A Kinetics Lab Experiment | Lab Reports Chemistry

A Kinetics Experiment

The Rate of a Chemical Reaction: A Clock Reaction

Andrea Deoudes

February 2, 2010

Introduction:

The rates of chemical reactions and the ability to control those rates are crucial aspects of life.

Chemical kinetics is the study of the rates at which chemical reactions occur, the factors that

affect the speed of reactions, and the mechanisms by which reactions proceed. The reaction rate

depends on the reactants, the concentrations of the reactants, the temperature at which the

reaction takes place, and any catalysts or inhibitors that affect the reaction. If a chemical

reaction has a fast rate, a large portion of the molecules react to form products in a given time

period. If a chemical reaction has a slow rate, a small portion of molecules react to form

products in a given time period.

This experiment studied the kinetics of a reaction between an iodide ion (I-1) and a

peroxydisulfate ion (S2O8-2) in the first reaction: 2I-1 + S2O8-2  I2 + 2SO4-2. This is a relatively

slow reaction. The reaction rate is dependent on the concentrations of the reactants, following

the rate law: Rate = k[I-1]m[S2O8-2]n.

In order to study the kinetics of this reaction, or any reaction, there must be an experimental way

to measure the concentration of at least one of the reactants or products as a function of time.

This was done in this experiment using a second reaction, 2S2O3-2 + I2  S4O6-2 + 2I-1, which

occurred simultaneously with the reaction under investigation. Adding starch to the mixture

allowed the S2O3-2 of the second reaction to act as a built in “clock;” the mixture turned blue

when all of the S2O3-2 had been consumed. Thus, the concentration of the S2O3-2 could be

measured over a period of time, using the initial concentration of S2O3-2 in the mixture,

calculated to be 0.00126M, and the final concentration, 0M. These values could be used to

calculate the rate according to the formula: Rate = -½ Δ[S2O3-2] / Δt, where Δt was the time

required for the color change to occur.

After solving for the reaction rate in the first 3 experiments, one could obtain the values for m

and n in the rate law equation (Rate = k[I-1]m[S2O8-2]n). The values of m and n were determined

by observing the change in the reaction rate that occurred as the result of a change in the

concentration of I-1 and S2O8-2, respectively.

In this experiment, the reactions were done at four different temperatures in order to observe the

effect of temperature on the rate of reaction. In studying the kinetics of a reaction, it is also

important to consider the effects that potential catalysts or inhibitors can have on the reaction

rate. These effects were observed by adding one drop of Ag+1 or Cu+2 solution to some of the

data runs and comparing the results to data runs with no additives.

The frequency factor and activation energy are two additional quantities that are important in

understanding the kinetics of any reaction. In order to measure these factors, the Arrhenius

equation, k = Ae-Ea/RT, can be used. After taking the natural log of both sides of the equation, the

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A Kinetics Experiment

The Rate of a Chemical Reaction: A Clock Reaction

Andrea Deoudes

February 2, 2010

Introduction:

The rates of chemical reactions and the ability to control those rates are crucial aspects of life.

Chemical kinetics is the study of the rates at which chemical reactions occur, the factors that

affect the speed of reactions, and the mechanisms by which reactions proceed. The reaction rate

depends on the reactants, the concentrations of the reactants, the temperature at which the

reaction takes place, and any catalysts or inhibitors that affect the reaction. If a chemical

reaction has a fast rate, a large portion of the molecules react to form products in a given time

period. If a chemical reaction has a slow rate, a small portion of molecules react to form

products in a given time period.

This experiment studied the kinetics of a reaction between an iodide ion (I

) and a

peroxydisulfate ion (S 2 O 8 -2) in the first reaction: 2I-1^ + S 2 O 8 -2^  I 2 + 2SO 4 -2. This is a relatively

slow reaction. The reaction rate is dependent on the concentrations of the reactants, following

the rate law: Rate = k[I-1]m[S 2 O 8 -2]n.

In order to study the kinetics of this reaction, or any reaction, there must be an experimental way

to measure the concentration of at least one of the reactants or products as a function of time.

This was done in this experiment using a second reaction, 2S 2 O 3 -2^ + I 2  S 4 O 6 -2^ + 2I-1, which

occurred simultaneously with the reaction under investigation. Adding starch to the mixture

allowed the S 2 O 3

of the second reaction to act as a built in “clock;” the mixture turned blue

when all of the S 2 O 3

had been consumed. Thus, the concentration of the S 2 O 3

could be

measured over a period of time, using the initial concentration of S 2 O 3 -2^ in the mixture,

calculated to be 0.00126M, and the final concentration, 0M. These values could be used to

calculate the rate according to the formula: Rate = -½ Δ[S 2 O 3 -2] / Δt, where Δt was the time

required for the color change to occur.

After solving for the reaction rate in the first 3 experiments, one could obtain the values for m

and n in the rate law equation (Rate = k[I-1]m[S 2 O 8 -2]n). The values of m and n were determined

by observing the change in the reaction rate that occurred as the result of a change in the

concentration of I

and S 2 O 8

, respectively.

In this experiment, the reactions were done at four different temperatures in order to observe the

effect of temperature on the rate of reaction. In studying the kinetics of a reaction, it is also

important to consider the effects that potential catalysts or inhibitors can have on the reaction

rate. These effects were observed by adding one drop of Ag

or Cu

solution to some of the

data runs and comparing the results to data runs with no additives.

The frequency factor and activation energy are two additional quantities that are important in

understanding the kinetics of any reaction. In order to measure these factors, the Arrhenius

equation, k = Ae

-Ea/RT

, can be used. After taking the natural log of both sides of the equation, the

resulting new equation is lnk = -Ea/RT + lnA. This is a linear equation in the form of y = mx + b,

where y =lnk, m = -Ea/R, x = 1/T, and b = lnA. The data from this experiment could be used to

create a plot of lnk vs. 1/T, and the linear trendline of the plot could be used to determine

activation energy and frequency factor.

Experimental Procedure:

The class was split into groups in order to determine which group was responsible for each of the

various data runs as listed in Table 1, which was taken from the laboratory manual.

Table 1.

All of the experiments were done using two clean and dry 50mL Erlenmeyer flasks, one labeled

“A” and the other labeled “B”. In order to ensure accuracy, each experiment was conducted

twice and the average reaction time from the two runs was used. Prior to each experiment, the

groups used a thermometer to determine the temperature at which the reaction would take place.

For the experiments performed at approximately room temperature, students completed the

reaction at the lab benches. For the experiments done at around 10˚C, 30˚C, and 40˚C, water

baths were used to achieve the desired temperature for the reaction.

The groups prepared the solutions as indicated by Table 1. The solution in Flask “A” was made

with varying volumes of 0.20% starch, 1.2x10-2M Na 2 S 2 O 3 , 0.2M KI, and 0.2M KNO 3. The

solution in Flask “B” was made with varying volumes of 0.2M (NH 4 ) 2 S 2 O 8 and 0.2M (NH 4 ) 2 SO 4.

After the solutions for the room temperature runs were prepared in the two flasks, the solutions

were mixed to create the reaction. One person acted as the mixer and one person acted as the

timer. The timer began timing as the mixer poured the solution from one flask into the other,

then back and forth a few times. Then, the mixer gently swirled the contents in the flask that

contained both solutions. The flask was swirled continuously until the color change was

observed. As soon as the color change was observed, the timer stopped the stopwatch and

13 0.0421 0.0211 10.5 313 6.00E- 05 6.75E- 02

14 0.0421 0.0211 51.5 313 1.22E- 05 1.37E- 02

Avg of 1,2,3 294 4.58E- 03

The natural log of the Arrhenius equation, lnk = -Ea/RT + lnA, can be applied to the Solution

Temperature and Rate Constant data from Table 2 to create Arrhenius plots of lnk vs. 1/T. Table

3 contains the data that corresponds with Figure A, the Arrhenius plot of reactions without

additives. Table 4 contains the data that corresponds with Figure B, the Arrhenius plot of

reactions for which Cu+2^ was added. Table 5 contains the data that corresponds with Figure C,

the Arrhenius plot of reactions for which Ag+1^ was added.

Table 3.

Experiment #

(No Additives) 1/T (K) lnk

Avg of 1,2,3 3.40E- 03 - 5. 4 3.53E- 03 - 6. 9 3.30E- 03 - 5. 12 3.19E- 03 - 4.

Figure A.

Table 4.

Experiment

(Cu+2added) 1/T (in K) lnk

5 3.53E- 03 - 4.

7 3.40E- 03 - 3.

10 3.30E- 03 - 3.

13 3.19E- 03 - 2.

y = - 6218.4x + 15.

8
6
4
2 0 3.10E- 03 3.20E- 03 3.30E- 03 3.40E- 03 3.50E- 03 3.60E- 03 ln ƙ 1/T

Arrhenius Plot of Reactions

Without Additives

Figure B.

Table 5.

Experiment (Ag+1^ Added) 1/T (in K) lnk 6 3.53E- 03 - 6. 8 3.40E- 03 - 5. 11 3.30E- 03 - 4. 14 3.19E- 03 - 4.

Figure C.

Experiments 2, 4, 9, and 12, were very similar runs, only differing in the temperature at which

the reaction was conducted. Thus, in order to observe the effects of temperature on the rate of

reaction, a plot of reaction rate vs. temperature was constructed for those four experiments

(Figure D).

y = - 3686.6x + 8.

5
4
3
2
1 0 3.10E- 03 3.20E- 03 3.30E- 03 3.40E- 03 3.50E- 03 3.60E- 03 ln ƙ 1/T

Arrhenius Plot of Reactions

With Cu

y = - 6019.6x + 14.

8
6
4
2 0 3.10E- 03 3.20E- 03 3.30E- 03 3.40E- 03 3.50E- 03 3.60E- 03 ln ƙ 1/T

A Kinetics Lab Experiment, Lab Reports of Chemistry

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Partial preview of the text

Download A Kinetics Lab Experiment and more Lab Reports Chemistry in PDF only on Docsity!

A Kinetics Experiment

The Rate of a Chemical Reaction: A Clock Reaction

Andrea Deoudes

February 2, 2010

Introduction:

The rates of chemical reactions and the ability to control those rates are crucial aspects of life.

Chemical kinetics is the study of the rates at which chemical reactions occur, the factors that

affect the speed of reactions, and the mechanisms by which reactions proceed. The reaction rate

depends on the reactants, the concentrations of the reactants, the temperature at which the

reaction takes place, and any catalysts or inhibitors that affect the reaction. If a chemical

reaction has a fast rate, a large portion of the molecules react to form products in a given time

period. If a chemical reaction has a slow rate, a small portion of molecules react to form

products in a given time period.

This experiment studied the kinetics of a reaction between an iodide ion (I

) and a

peroxydisulfate ion (S 2 O 8 -2) in the first reaction: 2I-1^ + S 2 O 8 -2^  I 2 + 2SO 4 -2. This is a relatively

slow reaction. The reaction rate is dependent on the concentrations of the reactants, following

the rate law: Rate = k[I-1]m[S 2 O 8 -2]n.

In order to study the kinetics of this reaction, or any reaction, there must be an experimental way

to measure the concentration of at least one of the reactants or products as a function of time.

This was done in this experiment using a second reaction, 2S 2 O 3 -2^ + I 2  S 4 O 6 -2^ + 2I-1, which

occurred simultaneously with the reaction under investigation. Adding starch to the mixture

allowed the S 2 O 3

of the second reaction to act as a built in “clock;” the mixture turned blue

when all of the S 2 O 3

had been consumed. Thus, the concentration of the S 2 O 3

could be

measured over a period of time, using the initial concentration of S 2 O 3 -2^ in the mixture,

calculated to be 0.00126M, and the final concentration, 0M. These values could be used to

calculate the rate according to the formula: Rate = -½ Δ[S 2 O 3 -2] / Δt, where Δt was the time

required for the color change to occur.

After solving for the reaction rate in the first 3 experiments, one could obtain the values for m

and n in the rate law equation (Rate = k[I-1]m[S 2 O 8 -2]n). The values of m and n were determined

by observing the change in the reaction rate that occurred as the result of a change in the

concentration of I

and S 2 O 8

, respectively.

In this experiment, the reactions were done at four different temperatures in order to observe the

effect of temperature on the rate of reaction. In studying the kinetics of a reaction, it is also

important to consider the effects that potential catalysts or inhibitors can have on the reaction

rate. These effects were observed by adding one drop of Ag

or Cu

solution to some of the

data runs and comparing the results to data runs with no additives.

The frequency factor and activation energy are two additional quantities that are important in

understanding the kinetics of any reaction. In order to measure these factors, the Arrhenius

equation, k = Ae

, can be used. After taking the natural log of both sides of the equation, the

resulting new equation is lnk = -Ea/RT + lnA. This is a linear equation in the form of y = mx + b,

where y =lnk, m = -Ea/R, x = 1/T, and b = lnA. The data from this experiment could be used to

create a plot of lnk vs. 1/T, and the linear trendline of the plot could be used to determine

activation energy and frequency factor.

Experimental Procedure:

The class was split into groups in order to determine which group was responsible for each of the

various data runs as listed in Table 1, which was taken from the laboratory manual.

Table 1.

All of the experiments were done using two clean and dry 50mL Erlenmeyer flasks, one labeled

“A” and the other labeled “B”. In order to ensure accuracy, each experiment was conducted

twice and the average reaction time from the two runs was used. Prior to each experiment, the

groups used a thermometer to determine the temperature at which the reaction would take place.

For the experiments performed at approximately room temperature, students completed the

reaction at the lab benches. For the experiments done at around 10˚C, 30˚C, and 40˚C, water

baths were used to achieve the desired temperature for the reaction.

The groups prepared the solutions as indicated by Table 1. The solution in Flask “A” was made

with varying volumes of 0.20% starch, 1.2x10-2M Na 2 S 2 O 3 , 0.2M KI, and 0.2M KNO 3. The

solution in Flask “B” was made with varying volumes of 0.2M (NH 4 ) 2 S 2 O 8 and 0.2M (NH 4 ) 2 SO 4.

After the solutions for the room temperature runs were prepared in the two flasks, the solutions

were mixed to create the reaction. One person acted as the mixer and one person acted as the

timer. The timer began timing as the mixer poured the solution from one flask into the other,

then back and forth a few times. Then, the mixer gently swirled the contents in the flask that

contained both solutions. The flask was swirled continuously until the color change was

observed. As soon as the color change was observed, the timer stopped the stopwatch and

13 0.0421 0.0211 10.5 313 6.00E- 05 6.75E- 02

14 0.0421 0.0211 51.5 313 1.22E- 05 1.37E- 02

The natural log of the Arrhenius equation, lnk = -Ea/RT + lnA, can be applied to the Solution

Temperature and Rate Constant data from Table 2 to create Arrhenius plots of lnk vs. 1/T. Table