringe to the inner side of the flask, then remove the syringe. Refill the syringe when necessary.

Focus Questions

1. For each kinetics run, construct a table giving
   • The value of [I-]₀ (see Background section for method of calculation).
   • Values of t (including t = 0, there should be 11 values for each run).
   • The value of [H₂O₂] for each t (see Background Section for method of calculation).
2. For each run, make appropriate plots to determine the order of reaction with respect to H2O2. To do this may require that you do further calculations.
3. Extract observed pseudo-order rate constants from your plots.
4. Make an appropriate plot of these rate constants to determine the order of the reaction with respect to I^- and H⁺.

• 2 100-mL Beakers
• 2 5-cm Watch glasses
• Graduated cylinders (5 and 25-mL)
• 4 25-mL Erlenmeyer flasks
• thermometer
• 3 2-mL graduated pipets
• Syringe pipet pump
• 1-mL plastic syringe
• 22.5 mL 0.200 M KI
• 30.0 mL buffer solution
• 5.0 mL 0.2% starch solution
• 7.5 mL standard Na₂S₂O₃ ( 0.2 M)
• 20.0 mL 0.4 M H₂O₂

Record all data in your notebook. Clean the required 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^-, S₂O₃^2-, H₂O₂, buffer, starch solution, distilled water (Table 1)).
2. Obtain appropriate amounts of the solutions (Table 1) and cover the supply beakers.
3. Record the exact concentrations of the Na₂S₂O₃, I^-, and peroxide solutions
4. Fill syringe with solution.

Concentration-time Curve for H₂O₂; Dependence of Reaction Rate on [H2O2]

Several kinetic runs will be carried out using the volumes of stock solutions given in Table 1. The procedure for carrying out Run 1 to determine a concentration-time curve will be outlined in some detail. The others are carried out in a similar manner.

Run 1:

1. To a clean, dry 25-mL Erlenmeyer flask, add the quantities of reagents 1-5 indicated under run 1 in the Table 1, and mix.
2. With a graduated pipet, add 0.1-mL of H₂O₂ stock solution to the flask, swirl, and simultaneously begin timing.
3. At the first appearance of blue color
   • record the time to the nearest second (t = t_l)
   • add 0.10 mL of S₂O₃^2- solution from the syringe. What happens?

     NOTE: At t₁ you are adding the second increment of S₂O₃^2- while recording the time required for the first increment (present initially - see step 7b) to react.

4. Continue as in 3). At each appearance of blue color,
   • Record the time;
   • Add a 0.10 mL increment of S₂O₃^2-, until you have added a total of 1.00 mL of S₂O₃^2- (includes the 0.1 mL present initially)
   • Record the temperature.

Use equations (4) and (5) to calculate the [H2O2] present at each measured time. Plot the concentration-time curve and express in words what the curve tells you.

Discuss among yourselves the following questions:

1. What quantity or quantities are needed to numerically define the reaction rate?
2. How can these quantities by determined from the concentration-time curve.
3. Is reaction rate constant as reaction proceeds? If not, how does it vary?
4. At what time is the reaction rate largest?
5. How much time is required for [H2O2] to decrease from its initial value to 1/2 of that? How much more time is required for it to fall to 1/4 the initial value?
6. How can you use the concentration-time profile to establish the mathematical relationship between the reaction rate at a particular time and the concentration of H2O2 at that time? Do this; write an equation to express the relationship.

Before proceeding, discuss and predict what you think the mathematical relationship between rate and the concentrations of other participants should be.

Dependence of Reaction Rate on [I-]. Carry out Runs 2-4 in the same fashion as Run 1. For each run, plot the concentration-time profile. Use your four concentration-time profiles to determine the dependence of rate on the iodide ion concentration.

Dependence of Reaction Rate on [H+]. If time allows, design a procedure to determine the order in H⁺.

Dependence of Reaction Rate on Temperature. If time allows, design a procedure to determine the temperature dependence of the rate.

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

• Clean all glassware with Alconox and water and shake dry.
• Rinse Pasteur pipets and graduated pipets by inserting the nozzle of your wash bottle in the top end and squirting water through. Aspirate dry.
• Return all borrowed equipment to the instructor before leaving lab.

Disposal Methods

Preparation Questions