CHY116 Kinetics of the Clock Reaction

Measuring the Rate of a Chemical Reaction

Learning Goals

This experiment makes use of a version of the classic 'iodine clock reaction.' There are different versions of this reaction, but in all, the iodide ion (I-) is oxidized to form elemental iodine (I2). In the presence of thiosulfate (S2O32-), I2 is immediately reduced back to I-. However, when the thiosulfate is depleted, I2 can then react with starch, which produces an observable change , and signals the end of the reaction. By manipulating the starting conditions, and timing the length of the reaction, the various factors affecting the kinetics of the chemical reaction can be investigated, and the rate law, rate constant, and activation energy can be established.

Overview

By performing a chemical reaction under varying conditions, you will explore what factors affect the rate of a clock reaction, calculate the rate constant and reaction order.

Lecture Connections

Experimental Background

Reactions occur at varying speeds ranging from the slow rusting of iron to the extremely fast decomposition of TNT. Factors such as concentration of the reactants, temperature, catalysis, and nature of the reactants affect the rate of reactions. In this experiment, you will investigate the rate of oxidation of iodide (I¯) to iodine (I2) by the oxidizing agent hydrogen peroxide (H2O2):

H2O2 + 2 I- + 2 H+ <--> I2 + 2 H2O (Reaction 1)

The rate law for this reaction is given by:

rate =k [H2O2]x[I-]y[H+]z (Equation 1)

where exponents x, y, and z and the rate constant k are to be determined. To do so, you (as a class) will change the starting concentrations of each reactant and see how the measured rate changes.

To measure the rate of Reaction 1, you will add two more ingredients that do not affect the rate of reaction, but allow it to be studied more easily. First, a starch indicator is used to signal the presence of iodine (I2). In the presence of a starch indicator, a solution of I2 is blue/black. To delay the color change and make it easier to time, a fixed and limited amount of thiosulfate ion, S2O3 2-, will be added to the2 3

reaction. Thiosulfate ions do not react with any of the reactants of Reaction 1, but do react rapidly with the product I2 (Reaction 2):

2 S2O32- (aq) + I2 (aq) <--> 2 I- (aq) + S4O62- (aq) (Reaction 2)

Limiting blue/black clear/colorless reagent

As long as S2O32- is available, the reaction mixture will appear clear/colorless, since as soon as I2 is formed it is converted back to I¯. Once all the S2O32- is consumed, however, I2 builds up, and the solution becomes blue/black.

Note that, for a given amount of thiosulfate, the time required to use up that thiosulfate is determined by the rate at which I2 is formed, i.e. the rate of Reaction 1. The faster the rate of Reaction 1, the sooner the thiosulfate is used up and the solution changes color.

If the reaction is repeated several times and the concentrations of the reactants are varied, the rate law and the order of the reaction can be determined. (See the “method of initial rates” discussion in the text, section 12.4.)

The rate of Reaction 1 can be written as.

The initial (t = 0) concentration of H2O2 is known. We can calculate the concentration of H2O2 when the solution turns blue by recognizing that this happens as soon as ALL the S2O32- is used up. Notice that for every H2O2 that reacts, one I2 is formed, and while S2O32- is present, two S2O32- are used to convert that one I2 back into iodide. Thus, two thiosulfates are used for every one peroxide molecule, so Δ[H2O2] = 2Δ[S2O32-] . Over the time period from t = 0 to the color change, [S2O32-] goes from its initial value to zero, so Δ[S2O32-] is just equal to zero (the initial concentration of S2O32-).

Calculations

Using class data, the rate law, order of the reaction, and rate constant k of the reaction under study can be calculated.

a.) Determination of the reactant concentrations

Use the formula M1V1 = M2V2 to determine the (diluted) concentration of reactants at the start of the reaction.

[H+] is controlled by a buffer system (Chapter 14) composed of a mixture of acetic acid (CH3COOH, abbreviated HAc) and acetate ions (CH3COO-, abbreviated Ac-). The actual concentration of H+ can be calculated from

[H+ ] = 1.8 x 10-5(HAc]/[Ac-])

You can change [H+] by adding HAc, changing the ratio of acid to acetate.

b.) Determination of the rate of the reaction

The rate of the reaction is determined by the time required for the blue starch complex to form. This is set by the amount of S2O32- present.

where [S2O32-] is the initial concentration of thiosulfate and Δt is the experimentally observed time required for the solution to turn blue/black.

c.) Determination of the order of the reaction

For a reaction with a rate law in the form of: rate = k[A]x[B]y[C]z, the order of the reaction for each reactant can be found by comparing two experiments in which the concentrations of only one reactant (for example A) changes – the rest of the values will cancel, so that:

Expressing the reaction rate in terms of the rate of disappearance of H2O2 we have

If the change in concentration of H2O2 is the same for both experiments, this simplifies to

where Δt1 and Δt2 are measured experimentally. From these equations, x can be determined by taking the log of both sides (remember: log Ax = x logA):

The reaction orders for A, B, and C (or H2O2, I-, and H+) can each be calculated, and the overall order of the reaction is the sum of each reactant’s order.

d.) Determination of the rate constant

Once the orders of the reaction x, y, and z are determined, the experimental rate constant for each run is calculated by substituting known values for rate and concentration into the rate expression

rate = k [H2O2]x[I-]y[H+]z

e.) Determination of the activation energy

The activation energy is calculated from the Arrhenius equation

k = Ae- Ea / RT

where k is the rate constant from Equation 1, A is a collision constant, Ea is the activation energy (in Joules), R is the gas law constant (8.314 J/K·mol), and T is the temperature in Kelvin. Taking the logarithm of both sides gives

ln k = ln A - Ea/RT

which shows that a plot of ln(k) vs 1/T should give a straight line of slope -Ea/R and intercept ln(A). Alternatively, from values of k measured at two temperatures,

where k1 is the rate constant measured at temperature T1 and k2 is the rate constant at T2.

If the concentrations of all reactants are the same at both temperatures, then the ratio k2/k1 can be replaced by the ratio time1/time2.

In the pre-lab discussion, you and your instructor will design a series of experiments to determine the kinetics of Reaction 1. Be prepared to contribute to this discussion, both by offering ideas and by critically evaluating the ideas of others.

Note that your group will develop a specific procedure for the particular experiments you will be performing.

Procedure

Your goal as a class is to determine the rate law and activation energy for the reaction

H2O2 + 2 I- + 2 H+ <----> I2 + 2 H2O (Reaction 1)

Your instructor will demonstrate the “iodine clock reaction”, a mixture of the above chemicals plus starch and thiosulfate. (See the background information for a discussion of the relevant chemistry and calculations.) The basic recipe used is:

Into a 100-mL beaker, measure

· 10.0 mL buffer solution (0.05 M acetic acid (HAc) and 0.05 M sodium acetate (NaAc))

· 10.0 mL 0.05 M potassium iodide (KI)

· 1 drop 2% starch indicator

· 5.0 mL 0.03 M sodium thiosulfate (Na2S2O3)

· 1 drop 0.1 M EDTA to chelate any unwanted metals present (omit if using a catalyst)

· water to make a total volume of 50 mL Stir thoroughly, then add

· 10.0 mL 0.9 M hydrogen peroxide (H2O2) quickly with stirring, and start timing.

Note that the order of adding reagents is important!

Additional chemicals available:

0.3 M acetic acid (HAc)

3% ammonium heptamolybdate (catalyst)

By changing the concentrations of the reactants you can determine the rate law, and the activation energy can be obtained from measurements at two different temperatures.

You will work in groups of 3 for this experiment. As a class you should decide what reaction conditions (concentration, temperature) should be changed by each group. Each group should plan to

1.) Do the basic recipe (above) one time and report reaction time on the board.

2.) Alter one specific variable in the basic recipe and measure the reaction time. Repeat. If the two times differ significantly from each other, repeat again. Report your average time on the board.

3.) Repeat step 2 for a different variable.

Dispose of all of your solutions in the appropriately labeled waste container.

The first task of your group is to devise a detailed procedure for your experiments. Have your instructor check your procedure before you begin work. Please bring a stopwatch to lab.

Record the purpose of your experiment and your planned procedure in your lab notebook before you begin work. Record all data collected by the group in your notebook, as you do the work. When assigning tasks, give everyone equal time in all aspects of the experiment. DO NOT run two or three tasks in parallel and share results afterwards. All group members should keep their own full record of procedures and results. There should be no need for loaning or sharing lab notebooks after the post-lab discussion.

After completing your experiments, decide (as a group) what conclusions you can draw from your own results. All groups will then meet for a Post-lab discussion to tabulate, summarize, and evaluate all results. During this discussion, copy all class data into your notebook. You will need it to answer the Post-lab questions.

Pre-lab Assignment

Submit the pre-lab notebook preparation and the pre-lab questions.

Notebook Preparation

Pre-lab template preparation: In your lab notebook, prepare the following information:

A brief (2-3 sentences) objective of the lab.

A table of glassware, equipment and chemicals to be used. Include relevant properties and safety information for each chemical. Use this helpful link for online MSDS http://siri.org/msds/ to find out the hazards associated with substances you use and make in this experiment.

Several “bullet points” summarizing the tasks involved in the procedure.

Pre-lab Questions

Pre-lab Questions 1 through 5: Complete the following Pre-Laboratory Questions before you come to lab and turn them in at the beginning of the pre-lab lecture. The concepts will help you in planning and understanding your experiments. Be sure to read and study the BACKGROUND material and procedure. Keep full, legible records of your work, data, and observations in your Laboratory Notebook. You will hand in the copies of your lab work with the Report Form.

Pre-lab Questions (pdf version of Pre-lab Questions)

Your instructor will collect your answers to these questions before the pre-lab discussion, and will evaluate them as a pre-lab preparation grade. Your instructor will not accept these questions after the pre-lab discussion.

1.) Given the basic recipe in the procedure above,

a) What is the initial concentration of each reactant in Equation 1 (H2O2, I-, H+)?

(See Calculations a and b in the Background information. Remember the dilution!)

[H2O2]0 = M

[I¯]0= M

[H+]0 = M

b) What is the initial concentration of S2O32-?

[S2O32-] = M

c) What is the concentration of S2O32- when the solution turns blue/black?

[S2O32-] = M

d) What is the change in concentration of H2O2 (Δ[H2O2]) from the start of reaction until the solution turns blue/black?

Δ[H2O2] = M

e) What is the rate of the reaction (with appropriate units) if it takes 35 seconds to turn blue/black?

2.) In the calculations for reaction order and activation energy, you use the ratio of two rates, rate1/rate2. As pointed out in the Background information, as long as Δ[H2O2] is the same for both experiments, you could use time2/time1 instead. (Note that it is rate1/rate2 and time2/time1). Which of the following changes to the basic recipe would make Δ[H2O2] NOT have the same value? Circle all that apply.

a) Using different total volumes for the two experiments

b) Changing the initial concentration of H2O2

c) Changing the initial concentration of S2O32-

d) Adding a catalyst

3.) For the sake of this problem, assume Reaction 1 is second order in H+. A solution containing iodide, thiosulfate, starch, and 2.5 x 10-4 M H+ turns blue 14.7 seconds after hydrogen peroxide is added. How long will it take the same solution to turn blue if [H+] = 1.25 x 10-4 M (all other quantities the same)?

4.) For the reaction

2 NO(g) + Br2(g) → 2 NOBr(g)

the rate law will be of the form

rate = k[NO]x[Br2]y

Given the following rate data, determine the values of k, x and y. (See Calculations c and d in the Background information.)

ClockPreLabQ4Table.xlsx

x =

y =

k =

5.) If the basic recipe above takes 35 seconds to turn blue/black at 20 °C but only 15 seconds at 35 °C, what is the activation energy? (see Calculation e in the Background information.)

Experimental Videos - COVID Adaptation

(1) Basic Recipe, Experiment #1

HT demonstrating the "control" reaction, 11.1 min

(2) Experiment #2

VPT changing [KI], 7.1 min

(3) Experiment #3

Time lapse, VPT doubles [H2O2] reaction takes 60 seconds, 0.2 min

(4) HT Board Talk

2.4 min

(5) Experiment #4

HT changing [H+], 9.4 min

(6) Repeat of "control"

3.2 min

(7) Experiment #5

VPT performs "control" reaction with catalyst, 2.3 min

(8) Experiment #6

HT performs "control" at 1.0 ˙C and Board Talk, 17.1 min

References

Roberts, Juian; Postma, James The rate of a Chemical Reaction Chemical Kinetics Laboratory Studies in General Chemistry from General Chemistry in the Laboratory, Third Edition, W.H Freeman and Company, New York, 1991.

Grading Rubric

(pdf)

Post Experimental Analysis complete the REPORT FORM on your own.

Report Form

(pdf version of Report form) Name

CHY 116: Kinetics Clock Reaction

Carry out any needed calculations in your lab notebook and transfer the answers to the spaces provided. Beside the answer, give the page number where your instructor can find the calculation.

1.) How does the concentration of iodide ion change during the initial portion of the reaction, while there is still unreacted thiosulfate present? Can you assume that the concentration of iodide ion was essentially constant during the timed portion of the experiment? Explain.

2.) For the reaction using the basic recipe,

a) calculate the percentage of hydrogen peroxide that was consumed during the timed reaction.

b) Can you assume that its concentration was essentially constant during the timed portion? Why or why not?

3.) How does the concentration of hydrogen ions change during the timed reaction? Can you assume it was essentially constant? Explain.

4.) Do we care whether the reactant concentrations were constant during the timed reaction? Why or why not?

5.) The rate law for the iodine clock reaction is of the form

rate = k [H2O2]x [I-]y [H+]z , in which rate can be expressed as ∆[H2O2]/∆t.

Using the class data, determine the values of x, y and z. Indicate the page number in your lab notebook where you carried out the calculations.

ANSWERS: x = ; y = ; z = . (PAGE # )

6.) Using data from the reaction of the basic recipe, calculate the value of k, the rate constant.

ANSWER: L/mol•s (PAGE # )

7.) Compare your value of k in Question 6 to the reported value of 0.0115 L/mol•s at 25 oC. If your value differs by more than 20 %, discuss possible reasons for the discrepancy. Do not blame blunders in lab -- think about reasons that your value might differ from the reported value.

8.) Consider the following proposed mechanisms for Reaction 1.

Mechanism A (one step):

H2O2 + 2 I- + 2 H+ <----> I2 + 2 H2O (one step)

Mechanism B (three steps)

H2O2 + I- --> OH- + HOI (rate-determining step)

H+ + OH- <---->H2O (fast)

HOI + H+ + I- --> I2 + H2O (fast)

Mechanism C (three steps)

H2O2 + I- --> OH- + HOI (fast)

H+ + OH- <----> H2O (fast)

HOI + H+ + I- --> I2 + H2O (rate-determining step)

a) Is mechanism A consistent with the class results? Why or why not?

b) Is mechanism B consistent with the class results? Why or why not?

c) Is mechanism C consistent with the class results? Why or why not?

9.) Use data from the basic reaction at room temperature and at a higher temperature to compute Ea, the Arrhenius activation energy for this reaction (in kJ/mol).

Turn in:

1.) All pages of this Report Form, and

2.) A typed brief (5-10 sentences) summary of this experiment. The following should be included in the summary:

1. A topic sentence(s) describing the goal of the experiment.

2. A description of the results/outcomes of the experiment – don’t include procedural details or intermediate calculations. Include uncertainties associated with numerical results. If literature values are available, include a comparison between the experimental results and literature values. Include % error if a literature value is available.

3. A description of source(s) of error for this experiment. (This is usually the uncertainty associated with whichever instrument was used to record the data.) Do not cite human error.

4. A conclusion sentence.