Iodine Clock Coursework

The Iodine Clock Investigation

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The Iodine Clock Investigation


Introduction

This is an investigation into the rate of a reaction and the factors
that contribute to how fast a reaction will take place. Through the
recording and analysis of raw data, this investigation also allows us
to apply generally accepted scientific rules and to test them against
results gained from accurate experimental procedures.

Aim

The aim of this experiment is to investigate the rate at which iodine
is formed when the concentration and temperature of the reactants are
varied, and to attempt to find the order and activation energy.

The Chemistry

'THE IODINE CLOCK' - This is the experiment that will be used to
investigate reaction rates, and it is a reaction between acidified
hydrogen peroxide and potassium iodide:

2H+(aq) + 2I¯ (aq) + H[-1] 2O2 (l) ÕI2 (aq) + 2H2O2 (aq)

Iodide ions are firstly oxidised by the hydrogen peroxide, as shown in
the above equation. The iodine that is then produced reacts
immediately reacts with thiosulphate ions as follows:

I2 (aq) + 2Na2S2O3 (aq) Õ 2NaI (aq) + Na2S406 (aq)

As soon as all of the thiosulphate ions have reacted with the iodine,
the excess iodine molecules react with the 2% starch solution that is
present in the reaction. This can be seen as an instant change in
colour, from a colourless solution, to a deep purple coloured
solution. This change in colour denotes the completion of the
reaction.

Factors affecting the rate of reactions:

All chemical reactions occur at a definite rate under particular
conditions. In order to increase the rate at which reactions occur,
the frequency at which reacting molecules collide must be increased.
This may be achieved in a number of ways:

1. By increasing the concentrations of reacting species.

2. By increasing the temperature.

3. By increasing the pressure (only really significant in reactions
involving gases).

4. By the use of a suitable catalyst.

5. In the case of solids, by reducing particle size and thus

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Related Searches

Clock         Starch Solution         Iodide Ions         Activation Energy         Chemical Reactions         Reaction Rates         Molecules         Collide        




increasing the effective surface area.

6. In particular cases the rate may depend on the influence of
electromagnetic radiation such as visible or ultra-violet light.

In this investigation the two factors affecting the rate of reactions
that will be investigated, are concentration and temperature.

The Effect Of Concentration On The Rate Of A Reaction

In a reaction that takes place between two substances, A and B, if we
were to

look at how quickly substance A is used, the rate of the reaction
would be the rate of change of substance A, symbol rA. The rate of
change of concentration of substance B, rB may also have been
measured.

The rate of change of the concentration follows a general mathematical
expression in the form:

Rate = k [A]a[B]b[C]c

This is known as the concentration rate equation where:

· The rate equation has units of mol dm-3s-1 (other units may be
used).

· The square brackets denote the concentration (in mol dm-3).

· The sum of all of the indices is called the overall order of the
reaction.

'k' is a constant of proportionality called the rate constant. The
units of k depend on the order of the reaction and can be worked out
from the rest of the rate equation.

Order of the Reaction

The kinetics of a reaction can be classified in terms of its order;
these are experimentally determined quantities related to the rate
expression.

Rate = k [A]a[B]b[C]c

The order with regard to a particular species (A, B or C) is equal to
the power to which the concentration of this species is raised to (a,
b or c)

The overall order of the reaction is equal to the sum of the powers of
the concentration.

Order = a + b + c

Where the rate expression shows the reaction to be dependent on the
concentration of one reactant, the concentration of the substance is
raised to the power zero. Therefore, the rate is independent of the
concentration.

i.e. Rate = k[A]0

It is said to have zero order kinetics. When the results of a zero
order reaction are plotted on graphs of concentration vs. time, and
concentration vs. rate, the following results are typical to zero
order reactions:

[IMAGE]

[IMAGE]




When the rate expression shows the reaction to be dependant on the
concentration of one reactant raised to the power one, it is said to
have first order kinetics.

The following graphs show results typical to first order reactions:

[IMAGE]

[IMAGE]




The first zero order graph showing concentration vs. time is such that
the time it takes for the concentration of the reactant to be halved
is constant; this is known as half-life.

When the rate expression shows the reaction to be dependant on the
concentration of one reactant raised to the power two, it is said to
have second order kinetics.

[IMAGE] The following graphs show results typical to second order
reactions:

[IMAGE]

[IMAGE]




The final graph showing concentration vs. time will have a half-life
that isn't constant, but one that increases dramatically as the
reaction proceeds. This graph is visibly much 'deeper' than a graph
showing a first order reaction.

N.B.

The graphs above only show the order of a reaction with respect to
substance A, and do not necessarily show the overall order of the
reaction.

The Effect of Temperature on the Rate of Reaction

As the temperature of a reaction rises, the average speeds of reacting
particles also increase. At high temperatures there are more
collisions per second and this results in an increase in the rate of
the reaction.

For many reactions a 10 K (°C) increase in the temperature will
approximately double the rate of the reaction. A graph showing how the
rate of a reaction increase when compared to temperature will have the
following general shape:

[IMAGE]




In many reactions between gases it is the actual collision of
particles that control the rate of the reaction and for a collision
which actually results in a reaction the kinetic energy possessed by
the colliding particles must have a minimum energy, E.

Activation Energy[IMAGE][IMAGE]Products[IMAGE][IMAGE][IMAGE][IMAGE]
All reactions have what is known as activation energy, EA which is the
minimum energy possessed by reacting particles that is required to
initiate a chemical reaction.

[IMAGE]

Energy

Reactants




The activation energy of a reaction may be derived from the 'Arrhenius
Equation'.

The Arrhenius Equation

It can be shown that for one mole of colliding particles at a
temperature t, the

rate of reaction can be found from the following equation:

ln k = C - EA (1/T)

R

Where:

k = rate constant

R = gas constant (8×314 JK-1mol-1)

EA = Activation Energy

C = ln collision rate

The Arrhenius equation shows that when the temperature rises, there is
a large increase in the value of the rate constant, k. This
corresponds to a large increase in the number of collisions occurring
with the necessary minimum energy or activation energy, EA.

The following graph shows very simply that when you look at the
distribution of energies amongst gas molecules at different
temperatures T, there is a much higher proportion of molecules with
the necessary activation energy, or minimum energy for reaction, Emin,
at the higher temperature, T2, than at the lower temperature, T1.

[IMAGE]

Number of molecules having a given energy




Energy E



The activation energy of a reaction can be calculated from
experimental values using the Arrhenius equation. When the ln (rate)
of a reaction is plotted against the rate of a reaction, the following
shaped graph is typical:

[IMAGE]


When the gradient of this graph is measured, the EA can be calculated
as:

Gradient = - EA

R

Key Variables

Before we can even begin to consider carrying out any practical
experiments, the factors of the reaction which when changed might
affect the rate of the reaction, must be identified. These factors are
known as the key variables, and deciding which to vary, and which to
keep constant during the experiment becomes very important when it
comes to calculating such things as order of reactions and activation
energies.

The following, are all the factors that should be considered when
investigating the rate of this reaction:

· The concentration of KI

· The concentration of H2O2

· The concentration of H2SO4

· The concentration of Na2S2O3

· The temperature of the reaction mixture

Trial Experiments

Problem 1

Before it can be decided which of these factors to be controlled in
the various experiments, I must acquaint myself with the reaction
itself as there are certain points that must be gained:

1. To determine exactly which reactants affect the rate of the
reaction.

2. To choose suitable concentrations of reactants so that the
reactions proceed neither too quickly, nor too slow.

In order to gain the required information, trial experiments must be
conducted.

Experiment 1

Aim:

In this initial trial experiment, the aim is firstly, to become
acquainted with the method and procedure of the experiment; and
secondly, to confirm the fact that hydrogen peroxide affects the rate
of the reaction.

Apparatus/Equipment:

In these trial experiments, the following apparatus will be used:

· H2O2 - 5 Vol.

· KI -1.0M

· H2SO4 - 1.0M

· Na2S2O3 -0.005M

· 2% starch solution (fresh)

· 50ml burettes (x4)

· conical flasks

· volumetric flasks

· stands and clamps

· funnels

· stopclock

Method

Two conical flasks were taken, in the first was placed 10ml each of
KI, H2SO4, and Na2S2O3. In the second conical flask was placed 20ml of
H2O2, and two squirts of 2% starch solution from a bulbous pipette.
These volumes were measured out using the burettes, with each
substance in a separate burette.

The flask containing the peroxide was then poured directly into the
second flask, and the stopclock was started simultaneously. The
stopclock was then stopped when the reaction was completed, the end
point could be observed with a sudden change in colour from an initial
colourless solution, to a deep purple solution. The time taken for the
reaction to finish was then recorded.

This method was then repeated in exactly the same way, except that the
concentration of H2O2 used was diluted down to 2.5 Vol. The results of
this second experiment were also recorded.

Results

Concentration of H2O2

(Vol.)

Time taken for reaction to reach

End point. (secs)

5.0

4.02

2.5

7.35

Analysis of results

In the carrying out of this initial experiment, two points can be
deduced:

1. the concentration of H2O2 has an effect on the rate of the reaction

2. in the main experiment, the concentration of H2O2 should perhaps be
reduced, as the time taken for the reaction to reach an end point was
relatively short. An actual concentration may be deduced through
further trial experiments as this may also be effected by the other
reactants.

Experiment 2

Aim:

The aim of the following trial experiments is to investigate as in the
first trial, which of the remaining substances affect the rate of
reaction (i.e. KI, and H2SO4). Through carrying out these experiments,
I must also determine which concentrations must be used in the main
experiment, bearing in mind the following points:

1. Whichever substance is being investigated, the other reactants must
be in excess

2. When the strongest concentrations are being investigated, the time
for the reaction to reach its end point must be long enough to start
and stop the timer with accuracy. On the other hand, when the weakest
concentrations are being used, the time taken must not be so long that
valuable laboratory time is wasted.

Method

The method used was the same as in the first set of trials.

Results

a) Varying the concentration of KI.

Concentrations used:

H2O2 - 2.5M

H2SO4 - 1.0M

Na2S2O3 - 0.05M

Concentration of KI

(M)

Time taken for reaction to reach

End point. (secs)

1.0

8.22

0.5

17.55

b) Varying the concentration of H2SO4.

Concentrations used:

H2O2 - 2.5M

KI - 0.1M

Na2S2O3 - 0.05M

Concentration of H2SO4

(M)

Time taken for reaction to reach

End point. (secs)

1.0

2.15

0.5

2.11

0.1

2.20

Analysis of Results

After carrying out a second set of trial experiments, and taking into
account the results from the initial trial experiments, the following
points can now be deduced:

1. The rate at which this reaction takes place, is dependant on the
concentrations of:

-H2O2

-KI

2. Varying the concentration of H2SO4 had no effect on the rate of
reaction, as can be seen from the results. The presence of H2SO4 is
required however for the reaction to proceed, it can therefore be
deduced that it is acting as a catalyst, and is not a factor that
should be investigated.

Problem 2

Before the effects of temperature on reaction rate can be investigated
in the main investigation, it is important that the procedure used is
accurate, and uses the laboratory time most efficiently. Therefore, I
must conduct trial experiments to decide the best method to use, and
become acquainted with that method.

Experiment 3

Aim:

To determine the best method of conducting experiments investigating
how temperature affects the rate of reaction.

Apparatus:

· H2O2 - 2 Vol.

· KI -0.1M

· H2SO4 - 1.0M

· Na2S2O3 -0.005M

· 2% starch solution

· 50ml burettes (x5)

· conical flasks

· volumetric flasks

· stands and clamps

· funnels

· stopclock

· water bath (i.e. bunsen; tripod etc.)

· test tubes

Method

A water bath was set up and brought to a temperature of 50°C. Two test
tubes were taken, in the first was placed:

- 5ml H2SO4

- 5ml Na2S2O3

- 5ml KI

In the second tube was placed 10ml of H2O2; both tubes were placed in
the water bath, and left until their solutions were both at a
temperature of 50°C. Once the contents of both tubes had reached this
temperature, the were mixed and left in the water bath. The time taken
for the reaction to finish was recorded, and the experiment was
repeated for a second time at 50°C.

Results

The were no valid results gained from these trials, however they
illustrated that there were large inaccuracies in the method

Analysis of Results

Although there were no valid results gained from these trials it
proved that the method was flawed. Gaining the correct temperature of
both solutions using the water bath heated by a bunsen, was almost
impossible, the temperature fluctuated constantly. By the nature of
the water bath, it also meant that only very few experiments could be
heated and prepared at any one time. Bearing in mind that there is
limited lab time, this is also not ideal.

Therefore, in the main experiment, a thermostatic water bath will be
used that will accurately maintain the temperature of the water, and
also be capable of holding a large volume of test tubes at any given
time.

Modifications to Key Variables

Now that it has been made clear exactly which substances affect the
rate of the reaction, it can be stated exactly which variables are to
remain constant, and which to vary.

The following are factors that are to remain constant throughout this
experiment:

· Concentration of H2SO4 - varying the concentration does not affect
the rate of the reaction, however the presence of H+ ions is required
as they act as a catalyst. Therefore a 1.0M concentration will be used
throughout this investigation.

· Concentration of Na2S2O3 - the presence of thiosulphate ions is
required in this experiment as they indicate when the reaction is
complete, reacting with the iodine molecules that are produced.

I2 (aq) + 2S2O3 ¯ (aq) Õ 2I¯ (aq) + S406 (aq)

· Pressure (only really applicable with gases)

· Temperature - this will be kept at room temperature when
investigating the ways in which varying the concentration affects the
rate. However, part of this investigation is to investigate how the
temperature affects the rate of reaction, therefore this will be
varied when investigating the effects of temperature.

· Concentration - the concentration of all of the reactants will be
kept constant only when investigating the effects of temperature on
the rate of a reaction.

The following are factors that are to be varied, each independently of
each other; throughout this investigation:

· Concentration of H2O2 - through conducting the trial experiments, it
has been found that the concentration of H2O2 affects the rate of
reaction; therefore the concentration will be varied, and
investigated.

· Concentration of KI - this has also been found to affect the rate of
reaction; therefore the concentration will be varied, and
investigated.

· Temperature of reactants - this is a factor that affects all
chemical reactions in the form of the Q10 rule; therefore its effects
will be investigated.

Risk Assessment

Before I can start conducting the main experiment, the issue of
laboratory safety must be considered so that any possible exposure to
risk is avoided. The chemicals, quantities and techniques that are to
be used will all be looked at, and assessments will be made.

Procedures

Chemicals:

- H2SO4 (1.0M)

- H2O2 (2 Vol.)

- KI (1.0M)

- Na2S2O3 (1.0M)

Techniques:

- filling of burettes with chemicals stated

- measuring out of specific volumes of chemicals, using burettes

- Initiating chemical reaction through the mixing of two prepared
volumes of chemicals.

Hazardous chemicals and nature

· H2SO4 - this is and acid, and therefore should be handled with care,
even though low concentrations are being used. It should not be
swallowed or allowed to penetrate the skin as it may have toxic
properties at these concentrations. Contact with the skin or eyes must
also be avoided as it may be slightly corrosive, but more likely it
would cause irritation to the skin.

· H2O2 - this has similar properties to H2SO4 and therefore the same
precautions will be taken.

· Final reaction mixture - this should be handled with care as it
contains a mixture of unreacted substances as well as iodine. It
should be treated as having the same hazardous properties as above,
but also containing iodine which although in this form is not
particularly hazardous, has the ability to stain skin and clothing.

Protective measures

When carrying out this investigation and handling these chemicals,
certain laboratory safety procedures must be carried out at all times:

· Labs coat must be worn

· Goggles must be worn

· Work areas must be cleared of any unnecessary objects (e.g. bags;
books; stools)

The procedures mentioned previously, are all areas where chemicals are
being used, and there is certain amount of risk. Therefore when
conducting these procedures, extra care and vigilance must be taken;
acknowledging these possible risks should result in safe laboratory
work.

Units of Concentration

The units of concentration that will be used to measure out H2SO4, Na2S2O3
and KI are mol dm-3 or 'M'.

In the case of H2O2 however, the concentration will not be measured in
mol dm-3, but in 'Vol.'. Hydrogen peroxide is sold commercially in '20
Vol.' and '10 Vol.' solutions. A 20 Vol. solution of hydrogen peroxide
is one that liberates 20 times its own volume of oxygen when heated.

2H2O2 ® 2H2O + O2

2 x 34g. 22.41 at standard conditions

If the solution is a 2 Vol. solution of H2O2, (2 x 34g) of H2O2 must
be contained in a volume of solution, which is ½ of the volume of
oxygen it produces.

Therefore:

A 2Vol. solution contains (34 x 2) x 2 g of hydrogen peroxide per
litre;

22.4

that is 6.07g per litre of water.

Therefore, if it was required to calculate the concentration of H2O2
in mol dm-3 for the use in an equation for example, this could now be
done.

For 2 Vol. solution of H2O2, the concentration in terms of mol dm-3
is:

Moles = Mass/g

Molar mass/g mol-1

So:

Moles = 6.07 = 0.18

34

[IMAGE][IMAGE][IMAGE]If:

Concentration/mol dm-3 = Moles

Volume/ dm-3

Then:

Concentration/mol dm-3 = 0.18 = 0.18

1

Main Investigation - Implementing

1) An Investigation Into The Rate At Which Iodine Is Formed When The
Concentration Of Reactants is Varied

Aim

The first part of this investigation is to investigate how varying
concentration of H2O2 and KI affects the rate of the reaction.

Apparatus/Chemicals

In this main investigation, the following chemicals and apparatus will
be used:

· H2O2 - 2 Vol.

· KI -1.0M

· H2SO4 - 1.0M

· Na2S2O3 -0.005M

· 2% starch solution

· distilled water

· 100ml burette

· 50ml burettes (x4)

· conical flasks

· volumetric flasks

· stands and clamps

· funnels

· stopclock

Method

The five burettes were set up using the stands and clamps, and each
was filled with one of the chemical required for the investigation;
distilled water was placed in the 100ml burette. Each burette was then
labelled accordingly. The first concentration that was investigated,
was that of H2O2.

Two conical flasks were taken; in the first was measured out;

- 10ml H2SO4

- 10ml Na2S2O3

- 10ml KI

In the second conical flask was measured out 20ml of H2O2.

Once the two flasks had been prepared, the contents were mixed
together, and the stopclock was turned on simultaneously. Once the
reaction was completed and the instantaneous colour change was seen,
the stopclock was stopped immediately. The time was then recorded, and
the experiment was repeated three more times.

This procedure was again repeated another four times, except in the
second conical flask was placed a different concentration of H2O2, it
contained:

- 18ml H2O2

- 2ml distilled water

This gave an overall concentration 1.8 Vol. of H2O2.

This method was repeated another eight times, and each time the volume
of H2O2 in the second conical flask was reduced by 2ml, and the volume
of distilled water was increased by two. For each concentration, the
experiment was repeated four times; this enabled means and errors to
be calculated.

Once all of these experiments had been completed, the concentration of
KI was then varied. This was conducted in exactly the same way as the
H2O2, However, the total volume of KI used was only ever 10ml; the
same number of different concentrations were still tested though.

All of the results were recorded, and tabulated.

2) An Investigation into the Rate at Which Iodine Is Formed when The
Temperature of Reactants Is Varied

Aim

The second part of this investigation is to investigate how varying
the temperature of the reactants, affects the rate of the reaction.

Apparatus

· H2O2 - 2 Vol.

· KI -0.1M

· H2SO4 - 1.0M

· Na2S2O3 -0.005M

· 2% starch solution

· 50ml burettes (x5)

· conical flasks

· volumetric flasks

· stands and clamps

· funnels

· stopclock

· water bath (thermostatic)

· test tubes

Method

A thermostatically controlled water bath was set up and brought to a
temperature of 50°C. Eight test tubes were taken; in each of the first
four test tubes was placed:

- 5ml H2SO4

- 5ml Na2S2O3

- 5ml KI

In each of the remaining test tubes was measured out 20ml of H2O2, and
all of the eight tubes were placed into the water bath. Once all of
the solutions in all of the test tubes had reached the temperature of
50°C, each of the experiments was then conducted.

Two tubes were taken each containing one of the different solutions;
the tube containing the peroxide was then added to the second tube
containing the mixture as described above. As soon as the two
solutions had been added, the tube containing all of the reactants was
placed back into the water bath, and the stopclock was started.

Once the reaction had finished (as seen with an instantaneous colour
change from colourless to a deep purple) the stopclock was stopped,
and the time was recorded. The method was then repeated using each of
the remaining prepared test tubes; all of the results were recorded.

Results

All of the results gained from all of the experiments conducted will
now be presented in a series of tables and graphs. In each case,
through conducting a large number of experiments it has been possible
to gain Maximum, Minimum, and Average results, displaying the
occurrence of any anomalies, or inaccuracies in the results. Where
possible this has been shown in the graphs, in the form of error bars.
Where there are anomalous results that are clearly errors when
compared to the other results gained, these results will be discarded,
and not included when averages are being taken. If results are to be
discarded, then this will be clearly stated.

Table 1- Varying the Concentration of Peroxide

Concentrations/volumes used:

Sulphuric acid - 1.0M; 10cm3

Sodium thiosulphate - 0.005M; 10cm3

Potassium iodide - 1.0M; 10cm3

Hydrogen peroxide

Cm3

Water

cm3

Resulting

Conc.

Vol.

Reaction time s

Taverage

T1

T2

T3

T4

20

0

2.0

8.08

8.08

8.46

8.42

8.26

18

2

1.8

9.69

9.84

10.72

10.33

10.15

16

4

1.6

11.36

11.25

11.02

11.31

11.24

14

6

1.4

11.86

12.84

13.59

12.96

12.81

12

8

1.2

14.91

14.57

15.04

15.50

15.01

10

10

1.0

18.74

18.87

18.69

19.29

18.90

8

12

0.8

23.43

22.80

22.69

23.43

22.87

6

14

0.6

32.28

36.01

34.65

32.28

34.55

4

16

0.4

58.78

56.78

60.78

56.78

58.75

2

18

0.2

163.79

171.91

183.37

183.56

175.66

[IMAGE]Concentration of Peroxide /Vol.

Concentration of Peroxide /M

Taverage

/s

1

T /s-1

2.0

0.1786

8.26

0.1212

1.8

0.1607

10.15

0.0986

1.6

0.1429

11.24

0.0890

1.4

0.1250

12.81

0.0780

1.2

0.1071

15.01

0.0666

1.0

0.0893

18.90

0.0529

0.8

0.0714

22.87

0.0437

0.6

0.0536

34.55

0.0289

0.4

0.0357

58.75

0.0170

0.2

0.0179

175.66

0.0057

Table 2 - Varying the Concentration of Iodide Ions

Concentrations/volumes used:

Sulphuric acid - 1.0M; 10cm3

Sodium thiosulphate - 0.005M; 10cm3

Hydrogen peroxide- 2Vol.; 10cm3

[IMAGE]Potassium

Iodide

cm3

Water

cm3

Resulting

Conc.

cm3

Reaction time s

Taverage

1

T /s-1

T1

T2

T3

T4

10

0

1

8.08

8.08

8.46

8.42

8.26

0.1211

9

1

0.9

8.88

9.39

9.58

9.76

9.40

0.1064

8

2

0.8

10.73

10.71

11.01

10.64

10.77

0.0928

7

3

0.7

12.16

12.42

12.84

12.39

12.45

0.0803

6

4

0.6

13.80

13.99

14.23

14.22

14.06

0.0711

5

5

0.5

17.68

17.71

17.55

17.62

17.64

0.0567

4

6

0.4

22.45

22.01

21.84

22.35

22.16

0.0451

3

7

0.3

29.40

28.67

28.36

28.32

28.69

0.0349

2

8

0.2

43.01

41.28

41.49

40.39

41.54

0.0241

1

9

0.1

74.13

77.41

72.24

74.67

74.61

0.0134

Table 3 - Varying the Temperature

Temp.

°C

Reaction time s

Taverage

T1

T2

T3

T4

15

109.49

87.86

108.68

108.13

103.54

20.5

70.41

88.78

72.22

75.06

76.18

25

55.68

53.54

52.52

50.40

53.02

30

31.43

28.23

28.68

28.05

29.10

35

25.46

29.59

22.18

23.32

25.14

40

15.93

14.46

14.59

14.76

14.94

45

11.63

10.90

10.12

11.12

10.94

50

6.68

4.22

5.15

4.56

5.15

[IMAGE][IMAGE]Temperature

K

1

K

1

Taverage

ln (rate)

288

0.003572

0.0097

-4.64

293.5

0.003407

0.0131

-4.34

298

0.003356

0.0189

-3.97

303

0.003300

0.0344

-3.37

308

0.003247

0.0398

-3.22

313

0.003195

0.0670

-2.70

318

0.003145

0.0914

-2.39

323

0.003096

0.1941

-1.64

Analysis And Conclusions of Results

Now that all of the experiments have been conducted, the results
gained that have been displayed in the form of tables and graphs can
now be analysed. It may now be possible to draw certain conclusions
from the results about the nature of the reaction that has been
investigated.

1) Order Of The Reaction And The Rate Equation

As already stated in the background information gained at the start of
this investigation, it is possible to determine the order of the
reaction with respect to each of the reactants, and also the reaction
as a whole.

The first and most obvious point to make regards the order of reaction
with respect to H2SO4 and Na2S2O3. It was made clear from early on in
the investigation through background knowledge and trial experiments
the concentration of these two substances does not have an effect on
the rate of the reaction. Therefore it can be deduced that the order
of reaction with respect to:

- H2SO4 = Zero Order

- Na2S2O3 = Zero Order

Through conducting trial experiments however, it was soon deduced that
the concentrations of the remaining two reactants (H2O2 and KI) did
effect the rate of the reaction, and this can be clearly seen in the
results. To determine the nature of how these substances affected the
rate of the reaction, it was necessary to plot a number of graphs.

The first set of graphs to be drawn show concentration against time,
and were used to calculate half-lives for the various reactions
(Graphs 1 and 2). The second set of graphs drawn were identical to the
first set, but rather than calculating half-lives, they were used to
find the gradient at five points along each graph (Graphs 2 and 5).
Calculating the gradient at these points gave the rate of the reaction
at five different concentrations. From this a final set of graphs were
drawn showing the rate of reaction against concentration (Graphs 3 and
6).

After studying the results gained, and drawing these graphs, it was
found that out of all of the results recorded, all were included in
the final graphs and used to draw conclusions from. There were no
significantly anomalous results recorded, as all of them seemed to
follow the same pattern when plotted on the graphs. I felt that there
was no reason to ignore or discard any of these results.

For H2O2, Graph 1shows that the half-life increases very gradually,
which initially might suggest that the reaction is second order with
respect to H2O2. However, this increase is very gradual, and when
Graph 3 is studied, it is clear that the rate of the reaction is
directly proportional to the concentration. After studying Graph 1
carefully and checking its accuracy alongside the actual results, I
have decided that because the increase in the half-lives is only
gradual and Graph 3 clearly shows that the rate is proportional to the
concentration, this reaction is first order with respect to H2O2.

Graph 1 holds a number of possible areas where errors may occur and
this could explain why the half-life is not constant. The actual
process of drawing the graphs provides a number of sources for error.
Drawing an accurate curve freehand is quite difficult requiring a
steady and smooth action, and although this skill can be improved with
practise, it still provides sources of error. Also, the calculating of
the half-lives requires further drawing using a ruler that is only
accurate to the nearest millimetre.

Therefore, after studying the results gained from the experiments,
using chemical knowledge, and considering where any errors could have
occurred in this analysis, it can be deduced the order of the reaction
with respect to H2O2 is first order.

For KI, the results shown in the graphs 4, 5, and 6 show a similar
problem as found with H2O2, Graph 4 shows a gradually increasing
half-life, and Graph 6 shows clearly that the rate is proportional to
the concentration of KI. Therefore, for the same reasons as already
stated above for the reactions concerning H2O2, I deduce that the
order of the reaction with respect to KI is also first order.

It is now possible to express these results in the form of a rate
equation, this will show the order of the reaction with respect to the
individual reactants, but will also allow us to determine the overall
order of the reaction.

For the reaction studied in this investigation, the rate equation is:

Rate = k [H2O2]1[KI]1[H2SO 4]0[Na2S2O3]0

Rate = k [H2O2]1[KI]1

Therefore, the overall order for the reaction is second order.

The units of the rate constant, k are:

mol-1dm3s-1

2) Effects of Temperature On Rate Of Reaction, And The Arrhenius
Equation

After obtaining the results from investigating how temperature effects
the rate of the reaction, Graph 7 shows the ln(rate) against 1/T (K-1).
By measuring the gradient of this graph, a value for the activation
energy EA can be calculated. The relationship being used is known as
the Arrhenius equation:

ln k = C - EA (1/T)

R

The gradient of the graph is equal to -EA

R

Once all of the points had been plotted onto this graph, a line of
best fit was then drawn, giving the average and most accurate value
for the gradient. However, in order to consider the possible errors
that I may have made, I also drew two more lines on the same graph
giving the values of the maximum and minimum gradients.

The calculation below shows the activation energy that was calculated
from the line giving the average gradient on Graph 7. This is the
value that will be used as the final answer.

The gradient of the Graph 7 = -2.5 / 0.000367 = -6811.99

Therefore, the activation energy, EA, for this reaction =

6811.99 x 8.314 = 56634.88 J mol-1

= 56.63 kJ mol-1

Below are the calculations showing the possible maximum and minimum
activation energies for this reaction.

Maximum gradient = -1.375 / 0.000087 = -15865.38

Therefore, the activation energy, EA, for this reaction =

15865.38 x 8.134 = 131904.77 J mol-1

= 131.90 kJ mol-1

Minimum gradient = -0.4 / 0.0005 = -800

Therefore, the activation energy, EA, for this reaction =

800 x 8.134 = 6507.2 J mol-1

= 6.51 kJ mol-1

It can therefore be shown quantitatively by just three values from a
single graph the possible errors that might exist in the results
gained.

Maximum

EA(kJ mol-1)

Average

EA(kJ mol-1)

Minimum

EA(kJ mol-1)

Largest Possible Error

131.90

56.63

6.51

125.39

Q 10 Rule

The results gained from investigating how the temperature affects the
rate of the reaction can now also be used to test the Q 10 rule.

The table below showing just a selection of results gained clearly
reinforces the generally accepted rule that as there is a 10° rise in
temperature, the reaction time is halved.

Temperature

°C

Average Reaction Time

S

15

103.54

25

53.04

35

25.14

Evaluation

Now that the experiment has been conducted and the results have been
analysed, I must begin to evaluate the investigation, assessing
methodology and results, and identifying both errors and their
sources.

Firstly, the reliability of the results should be assessed; can these
results be relied upon to give conclusions that show the true
patterns, or trends that actually occur in this reaction? I believe
that there is no reason to doubt the methodology behind this
investigation, the results gained show what was required to see how
concentration and temperature affect the rate of reaction. Whether the
results compiled show exactly what was happening when the experiments
were being conducted, is another question because there may be some
doubt about the accuracy of the results.

After assessing the methods and procedures used throughout this
investigation, areas where errors may have occurred have been
identified, and this may explain some of the uncertainty experienced
when analysing the results. Firstly, the various solutions that were
being used were not all taken from the same batch of solutions. Due to
the allotment of laboratory time, it was impossible to use the same
batches of solutions throughout the whole experiment. This was because
during the periods when experiments were not being conducted, it was
possible that the solutions that were being used may have 'gone off',
and therefore new batches had to be made up. Therefore, each time a
new batch was made up, to say that it was exactly the same
concentration as the previous batch would be impossible. And so this
is a very real area where errors in the results may have occurred.
Modifications that could be made to increase the accuracy would be to
only make up one large batch of solutions, and conduct all of the
experiments in one go, taking up no more time than perhaps 48 hours.

Secondly, during the preparations of each experiment, the solutions
were measured out using burettes that measured to the nearest 0.1cm3
only, and relied on the careful control of the tap. Therefore there is
a possibility that volumes were not always measured to the accuracy
capable, this may have been due to bad technique, or possibly the fact
that there was limited time and a certain amount of pressure to
complete all of the practical work. Modifications that could be made
to perhaps increase the accuracy of the volumes measured could include
using more accurate burettes; spending more time on both practising
the technique, and conducting the actual experiment.

Deciphering exactly when the reaction was complete was not always as
clear as expected, and this may account for any possible errors. For a
number of the experiments conducted, the colour change that indicates
when the reaction is complete, was not always as instantaneous as
previously described. On these occasions, the colour change was
relatively slow, and this made it difficult to determine when exactly
the whole of the solution had changed colour. Therefore there may have
been some variation in actually deciding the end point of the
reaction, which may have lead to errors and inaccuracies in the
results. I am not sure why there was variation in the times it took
for the actual colour change to occur, and therefore modifications to
the method cannot be made, however this opens up the possibility for
further investigation.

When the ways in which temperature affected the rate of reaction was
investigated, it was decided to use a thermostatically controlled
water bath after modifying the original procedure. Although this was
much more accurate than using a bunsen to heat the water, the accuracy
of the water bath to maintain the desired temperature is questionable.
The water bath that was being used seemed to be temperamental and not
always particularly accurate, this may have lead to errors in the
results. Therefore a modification that might be made could be to
perhaps use a more sophisticated and reliable water bath.

Thus, as shown above there numerous possibilities where error may have
occurred. These areas of error must therefore be used to explain why
the results appeared as they did, making it difficult draw definite
conclusions.

Bibliography

'Chemistry Students Book' - Nuffield Advanced Science

'Chemistry In Context' - Graham Hill And John Holman

'Basic Inorganic Chemistry' - Cotton, Wilkinson, Gaus

'Inorganic And Physical Chemistry' - A. Holderness

'Advanced Chemistry' - P.R.S. Murray

'Chemistry In Focus' - John Andrew And Paul Rispoli

'Chemistry' - Ken Gadd And Steve Gurr



EXPERIMENT 10 Chemical Kinetics: Iodine Clock Reaction Objective: The complete concentration and temperature dependence of the reaction rate for the reaction between peroxydisulfate ion and iodide ion in aqueous solution will be determined. Introduction: Peroxydisulfate ion reacts with iodide ion in aqueous solution to give iodine and sulfate ion: 2 I-+ S 2 O 8 2-→ I 2 + 2 SO 4 2-(1) In this experiment you will determine the rate at which reaction (1) occurs. This rate can be expressed in the form of a rate law, Rate = k I " [ ] x S 2 O 8 2 " [ ] y where x and y are the orders of reaction in iodide and peroxydisulfate ions, respectively, and k is the specific rate constant; the specific rate constant is a function of temperature. The initial rate of this reaction will be determined by measuring the time required to generate a certain amount of iodine, the same amount in every trial, according to reaction (1). Two other reactions will be used to signal when this constant amount of iodine has been produced. The first of these two reactions involves thiosulfate ion, S 2 O 3 2-, a certain, identical amount of which will be added to the reaction mixture in each trial. Thiosulfate reacts with iodine as fast as iodine is produced, converting it back to iodide ion: 2 S 2 O 3 2-+ I 2 → S 4 O 6 2-+ 2 I-(2) Because reaction (2) is so fast, relative to reaction (1), iodine will not have a chance to build up in solution until all of the thiosulfate has been consumed. Any buildup of iodine in the solution indicates that the thiosulfate has been used up, and that, of course, means that the constant amount of iodine has been produced. To visually detect the presence of excess iodine in the solution you will add some starch to one of the solutions before mixing the iodide and peroxydisulfate. As soon as iodine begins to build up in solution it will react to form a dark-blue complex with starch: I 2 + starch → I 2-starch complex (3) The time period as measured from the time of mixing the peroxydisulfate with the iodide in the presence of starch and thiosulfate to the appearance of the dark-blue color is the time it

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