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Lab Manual | Clock Reaction

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Reaction Kinetics: The Iodine Clock Reaction

Introduction

The “clock reaction” is a reaction famous for its dramatic colorless-to-blue color change, and is often

used in chemistry courses to explore the rate at which reactions take place. The color change occurs

when I2 reacts with starch to form a dark blue iodine/starch complex. The ability to record the time

at which the blue complex appears allows the rate of reaction to be determined accurately with a

stopwatch.

In this experiment, the rate law for a reaction is determined using the method of initial rates. The

effect of concentration on the rate of this reaction is determined by measuring the initial reaction rate

at several reactant concentrations. You will also examine the effect of a catalyst on the reaction rate.

Lastly, you will investigate the effect of temperature on the rate of this reaction, which will allow

you to determine the activation energy.

The Clock Reaction

The primary reaction to be studied is the oxidation of the iodide ion by the bromate ion in aqueous

solution:

Equation 1

This reaction will be run in the presence of a known amount of S2O3 2-

(thiosulfate), which reacts

very rapidly with I2. As long as S2O3 2-

is present, I2 is consumed by S2O3 2-

as fast as it is formed.

This competing reaction prevents the I2 produced from our reaction of interest from reacting with

starch, so no color change is observed until the thiosulfate is completely used up. The "clock"

reaction is the reaction of a very small amount of S2O3 2-

(thiosulfate) with the I2 produced in the

primary reaction:

Equation 2

The “clock” reaction will signal when the primary reaction forms a specific amount of I2. The

amount of I2 formed before the color change can be calculated from the known amount of S2O3 2-

added using the molar ratio in Equation 2. To find the rate of Equation 1, the change in the

concentration of I2 is monitored over time. Below, [I2] is the change in the concentration of I2, and

t represents the change in time:

Equation 3

Recall that:

BC CHEM& 162

Lab Manual | Clock Reaction

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Equation 4

As soon as all of the S2O3 2-

ions have reacted, the I2 still being formed (Equation 1) starts to

accumulate and reacts with starch. Starch serves as an indicator to help us “see” the I2, since the

interaction between starch and I2 forms a blue starch-iodine complex. Thus, "∆t" is simply the time

elapsed between mixing the reagents and the appearance of the blue color. Because the S2O3 2-

ion

concentration in the reaction mixture is known, you can calculate "∆[I2]" using the stoichiometry of

the “clock” reaction. Since the same amount of S2O3 2-

should be added to each run, ∆[I2] should also

be the same for each run. The amount of time for the appearance of the blue color varies with initial

reactant concentrations, with temperature, and in the presence of catalyst, so t is not constant.

Notes Regarding the Initial Concentrations, the Dilution Formula, and Ion Molarities:

The initial concentrations of reactants are calculated for the moment at which they are mixed. At

that time, the solutions have mutually diluted each other (raised the volume of total solution, with or

without adding moles of the solute), but have not yet started disappearing from solution (via

reaction). For each ion in solution, a new molarity must be calculated that takes into consideration

the new total volume of the solution, and the other ions that were added. The concentrations and

volumes of the reactants which should be used are given in Table 1. Care should be taken to record

the concentrations of the actual stock solutions provided. Also, care should be taken to use the

volumes indicated in Table 1, and to note the exact volumes used in each trial to the proper

precision.

To determine the concentration of an ion in solution, consider the stoichiometric relationship

between the ionic compound and the number of ions formed in solution. For example, a 0.20 M

solution of KI releases 0.20 M I - and 0.20 M K

+ ions. The situation would be different if the source

of I - was CaI2, since 2 moles of I

- ions would be released for each mole of CaI2 that dissociates.

Calculate the concentration of each reactant after combining the solutions, but before the chemical

reaction begins. Note that if all volumes of solution used were precisely as indicated in Table 1, and

if volumes were additive, the total volume of the reaction mixture for each trial would be 10 mL. To

simplify both the calculations and the experimental procedure, the concentrations of the reactants in

the reaction mixture should be calculated as if volumes were additive. Thus, in the dilution formula,

M1V1 = M2V2, V2 is approximately 10 mL, and V1 is the volume of the individual solution added to

the mixture.

Preliminary Calculations Involving the "Clock” Reaction

Using the dilution formula, the concentration of S2O3 2-

in the mixture is 2.0 x10 -4

M. According to

the stoichiometry of the clock reaction in Equation 2, the number of moles of I2 is one-half the the

number of moles of S2O3 2-

. The blue color will appear when 1.0 x10 -4

M of I2 (with two significant

figures) has been formed by the primary reaction. This number remains constant in all Runs, and so

provides ∆[I2] in all of rate calculations.

All Runs: ∆[I2] = 1.0 x 10 -4

M I2

Reaction Rate

The rate law for Equation 1 will be determined by measuring the initial rate of reaction with varying

initial reactant concentrations. The concentration of S2O3 2-

in the reaction mixture is very small

BC CHEM& 162

Lab Manual | Clock Reaction

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compared to the other reactants present, such that the measured rate is the initial rate of the reaction.

The rate law for Equation 1 is given below:

Equation 5

In Equation 5, k is the rate constant, and x, y, and z represent the order of I – , BrO3

– , and H

+ ,

respectively. These values will be determined experimentally from your data.

Note: When the solution turns blue, only a very small percentage of the reactants have been used

up. Thus, the initial concentrations of reactants can be used in the rate law with very little error.

Typically, the error is less than 5%.

Effect of Concentration on the Reaction Rate: Finding the Rate Law

The rate law is determined using the method of initial rates. The following example will illustrate

how to find a reaction order using the method of initial rates.

Example: The following data was obtained for the reaction: A + B  C

Experiment [A], M [B], M Rate (M/s)

1 0.020 0.10 1.20

2 0.030 0.10 1.80

3 0.030 0.25 11.25

The general rate law for this example is

Rate = k[A] x [B]

Since [A] changes between Experiment 1 and 2, while [B] remains constant, the order for A is

obtained by taking the ratio of the rates from these two experiments:

Since k is constant at a given temperature and [B] y is constant for Experiments 1 and 2, the equation

simplifies to:

or 1.50 = 1.5 x

Thus, x = 1 for this example.

Unfortunately, experimental results are not usually that "clean", and a more sophisticated method is

needed to find x. Mathematically, solving for exponents requires the use of logarithms. Taking the

log of both sides of the equation above yields:

Rearranging this equation to solve for x yields

.10][k[0.020]

.10][k[0.030]

M/s1

M/s

801 =

. ÷÷ ø

ö çç è

æ =

[0.020]

[0.030]

M/s1

M/s

801

( ) ( )1.5logx1.50log ×=

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Lab Manual | Clock Reaction

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x = = 1.0 ≈ 1

Experiments 2 and 3 may then be used to find the order for B, as shown below

By cancelling out the common terms and dividing the rate and concentration values, we obtain

6.25 = 2.5 y

Taking the log of both sides and rearranging to solve for y gives

y = = 2.0 ≈ 2

Effect of Temperature on the Reaction Rate: Determination of Ea

The Arrhenius equation describes the relationship between the rate constant (k), the Kelvin

temperature (T), the activation energy (Ea), and the frequency factor(A) :

ln k = - + ln A Equation 6

y = m x + b

The reaction is run at four different temperatures, and for each, the rate constant is calculated. From

these data, construct an Arrhenius plot of ln k vs (1/T) by mapping the four data points onto the

equation of a line, as shown under Equation 6. Use the slope of the Arrhenius plot to determine the

value of the activation energy, Ea.

Effect of a Catalyst on the Reaction Rate

A comparison of the reaction rate with and without a catalyst will demonstrate catalytic action.

*Waste handling: Keep a large beaker at your station to collect the clock reaction waste. When

you have completed all trials, pour the contents of the waste beaker into the waste container in the

fume hood.

Safety Precautions

CAUTION: I2 is toxic and corrosive! It can damage eyes, skin, and clothes on contact. It

is readily absorbed through skin and harmful if inhaled in high concentrations.

Disposal

Dispose of all chemicals and solutions in the appropriately marked waste container in the fume hood.

( ) ( )1.5log 1.50log

.10][k[0.030]

.25][k[0.030]

M/s

801

2511 =

( ) ( )2.5log 6.25log

è ç

ø ÷

BC CHEM& 162

Lab Manual | Clock Reaction

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Procedure

It is recommended that you start the HOT water bath for part II immediately.

Make sure that the necessary glassware is clean and dry.

Label your glassware.

Record the concentrations of all stock solutions provided in your data table.

Record experimental values with the units and significant figures appropriate for the observation.

Use the precision of the measuring device to get the correct number of significant figures.

PREPARE SOLUTION A & SOLUTION B FOR ONLY ONE TRIAL AT A TIME.

Part I: Effect of Concentration on Reaction Rate

1. Prepare solutions A and B separately for Run 1 in TABLE 1 below, using the stock solutions

provided.

2. Transfer solution A to a clean, dry, appropriately sized beaker. Add a magnetic stir bar to the

beaker, place the beaker on the magnetic stirplate, and adjust the speed of the stirrer to obtain a slow

but steady speed.

3. Simultaneously

* Add solution B to the reaction beaker as rapidly as possible.

* Begin timing as solutions A & B are combined.

4. Watch the solution continuously, when the solution initially turns blue-black:

* Note and record the elapsed time.

* Note and record the temperature of the solution.

5. Use forceps to remove the stir bar from the beaker. Between trials, rinse the beaker and the stir

bar with tap water, and then DI water.

6. Repeat the steps above for Runs 2 – 5 in Table 1 below.

Part II: Effect of Temperature on Reaction Rate

1. Prepare a cold (0 °C) water bath, a cool (10 °C) water bath, and a hot (40 °C water bath). Select

appropriately sized beakers for your temperature baths.

For each water bath:

1. Prepare solutions A and B separately for Run 1 in TABLE 1 below, using the stock solutions

provided.

2. Allow the solutions to remain in the water bath (separately, not yet mixed!) until their

temperatures have equilibrated (about 10 minutes), and measure the temperature of Tube A.

3. Transfer solution A to a clean, dry beaker. Add a magnetic stir bar to the beaker, place the beaker

on the magnetic stirplate, and adjust the speed of the stirrer to obtain a slow but steady speed.

4. Simultaneously

* Add solution B to the reaction beaker as rapidly as possible.

* Begin timing as solutions A & B are combined.

5. Watch the solution continuously, when the solution initially turns blue-black:

* Note and record the elapsed time.

* Note and record the temperature of the solution.

6. Use forceps to remove the stir bar from the beaker. Between trials, rinse the beaker and the stir

bar with tap water, and then DI water.

BC CHEM& 162

Lab Manual | Clock Reaction

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Part III: Effect of Catalysis on Reaction Rate

1. Prepare solutions A and B separately for Run 1 in TABLE 1 below, using the stock solutions

provided.

2. Add one drop 0.5 M ammonium molybdate to solution B prior to mixing.

3. Transfer solution A to a clean, dry beaker. Add a magnetic stir bar to the beaker, place the beaker

on the magnetic stirplate, and adjust the speed of the stirrer to obtain a slow but steady speed.

4. Simultaneously

* Add solution B to the reaction beaker as rapidly as possible.

* Begin timing as solutions A & B are combined.

5. Watch the solution continuously, when the solution initially turns blue-black:

* Note and record the elapsed time.

* Note and record the temperature of the solution.

6. Use forceps to remove the stir bar from the beaker. Between trials, rinse the beaker and the stir

bar with tap water, and then DI water.

TABLE 1: Intended Composition of Reaction Mixtures

Solution A Solution B

Run

No. 0.2%

Starch,

drops

0.001 M

Na2S2O3,

0.01 M

KI, mL

H2O, mL 0.04 M

KBrO3, mL 0.1 M HCl,

1 4 drops 2 2 2 2 2

2 4 drops 2 4 0 2 2

3 4 drops 2 2 0 4 2

4 4 drops 2 2 0 2 4

5 4 drops 2 1 0 3 4

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Lab Manual | Clock Reaction

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Data

TABLE 2: Experimental Data

Reagents

Solute: Starch Na2S2O3 KI KBrO3 HCl

Concentration:

Part I

Solution A Solution B Time Temperature

Reaction

Mixture

Starch Na2S2O3 KI H2O KBrO3 HCl

Run 1

Run 2

Run 3

Run 4

Run 5

Part II Solution A Solution B Time Temperature

Reaction

Mixture

Starch Na2S2O3 KI H2O KBrO3 HCl

Run 1

Part III Solution A Solution B Time Temperature

Reaction

Mixture

Starch Na2S2O3 KI H2O KBrO3 HCl

Run 1

BC CHEM& 162

Lab Manual | Clock Reaction

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Calculations

Perform the following calculations in your lab notebook.

1. Use the dilution formula to calculate the concentration of S2O3 2–

, I – , BrO3

– , and H

+ , for each

Trial, and record these values in Table 1: Calculations. Show a sample calculation below for

Trial 1. Respect significant figures based on the metric equipment used and during calculations.

Show the units. These values will be determined experimentally from your data.

2. Convert the time for each trial to seconds. Calculate the rate of reaction (∆[I2]/∆t) for each trial and record these values in Table 1. Recall, ∆[I2] may be approximately 1.0 x 10

-4 M I2 for every

trial. Show a sample calculation below for Run 1. Respect significant figures based on the

metric equipment used and show the units.

3. Use Trials 1 and 2 to determine x, the order of I–. Show calculated value with significant figures and then state the order for I

– , rounding the value for x to the nearest whole

number.

4. Use Trials 1 and 3 to determine y, the order of BrO3 – . Show calculated value with significant

figures and then state the order for BrO3 – , rounding the value for y to the nearest whole

number.

5. Use Trials 1 and 4 to determine z, the order of H+. Show calculated value with significant figures and then state the order for H

+ , rounding the value for z to the nearest whole

number.

6. Calculate k for each of Trials 1 to 4, and calculate the average of these values.

7. Write the complete Rate Law, including the experimentally numerical value of k, x, y, and z. Use this equation to calculate an expected rate for Trial 5, and compare it to the experimentally

observed rate.

8. Calculate the Kelvin Temperature for each Trial. Using the data from part II with the data for Trial 1 from Part I, prepare an Arrhenius plot with ln k on the y-axis and (1/T) on the x-axis

with Excel. (Note that "T" needs to be expressed in Kelvin rather than °C.) This plot should

have four points, the original room temperature data for Trial 1, and the runs you carried out at

the three additional temperatures. Print a full page graph with the straight line equation showing

the slope and intercept. (Refer to the Excel Tutorial from earlier in the class if necessary.) The

graph must have properly labeled axes and a descriptive title. Attach the graph to this report.

9. Use the Arrhenius plot to calculate the activation energy, Ea, for this reaction, and show the calculation below. Respect significant figures based on the linear regression analysis and show

units. Note that the slope of this Arrhenius plot is equal to Ea/R, where the gas constant, R, is

8.314 .

mol ×K

BC CHEM& 162 Name _______________________________

Lab Manual | Clock Reaction Section _______

Page 9 of 11

Report Sheets Reaction Kinetics: The Iodine Clock Reaction

Calculations

In the tables below, neatly copy your calculated values.

Sample Calculations In the space provided below, neatly write a copy of calculations 3, 4, 5, 6, and 9 (showing all work).

TABLE 3: Calculations

Part I

Reaction

Mixture

Time Rate Temperature

Run 1

Run 2

Run 3

Run 4

Run 5

Part II

Reaction

Mixture

Time Rate Temperature

Run 1

Part III

Reaction

Mixture

Time Rate Temperature

Run 1

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Lab Manual | Clock Reaction

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Results The rate law for Equation (1) was experimentally determined to be __________________________

The activation energy (Ea) for Equation (1) was experimentally determined to be _______________

Attach a copy of the graph of your Arrhenius plot to this report.

Discussion / Follow–up Questions

1. Compare the average reaction time for Trial 1 from Part I (without a catalyst) to the reaction time of Trial 1 from Part I II (with a catalyst), include the values. How did the addition of a

catalyst affect the rate of reaction? Write in full sentences, using proper grammar.

2. Compare your average reaction time for Run 1 at room temperature to your reaction times at the colder and hotter temperatures. What effect does changing the temperature have on the

rate of reaction? Write in full sentences, using proper grammar.

3. Theoretical Error: What assumptions or approximations may have contributed to error in your results?

4. Procedural Error: What parts of the experimental procedure may have contributed significantly to error in your results?

5. Human Error: What failures on your part to effectively follow the procedure as written do you think may have contributed to error in your results? (There should not be anything to

write here, but if you did mess-up and for some reason did not redo the portion(s) of the

experiment required to obtain uncorrupted data, the mistake(s) should be identified here.)

BC CHEM& 162 Name _______________________________

Lab Manual | Clock Reaction Section _______

Page 11 of 11

Pre-Laboratory Exercise Reaction Kinetics: The Iodine Clock Reaction

1. Calculate the initial concentration of the ions below in the reaction mixture for each reaction (Runs 1 – 5) using the volumes and concentrations given in Table 1. Show a

sample calculation for Run 1.

2. During the experiment, after each trial, where will the reaction mixture and runoff of rinsing your glassware be collected ?

3. If different than your answer to question 2, after the experiment where will the reaction mixture and runoff of rinsing your glassware be collected?

Run

No. Initial [S2O3

2– ]

(M)

Initial [I – ]

(M) Initial [BrO3

– ]

(M)

Initial [H + ]

(M)