Experiment 1 The Iodine “Clock” Reaction - CHEMISTRY …proffenyes.weebly.com/uploads/2/5/2/3/25237319/experiment_1_-_the... · Rate = k [A]m [B]n where: k: is a ... log 2.00

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Experiment 1 – The Iodine “Clock” Reaction ABSTRACT

I. THE RATE LAW

1. The Effect of Initial Concentration of Reactants on Reaction Rate

2. Reaction Rates

3. Reaction Orders

4. The Rate Constant

II. THE EFFECT OF TEMPERATURE ON REACTION RATE

1. Determination of Activation Energy

2. Determination of the effect of 100C temperature increase on Reaction Rate.

Chemistry 102

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EXPERIMENT 1

REACTION RATE, RATE LAW, AND ACTIVATION ENERGY

THE IODINE “CLOCK” REACTION


PURPOSE:

I. To determine the Rate Law for the following chemical reaction:

2 I-(aq) + S2O82-(aq) I2 (aq) + 2 SO4

2- (aq)

iodide ion persulfate ion iodine sulfate ion

II. To determine the Effect of Temperature on the Reaction Rate

PRINCIPLES: I. THE RATE LAW

1. The Effect of Initial Concentrations of Reactants on Reaction Rate

The rate of a chemical reaction is a measure of how fast a chemical reaction occurs.

The Reaction Rate can be determined experimentally by measuring the change in

concentration of the reactants of products, divided by the change in time.

During the course of the reaction, the reactants are used up to produce products.

As a result, the concentration of the reactants (I- and S2O82-) decreases and the

concentration of the products (I2 and SO42-) increases accordingly.

In this experiment the Reaction Rate will be calculated by dividing the experimentally

determined increase in concentration of one of the products (elemental iodine, I2), by the

corresponding time interval:

Δ I2

Rate =

Δ t

The experimental determination of the increase in concentration of iodine (I2), during a

corresponding time interval, can be easily monitored, since the presence of even small

amounts of iodine can be detected by virtue of the intensely blue colored complex formed

between iodine and starch.

2 I-(aq) + S2O82-(aq) I2 (aq) + 2 SO4

2- (aq) (Reaction 1)

iodide ion persulfate ion iodine sulfate ion

Reacts with starch

to form

a deep-blue complex

One creative way of measuring the rate of formation of iodine is to couple the reaction in

which the iodine is formed (Reaction 1) with a much faster reaction that consumes all of

the iodine (Reaction 2)

I2(aq) + 2 S2O32-(aq) 2 I-(aq) + S4O6

2-(aq) (Reaction 2)

thiosulfate ion

Reaction 2 immediately consumes the I2 generated in the first reaction, until all of the

S2O32- (thiosulfate ion) is used up. When all of the S2O3

2- is consumed, I2 builds up and

reacts with starch to form the deep blue Starch-Iodine Complex, according to Reaction 1

given above:

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Chemistry 102

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EXPERIMENT 1




The appearance of the deep-blue complex tells us that at this point in time (tcolor),

sufficient I2 has been produced by Reaction 1 to use up all of the S2O32- (thiosulfate

ion) originally added. From this observation, we can calculate the concentration of I2

(Iodine) produced by Reaction 1, by noting that according to the stoichiometry of

Reaction 2:

1 mole of I2 reacts with 2 moles of S2O32- (thiosulfate), or

½ mole of I2 reacts with 1 mole of S2O32- (thiosulfate).

It follows that at the time the deep blue color (tcolor) appears:

[S2O32-] originally added and used up

[I2] produced =

2

2. Reaction Rates

If we know the initial concentration of the thiosulfate ion (S2O32-), that is the same for each

experiment, and remember that it is all used up when the color of the solution changes, then

we know that half the amount of I2 was also consumed in Reaction 2.

This means that the change in the I2 concentration is equal to half the starting (initial)

concentration of the thiosulfate ion S2O32-) and it remains constant throughout all the

experiments. In essence, the same amount of I2 is produced in each experiment at the time the

color changes, but it takes varying times for this to occur, since the tcolor depends on the

reaction conditions (concentration of reactants and temperature).

[S2O32-]

Δ [I2] =

2

The time of the color change (tcolor) is also the time that passed during the reaction (Δt).

It follows that the rate of any of the reactions can be calculated as:

Δ [I2] [I2] produced at tcolor

Rate = =

Δt tcolor


Chemistry 102

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EXPERIMENT 1




3. Reaction Orders

The rate of a reaction depends on the concentration of one or more of the reactants.

Consider a reaction with a pattern similar to the reaction studied in this experiment.

aA + bB cC + dD

REACTANTS PRODUCTS

As long as the reverse reaction is negligibly slow, the relationship between the Rate of

Reaction and the Concentration of Reactants can be expressed by a mathematical

expression called the Rate Law:

Rate = k [A]m [B]n where:

k: is a proportionality constant called the Rate Constant,

m is the Reaction Order with respect to Reactant A, and

n: is the Reaction Order with respect to Reactant B.

The values of the reaction orders (“m” and “n”) determine the dependence of the reaction

rate on concentration of the respective reactants. Reaction orders commonly have one of

the following values: 0, 1, -1, or 0.5. Reaction orders (such as “m” and “n”) can only be

determined experimentally and they are NOT related to the coefficients of the balanced

chemical equation (such as “a”, “b”, “c” and “d”). The examples below illustrate how the

Reaction Orders (“m” and “n”) can be determined experimentally for a reaction involving

two reactants (A and B), such as the reaction studied in this experiment.

The reaction orders (“m” and “n”) with respect to the two reactants (A and B) are

determined by measuring the initial rate for several reaction runs with varying

concentrations of one reactant (for example A) independently of the concentration of the

other reactant (B). This allows us to determine the dependence of the rate on the

concentration of [A] and the numerical value of the Reaction Order with respect to

reactant A (“m”).

Temperature Reaction

Run

[A}

M

[B]

M

Initial Rate

M/s

Room

Temperature

1 0.0300 0.0450 4.83 x 10-6

Room

Temperature

2 0.0600 0.0450 9.66 x 10-6

x 2 Constant

x 2

Between the first two experiments (1 & 2), the concentration of [A] doubles, the

concentration of [B] stays constant and Reaction Rate doubles. It follows that the initial

rate is directly proportional to the initial concentration of [A]. The reaction is therefore of

the first order with respect to [A] and m = 1.


Chemistry 102

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EXPERIMENT 1




Experimental data seldom provide such obvious and easy to interpret numbers. Reaction

orders calculated from experimental data can be calculated by substituting two initial

concentrations and the corresponding initial rates into a ratio of the rate laws.

To calculate “m”

Rate 2 k [A]m[B]n 9.66 x 10-6 M/s k [0.0600 M]m [0.450 M]n

= =

Rate 1 k [A]m[B]n 4.83 x 10-6 M/s k [0.0300 M]m [0.450 M]n

Canceling out similar terms and doing the calculations yields: 2.00 = 2.00m

To find “m”, take the log of both sides of the equation and solve for “m”

log 2.00 0.301

log 2.00 = log (2.00m) log 2.00 = m log 2.00 m = = = 1

log 2.00 0.301

The Reaction Order with respect to [B], (“n”) can be calculated in a similar manner by

using the data obtained for Reaction Runs 3 & 4.

Temperature Reaction

Run

[A}

M

[B]

M

Initial Rate

M/s

Room

Temperature

3 0.0500 0.0150 1.33 x 10-6

Room

Temperature

4 0.0500 0.0300 2.66 x 10-6

Constant

x 2 x 2

Between the last two experiments (3 & 4), the concentration of [B} doubles, the

concentration of [A] stays constant and Reaction Rate doubles. It follows that the initial

rate is directly proportional to the initial concentration of [B]. The reaction is therefore of

the first order with respect to [B] and n = 1.

It follows that for this reaction the Rate Law is:

Rate Law = k [A]1[B]1


Chemistry 102

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EXPERIMENT 1




4. The Rate Constant

The Rate Constant “k” is characteristic for every reaction and it is independent of the

concentrations of the reactants. However, “k” is temperature dependent.

The rate constant, “k” can be calculated, after the reaction orders have been determined

experimentally, from any of the reactions that have been run at the same temperature (for

example at room temperature, commonly abbreviated R.T.).

In this experiment data for four reactions run with different initial concentrations of

reactants but at the same temperature (room temperature) provide the experimental data

needed for the calculation of the Rate Constant, k, at R.T. In order to obtain an accurate

value for the rate constant, “k” will be calculated for every one of the four experiments

and the average value of the rate constant will be reported in the final expression of the

Rate Law.

An example of calculation of “k” at room temperature is given below for a reaction with

a pattern similar to the reaction studied in this experiment.

aA + bB cC + dD

To calculate “k”:

First solve the Rate Law for k: Rate

Rate = k [A]m [B]n Solve for k k =

[A]m [B]n

Next, substitute the Initial Concentrations of the Reactants and the experimentally determined Reaction Orders and Initial Rate

Assume:

[A] = Initial Concentration of Reactant A = 0.300 M

[B] = Initial Concentration of Reactant B = 0.450 M

m = Reaction Order with respect to Reactant A = 1

n = Reaction Order with respect to Reactant B = 1

Initial Rate = 4.83 x 10-6 M/s

4.83 x 10-6 M/s

k = = 3.58 x 10-3 M-1 . s-1

[0.0300 M]1 [0.0450 M]1


Chemistry 102

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EXPERIMENT 1




II. THE EFFECT OF TEMPERATURE ON REACTION RATE

Changing the temperature changes the Rate of Reaction, by the effect the temperature change

has on the Rate Constant, k. In fact, the Rate Constant, k is only constant when the

temperature remains constant.

The Arrhenius equation quantifies the temperature dependence of the rate constant in the

following form:

Ea

-

RT

k = A e

where:

k = The Rate Constant

Ea = The Activation Energy

is an energy barrier that must be surmounted for the reactants to be

transformed into products.

at a given temperature, the higher the Activation Energy, the slower the

reaction rate.

A = The Frequency Factor represents the number of times that the molecules of reactants approach the

activation barrier per unit time.

R = The universal gas constant = 8.314 J/K . mol

T = The absolute temperature e = 2.718

The factors included in the Arrhenius equation are important quantities in understanding the

kinetics of reactions. This experiment aims to:

1. To determine the numerical value of the Activation Energy, Ea for the reaction

being studied and,

2. To determine the ratio by which reaction rate increases when the temperature

increases by 100C (arbitrarily chosen from 200C to 300C)


Chemistry 102

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EXPERIMENT 1




1. Determination of Activation Energy, Ea

Since the temperature dependence of the reaction rate is contained in the rate constant,

we need to focus on the effect of temperature on the Rate Constant, “k”.

This is achieved by obtaining, analyzing and interpreting kinetic data obtained in the

laboratory. The experimental data is obtained by:

Running several experiments with the same initial concentrations of reactants but at different temperatures and determining the corresponding Reaction Rates

Calculating from experimental data the values of the Rate Constants, k, for these

experiments.

The experimental data obtained is analyzed and interpreted by using the Arrhenius

equation.

Recall that according to the Arrhenius Equation:

Ea Taking the natural log of both sides of this equation yields:

- Ea Ea 1

RT ln k = ln A - OR ln k = - + lnA

k = A e RT R T

The equation we have obtained has the form of the equation of a straight line: y = b + mx

Ea 1

ln k = - + lnA

R T

y = (m x) + b

Dependent Slope Independent

Variable Variable

Vertical Horizontal

Coordinate Coordinate

A plot of the natural log of the Rate Constant, k (ln k) versus the inverse of the Absolute

Temperature (1/T) yields a straight line with a slope = – Ea/R.


Chemistry 102

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EXPERIMENT 1




Such a plot is called an Arrhenius plot and is commonly used in the calculation of kinetic data,

such as the Activation Energy, Ea.

ln k

Ea

Slope = -

R

1

(K -1)

T Figure 1

The Activation Energy, Ea can be calculated from the slope (slope = - Ea/R) of the straight line.

2. Determination of the effect of a 100C temperature increase on Reaction Rate.

The steps involved in this calculations are are listed below:

Convert t1 = 200C and t2 = 300C respectively to 0K (T1 and T2)

Calculate 1/T1 and 1/T2 (K-1)

Read from the Arrhenius plot, the corresponding values for ln k (ln k1 and ln k2):

a

ln k2

ln k1

1 1 (K-1)

T2 T1

Figure 2

Take the anti ln of the above values to obtain: “k1” at 200C, and

“k2” at 300C

Calculate the two reaction rates by substituting the respective “k” values and the initial reaction concentrations (recall that the concentrations of reactants are the

same as in Reaction Run # 1)

Divide the reaction rate at 300C by the reaction rate at 200C


Chemistry 102

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EXPERIMENT 1




PROCEDURE: PART I: DETERMINATION OF THE RATE LAW

The table below summarizes the volumes of reactants and the reaction conditions under which

the experimental data will be obtained for all parts of the experiment.

TABLE I

ALL KINETIC RUNS

250 mL Reaction Flask 50 mL Flask

Run

Nr.

Temp KI

0.200 M

Reactant 1

Na2S2O3

0.00500 M

KCl

0.200 M

Soluble

Starch

(NH4)2S2O8

0.100 M

Reactant 2

(NH4)2SO4

0.100 M

CuSO4

0.100 M

1 Room

Temp

20.00mL 10.00 mL 0.00 mL 3-4

drops

20.00 mL 0.00 mL 0.00 mL

2 Room

Temp

10.00 mL 10.00 mL 10.00 mL 3-4

drops

20.00 mL 0.00 mL 0.00 mL

3 Room

Temp

20.00 mL 10.00 mL 0.00 mL 3-4

drops

10.00 mL 10.00 mL 0.00 mL

4 Room

Temp

20.00 mL 10.00 mL 0.00 mL 3-4

drops

5.00 mL 15.00 mL 0.00 mL

5 About

00C

20.00 mL 10.00 mL 0.00 mL 3-4

drops

20.00 mL 0.00 mL 0.00 mL

6 About

100C

20.00 mL 10.00 mL 0.00 mL 3-4

drops

20.00 mL 0.00 mL 0.00 mL

7 About

40oC

20.00 mL 10.00 mL 0.00 mL 3-4

drops

20.00 mL 0.00 mL 0.00 mL

The table below summarizes the volumes of reactants to be used in making up the four

reaction mixture for the reactions run at room temperature

TABLE II

Kinetics Runs at Room Temperature and Varying Concentrations

250 mL Reaction Flask 50 mL Flask

Run

Nr.

KI

0.200 M

Reactant 1

Na2S2O3

0.00500

M

KCl

0.200 M

Soluble

Starch

(NH4)2S2O8

0.100 M

Reactant 2

(NH4)2SO4

0.100 M

CuSO4

0.100 M

1 20.00 mL 10.00 mL 0.00 mL 3-4 drops 20.00 mL 0.00 mL 0.00 mL




The actual procedure for carrying out each reaction run the same for all runs and it is described

in detail for Reaction Run Nr. 1


Chemistry 102

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EXPERIMENT 1




a. Filling the 250 mL Erlenmeyer Flask (Reaction Flask)

Use a buret to add exactly 20.00 mL of 0.200 M KI into the 250 – mL Erlenmeyer Flask, which we shall refer to as the Reaction Flask.

Use a buret to add exactly 10.00 mL of 0.00500 M Na2S2O3 into this flask.

Add 3 or 4 drops of starch solution to the flask.

Mix the contents of the flask by swirling the flask on the bench top.

b. Filling the 50 – mL Erlenmeyer Flask

Use a buret to add exactly 20.00 mL of 0.100 M (NH4)2S2O8 into this flask.

c. Temperature readings

Read the temperature of the solution in the 250 – mL flask with a rinsed, dried

thermometer. Record the temperature to the nearest degree.

Remove the thermometer, rinse and dry it, and measure the temperature of the solution in the 50 – mL flask. The solutions in both flasks should be at the same temperature ± 20C,

since both solutions have been kept at room temperature. Record the temperature as

Room Temperature (RT) to the nearest degree. If there is a slight difference between the

temperature readings in the two flasks (due to the limited accuracy of the thermometers),

record the average temperature as the initial temperature (Room Temperature).

d. Mixing and Timing

Figure 1 Figure 2 Figure 3

Have a timer available. While you start swirling the Continue swirling the Pour the solution from the solutions, your partner should solution until the blue

50 – mL flask into the 250 mL start the timer. color appears.

reaction flask. You should leave the 50 – mL Note and record the time

flask over the mouth of the at which the blue color

reaction flask, as shown above. appears, to the nearest

0.1 of a second Photos by Andrew Huertas

Figures 3a, 3b & 3c

Check the temperature of the mixture. Note and record the temperature as the final

temperature. The temperature of the reaction mixture should be recorded as the average

of the initial and final temperatures


Chemistry 102

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EXPERIMENT 1




If the timing was faulty, repeat the entire run.

e. Repeat the experiment with the other three mixtures in the table

Both flasks should be rinsed with tap water and deionized water and drained between

experiments.

Burets should be used in measuring the volumes of the following five solutions: KI, Na2S2O3, KCl, (NH4)2S2O8 and (NH4)2SO4

NOTE:

Although KCl and (NH4)2SO4 are not reactants in the reaction we are studying, these

reagents serve to maintain the effective concentrations of all ions (“ionic strength”) at a

constant level. Holding the ionic strength constant removes the dependence of the reaction

rate on variations in the solvent.

PART II. THE EFFECT OF TEMPERATURE ON REACTION RATE

To obtain experimental data reflecting the dependence of the Reaction Rate on Temperature, the

reaction will be carried out at several different temperatures, but with the same concentration of

all reactants as in ReactionRun # 1.. Table III below indicates the reaction conditions:

TABLE III

Reaction Mixtures at different temperatures

Reaction

Run

Temperature

Range

Reaction Mixture

5 About 00C Same as in Rxn Run # 1



Photo by Andrew Huertas Reaction Run # 5 (about 00C) Figure 4

The reaction is carried out by adding the same volumes of the same solutions as in

Reaction Run # 1.

Prepare mixtures of ice-cold water mixtures in a 250 mL beaker (for the 50 mL flask) and a 600 mL beaker (for the 250 mL flask), which will be used to cool the contents

of the two flasks, when immersed in the ice/cold water slurry.

Place one thermometer in each of the flasks and swirl the contents of the flasks.

Immerse the flasks in the beakers containing the ice/cold water slurry. While swirling

the two flasks gently, follow the temperature readings on the two thermometers and

allow the temperature of the solutions to stabilize. Make sure that the bulbs of both

thermometers are immersed in the respective solutions. This is particularly important

for the temperature reading in the 50 mL flask and it may require to tilt this flask,

while immersed in the ice/cold water slurry. Your aim is to reach a temperature close

to 0 0C (± 20 C) for both solutions and that this temperature be the same (or almost the

same) for the two solutions contained in the two flasks. It is not imperative that the

temperature be exactly 00C (difficult to achieve) but it is important that the

temperature of the two solutions be as close as possible to each other.


Chemistry 102

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EXPERIMENT 1




HINTS

The ice/cold water slurry should be mixed with a glass rod or a spatula.

The flasks should be fully immersed in the ice/cold water slurry.

The contents of the flasks should be gently mixed by swirling.

If the temperature reached is not low enough, the temperature of the solutions

may be further lowered by adding and mixing a spoonful of rock salt into the

ice/cold water slurry.

When you are confident that the temperatures of the two solutions are about the same and within ± 20 C from 00 C

Note and record this temperature (Initial temperature), and

Proceed to pour the solution from the 50 – mL flask into the 250 mL reaction

flask, while keeping the 250 mL flask immersed in the ice/water slurry and the

thermometer immersed in the flask.

At this point, it is best to divide the tasks of the two team members as follows:

One team member is responsible for: o Starting the stopwatch,

o Observing and recording the time at which the color change occurs

to the nearest 0.1 of a second.

The other team member is responsible for:

o Keeping the 250 mL flask immersed in the ice/cold water slurry for the

entire reaction interval.

o Mixing the contents of the 250 mL reaction flask by swirling the flask,

while keeping the thermometer immersed in the reaction mixture.

o Keeping track of the temperature of the reaction mixture, and

o Recording the temperature at the exact time at which the color change

occurred (final temperature) of the reaction mixture)

o The temperature of the reaction mixture should be recorded as the

average of the initial and final temperatures for this reaction run.

Reaction Run # 6 (about 100C)

Repeat the experiment at about 100C.

Note and record the initial temperature before mixing, the time required for the reaction and the final temperature of the reaction mixture at the time the solution

turned blue.

Record the temperature of the reaction mixture for this run as the average of the initial and the final temperature.

Reaction Run # 1 (Room Temperature)

Data for this Reaction Rate is already available from Part I.

Reaction Run # 7 (about 400C)

Repeat the experiment at about 400C.

Note and record the initial temperature before mixing, the time required for the reaction and the final temperature of the reaction mixture at the time the solution turned blue.

Record the temperature of the reaction mixture for this run as the average of the initial and the final temperature.


Chemistry 102

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EXPERIMENT 1




REPORT FORM

NAME: ______________________ Date: _____________ Partner: ______________________

PART I: DETERMINATION OF THE RATE LAW

1. Initial Concentrations of Reactants

For each kinetic run, calculate the initial concentration of the reactants:

[I-] and [S2O82-].

These calculations should be done prior to performing the

experiment! Since the reaction takes place in a total volume of 50.00 mL, this volume must be taken

into account in calculating the initial concentration of the two reactants.

For example, in Run 1, since the 20.00 mL of 0.200 M KI added reacts in a total volume

of 50.00 mL, the initial concentration of [I-]0 can be calculated as follows:

20.00 mL

[I-]0 = 0.200 M KI = 0.0800 M KI

50.00 mL

Similarly, in Run 1, the initial concentration of [S2O82-]0 is calculated as follows:

20.00 mL

[S2O82-]0 = 0.100 M (NH4)2S2O8 = 0.0400 M (NH4)2S2O8

50.00 mL

On the next page, carry out similar calculations for all other initial values of the two

reactants and complete the appropriate columns in DATA TABLE III.

You are required to:

Show all your calculations

Include units in your calculations, and

Express all measured quantities (including your answer) in the appropriate number of significant figures.


Chemistry 102

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EXPERIMENT 1




DATA TABLE I

[I-]0 [S2O82-]0

Run 1

Run 2

Same as in Run 1

Run 3

Same as in Run 1

Run 4

Same as in Run 1

2. Reaction Rates

In order to determine the Reaction Rates, the following quantities must be known:

Initial concentration of [S2O32-] added and completely used up (see below)

Concentration of [I2] produced, at the time the deep blue color (tcolor) appears.

(see below); Enter this value in the Data Table I on the next page.

The time in seconds when the deep blue color (tcolor) appears for each Reaction Run

(determined experimentally). Enter these data in DATA TABLE III.

DATA TABLE II

[S2O32-]added

[S2O32-]

[I2]produced =

2

All Runs


Chemistry 102

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EXPERIMENT 1




DATA TABLE III

Initial Concentrations and corresponding Reaction Rates

RATE*: Express Rate as (A x 10-6) throughout the entire experiment

Run

Nr.

Temp.

(0C)

(Room

Temp)

[I-]0

(M)

[S2O82-]0

(M)

TIME

(tcolor)

(s)

RATE

Expressed as A x 10-6

[I2]produced

tcolor

(M . s-1)

1

2

3

4


Chemistry 102

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EXPERIMENT 1




3. Reaction Orders with respect to Reactants

The general formula for the Rate Law for the reaction studied is:

Rate1 = k [I-]m [S2O82-]n

Rate2 = k [I-]m [S2O82-]n

When performing your calculations you are required to:

Show all calculations neatly, in a well-organized manner and in detail (follow the format of the example shown on page 4 or the sample calculation presented in

your textbook (page 605)

Round off your answer to an integer.

a. Reaction Order with respect to [I-}

From Run 1 and Run 2

Value of “m” rounded off to an integer =

b. Reaction Order with respect to [S2O82-]

From Run 3 and Run 4

Value of “n” rounded off to an integer =


Chemistry 102

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EXPERIMENT 1




k1 =

4. Rate Constant “k” at Room Temperature

The Rate Constant, “k” can be calculated by substituting the known values of reactant

concentrations, the reaction orders and the corresponding reaction rates in the formula

of the Rate Law. Refer to the example on page 5.

Write the equation for the Rate Law: Rate =

DATA TABLE IV

Calculate “k” for Reactions 1 through 6. Show your calculations and include units!

Run 1

k1 =

Run 2

k2 =

Run 3

k3 =

Run 4

k4 =

DATA TABLE V

Units Summary of Rate Constants “k” for Reaction Runs at Room Temperature

Run Nr. 1 2 3 4

k

( )

k (Average)

( )

Write the complete form of the Rate Law:

Include the formulas of both Reactants. Do not include numerical values

Include the respective Reaction Orders for both Reactants

Include the experimentally determined numerical value of “k”, expressed in the correct units.

RATE =


Chemistry 102

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EXPERIMENT 1




PART II: THE EFFECT OF TEMPERATURE ON REACTION RATE

To determine the quantitative relationship between Reaction Rate and Temperature you will

be plotting an Arrhenius Plot (ln k versus 1/T) to determine the:

1. Activation Energy (Ea) and

2. The ratio by which the reaction rate increases when the temperature is increased by 100C

(arbitrarily chosen from 200C to 300C)

1. Determination of the Activation Energy, Ea

Summarize your experimental data for the Reactions Runs with the same

concentration of all reactants (as used in Reaction Run # 1), but run at different

temperatures. In this manner, the corresponding reaction rates for these runs will be

affected by temperature only.

DATA TABLE VI

Run

Nr.

Temp.

Range

Initial

Temperature

of

Reactants

(0C)

Final

Temperature

of

Reaction

Mixture

(0C)

Average

Temperature

of

Reaction

Mixture

(0C)

Time

(tcolor)

(s)

RATE

[I2]produced

tcolor

(M . s-1)

5 About

00C

6 About

100C

1 Room

Temp.

7 About

400C

Calculate the Rate Constants for Runs 5, 6 & 7 from the initial concentrations of

reactants and the corresponding reaction rates.

DATA TABLE VII

(Calculation do not need to be shown)

Run

Nr.

Recorded

Average

Temp.

(K)

[I-]0

(M)

[S2O82-]0

(M)

RATE

(M . s-1)

Rate Constant

k

(M-1 . s-1)

(Expressed as A x 10-3)

5 About

00C

6 About

100C

1 Room

Temp

Calculated Average

7 About

400C


Chemistry 102

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EXPERIMENT 1




Calculate and collect the data that will be used for plotting ln k versus 1/T (The

Arrhenius Plot)

DATA TABLE VIII

(Calculations do not need to be shown)

Run

Nr.

1

T

(K-1)

ln k

5 About

00C

6 About

100C

1 Room

Temp

7 About

400C

Plot a graph of ln k vs. 1/T.

A sample graph is attached to your Report Form. Please follow the format, scale, data

reporting style and all the other details included in the sample graph.

Calculate the slope of the graph and show all the calculations on the graph.

Calculate the Activation Energy (Ea) from the slope of the graph

Please show calculations and include units.

J

Slope: __________ Recall: R = 8.314

mol . K

Ea = J/mol Ea = KJ/mol


Chemistry 102

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EXPERIMENT 1




2. Determination of the effect of 100C temperature increase on Reaction Rate

Collect your data:

t1 = 200C T1 = _____ K 1/T1 = _______ (K-1)

t2 = 300C T2 = _____ K 1/T2 = _______ (K-1)

Read the corresponding values for ln k1 and ln k2 from your Arrhenius plot.

Indicate these values and the source of these readings on your graph

(as shown on Figure 2, page 8)

Follow the guidelines given below Figure 2 (page 8) to calculate the ratio between the

Reaction Rate at 300C and the Reaction Rate at 200C

Show all calculations neatly, in a well-organized manner and in detail

Include units in your calculations

Round off your answer to the nearest integer.

State your conclusion:

______________________________________________________________________________

______________________________________________________________________________

Bibliography:

1. Nivaldo J. Tro, “Chemistry: A Molecular Approach”, Third Edition

2. R.A.D. Wentworth “Experiments in General Chemistry”, Sixth Edition

3. James M. Postma & all, “Chemistry in the Laboratory”, Seventh Edition


Documents

Experiment 1 The Iodine “Clock” Reaction - CHEMISTRY …proffenyes.weebly.com/uploads/2/5/2/3/25237319/experiment_1_-_the... · Rate = k [A]m [B]n where: k: is a ... log 2.00