kinetics of the acid catalysed reaction between iodine and propanone final lol

Chemistry Individual Investigation Yann Perusset 13F

1

Kinetics of the Acid Catalysed reaction between Iodine and Propanone I shall base my investigation around the following reaction: I2(aq) + CH3COCH3 (aq) CH3COCH2I (aq) + H⁺ (aq) + I⁻ (aq)

Aims of the investigation By the end of this investigation I hope to have worked out the rate equation and the rate constant for the above reaction. From the rate equation I will than propose a possible mechanism of the reaction. Lastly, with all information and data already gathered, I will calculate the activation enthalpy of the reaction. The general steps that I will take to achieve my aim are:

1. With use of a colorimeter, I will measure the absorbance of known concentrations of iodine solutions to plot a calibration curve of absorbance against concentration of iodine

2. To deduce the orders of reaction with respect to Iodine, Sulfuric Acid and Propanone a graph of concentration of solution against time would be created again with use of a colorimeter. I can then work out the rate equation and rate constant. (via initial rates method)

3. Propose a possible mechanism which is consistent with the rate equation 4. Finally, I will vary the temperature of the reaction and record new rate constants in order to

calculate the activation enthalpy by using the Arrhenius Equation.

Background Theory Collision Theory

1

The basic idea of collision theory is that reactions occur when the particles of the reactants collide, provided they collide with a combined minimum kinetic energy that must be supplied to enable the bonds in the reactants to stretch and break. This minimum amount of kinetic energy is known as the activation enthalpy (EA). The graph below shoes an enthalpy profile for an exothermic reaction.

The theory also helps explain that the more collisions that occur above the activation enthalpy the faster the reaction is as there are more successful collisions. The graph below helps show the concept:


2

To increase the number of particles with enough energy for a successful collision the temperature can be increased. This concept is again shown in the distribution curve shown above, with T1 being the colder temperature and T2 being the warmer. The area under the curve past the EA shows the number of particles with enough energy to form a successful collision. Basically, reactions go faster at higher temperatures because a larger proportion of the colliding molecules have the minimum activation enthalpy needed to react. Many factors can affect the overall rate of a reaction, the main factors being concentration, temperature, particle size and the presence of a catalyst. When I come to working out the rate equation and I increase the concentration of one of the reactants it is important to keep all other factors the same to determine what effect that particular solution has on the rate without any outside influences. An increase in the concentration of a solution increases the amount of particles within it, the more particles there is, the more likely a successful collision. Order of reaction, rate constants and rate equations

2+3

Rate Equations and rate constant With a rate equation, it is possible to see how the concentration of each of the reactant affects the rate of a reaction. An example of a simple rate equation is shown below.

Rate=k[A]x[B]y

The total order of a reaction is the sum of the orders of the individual reactants. So for this reaction it would be the sum of x and y. From the rate equation it is also possible to find out the rate-determining step. From looking at the order of each of the substances it will tell you the relative number of moles of each substance involved in the rate determining step. For example, assume substance A in the equation above is first order, this means that there will be one substance in the rate-determining step. From this, it would then be possible to create a mechanism for the reaction. The rate constant can be calculated by rearranging the rate equation. Below is an example of the rate equation above but rearranged in terms of the rate constant, k:

k=

To obtain a reliable value for the rate the initial rates method would be used (shown below) Once the rate equation is worked out, it can be easily linked to the reaction mechanism by looking at the rate determining step. Colourimetery

4

The rate of a reaction is the change in concentration of a reactant or product divided by the time taken for the change to occur. Therefore to calculate the rate, you have to be able to measure the concentration of a solution. To do so I shall use the process of Colourimetery. A colorimeter can be used to measure the change in colour of a reaction. It works on the principle that coloured solutions (like iodine that I will be using) absorb certain wavelengths of light. The amount of light that is absorbed by the solution is known as the absorbance of the solution. The absorbance is proportional to the concentration of the solution used (Abs Conc) Known concentrations of the coloured solution are used to produce a calibration curve which can then be used to find the concentration of any absorbance value that has that coloured solution within it.


3

Theory of initial rates (initial rates method)2

Initial rates are used to measure the rate of reaction before any of the reactant has been used up. This is done by drawing a tangent at t=0 on a time against concentration graph, and working out the rate from the gradient of the tangent while assuming that the reactants have not changed from their starting concentrations. The graph below shows how to obtain an initial rate from a tangent:

Once the initial rate of a reaction is worked out, a graph of rate against concentration of reactant can be drawn. This would give the order of reaction with respect to a particular chemical. The graphs below show what the graph would look like depending on its order.

0 order 1

st Order

2

n Order


4

The Arrhenius Equation5+6+9

The Arrhenius Equation was first proposed my Dutch chemist J.H van ‘t Hoft in 1884 but later developed into a physical justification by Sweden’s Svante Arrhenius, the Arrhenius equation links the rate equation with reaction temperature. The equations is shown below:

Where k is the rate constant, Ea is the activation enthalpy, A is the pre-exponential factor, R is the gas constant and T is the temperature in Kelvin. The Ea of the reaction can be calculated by first working out the rate constant of the reaction at a variety of temperatures and plotting a graph of Ink against 1/T. The gradient of the graph would give you the Ea of the reaction. (below) The equations below give the Arrhenius Equation rearranged in terms of Ea. This can also be used to calculate the Ea of the reaction.

Ea=(lnk-lnA)RT The gradient gives the activation enthalpy of the reaction


5

Method Method for constructing calibration curve

1. Make up 0.005, 0.0005, 0.00005, 0.000005, 0.0000005 moldm-3

of iodine solution into small test tubes via distilling 10cm

3 of the previous concentration into a 100cm

3 volumetric flask

2. Label each small test tubes with allocated concentrations 3. Add 5cm

3 of distilled water to a very clean small test tube and reference the colorimeter

4. Add 5cm3 of 0.05 moldm

-3 iodine solution into a small test tube and take a reading (a reading

of 1.70 and 1.90 is needed for the most concentrated iodine solution) 5. Repeat readings for all concentrations (reference the colorimeter each time) 6. Plot a graph for absorbance against concentration (for a better scale –log[-I2] is used)

Apparatus Used Colorimeter with a 420-490nm filter- iodine absorbs strongly between these wavelengths 10cm

3 glass pipette- to make accurate concentrations with minimal uncertainty

100cm3 volumetric flask- to thoroughly dilute the iodine

Small test tubes- to fit in the colorimeter Deducing the order of reaction with respect to iodine

1. Dilute the 0.05moldm-3

iodine solutions into 0.01moldm-3

and fill around 20cm3 into a 50cm

3

glass beaker 2. Get 2.00moldm

-3 sulfuric acid and fill around 20cm

3 into a 50cm

3 glass beaker

3. With one experiment at a time, use the volumes shown in the table below to fill small glass test tubes. (don’t add the propanone until step 7)

Experiment Volume of 0.01 moldm

-3

iodine solution (cm

3)

Volume of 2.00 moldm

-3

sulfuric acid solution (cm

3)

Volume of 100% neat propanone solution (cm

3)

Volume of distilled water

Total volume of mixture (cm

3)

1 3.0 1.5 0.5 0 5.0

2 2.5 1.5 0.5 0.5 5.0

3 2.0 1.5 0.5 1.0 5.0

4 1.5 1.5 0.5 1.5 5.0

5 1.0 1.5 0.5 2.0 5.0

6 0.5 1.5 0.5 2.5 5.0

4. Shake iodine, acid and water to mix 5. Reference the colorimeter as described above 6. Add the allocated amount of propanone quickly but carefully into the mixture and place in

the colorimeter (quick shake to insure the propanone doesn’t stay at the top) and start recording the absorbance every 10 seconds.

Time (s) Abd (expt1) Abd (expt2) Abd (expt3) Abd (expt4) Abd (expt5) Abd (expt6)

0

10

20

30

40

50

60

70

80

90

100

110

120

7. Repeat for all 6 experiments 8. With use of the calibration curve graph, for each experiment complete the table below


6

9. Draw a graph for each experiment of concentration of iodine against time, working out the rate of the reaction with use of the initial rates method (in background theory section)

10. A graph of rate of reaction against initial concentration can now be drawn. The shape of the graph should indicate the order of reaction with respect to iodine Repeat readings should be carried out until a set of two/three close matching results are achieved. Apparatus Used -Colorimeter with a 420-490nm filter- iodine absorbs strongly between these wavelengths. Has to be the same one as use for the calibration curve. -10cm

3 glass pipette- to create the 5cm

3 mixtures in the small

test tubes 100cm

3 volumetric flask- to thoroughly dilute the iodine and propanone

Small test tubes- to fit in the colorimeter 50cm

3 glass beakers

Stopwatch Deducing the order of reaction with respect to Sulfuric Acid and Propanone The exact same process is used again but instead only varying the solution of either the Sulfuric Acid. The tables below show what measurements to use: For deducing the order of reaction with respect to sulfuric acid


-3

iodine solution (cm

3)


-3


3)


3)



3)

1 0.5 3.0 1.5 0 5.0

2 0.5 2.5 1.5 0.5 5.0

3 0.5 2.0 1.5 1.0 5.0

4 0.5 1.5 1.5 1.5 5.0

5 0.5 1.0 1.5 2.0 5.0

6 0.5 0.5 1.5 2.5 5.0

Deducing the order of reaction with respect to propanone can be found online as repeating the same experimental process would be unnecessary.

Time (s) Abs Corresponding –log[-I2]

concentration

0

10

20

30

40

50

60

70

80

90

100

110

120


7

Deducing the activation enthalpy of the reaction This is what the small test tubes contain.


-3

iodine solution (cm

3)


-3


3)


3)


3)

1 3.0 1.5 0.5 5.0

For this reaction it is the same process as before. Yet done at different temperatures.

1. In small test tubes, create as many mixed solutions that will fit in an aluminium test tube rack 2. Put the solutions into a water bath and use the dial to heat up the water bath 3. In a separate test tube, the propanone has to be heated at the same temperatures as the

iodine sulfuric acid solution 4. The temperature of the small test tube solutions has to be taken and not the temperature of

the water bath as they would vary 5. Reference the colorimeter with distilled water and a clean test tube. 6. When the test tube solution reaches the allocated temperature, it is taken out of the water

bath, wiped dry, and placed in the colorimeter as quickly and as safely as possible 7. With use of a graduated 1cm pipette get 0.5 cm of propanone at the same temperature as

the iodine/sulphuric acid solution and add it to the small test-tube in the colorimeter. 8. Once propanone is added, start the stopwatch, and quickly shake the small test tube to

ensure the propanone does not stay at the top 9. Take absorbance readings every 10 seconds

This is to be repeated for all temperatures from 25°C to 50°C going up in 5°C each time. The reason why the water bath can’t be heated above 50°C because the boiling point of propanone is around 54°C so the experiment wouldn’t work. With use of the Arrhenius Equation, the EA can be worked out. Apparatus Used Colorimeter with a 420-490nm filter (same as used for the calibration curve) Water bath Thermometer 2cm graduated pipette for Sulphuric Acid 5cm graduated pipette for Iodine solution 1cm graduated pipette for propanone Stopwatch Apparatus used and why Graduated pipettes- for high accuracy and minimal percentage error when creating test solutions at a small scale Water bath- a controlled way to get to a wanted temperature and heat up solutions Volumetric flasks- an accurate way to make solutions of a wanted concentration


8

Risk Assessment7+8

Chemicals Used Propanone: Hazard/Risk

Contact with the eyes can cause serious, permanent damage.

Propanone is highly flammable, and presents a serious fire risk

Propanone is harmful if you swallow or inhale it. Long-term exposure, for example through breathing in the fumes, can cause liver damage

Repeated skin exposure may lead to defatting and irritation Control Measures

Always wear safety glasses

Ensure that ventilation is good so that you do not breathe in the vapour

Remove all sources of ignition from the working area

Always store bottles of propanone in a flame-proof cabinet

Do not leave propanone, or solutions containing it, standing unattended on a bench where they might get knocked over or set alight

Iodine: Hazard/Risk

Iodine is very toxic if swallowed or inhaled

Iodine has a significant vapour pressure at room temperature which can lead to the build-up of dangerous levels of iodine vapour

Exposure to iodine may lead to reproductive damage

Iodine may be absorbed through the skin Control Measures

Wear safety glasses

Wear gloves if skin contact is likely Sulfuric Acid: Hazard/Risk

Contact with the eyes or skin can cause serious permanent damage

Concentrated solutions of acid are extremely corrosive Control Measures

Always wear safety glasses

Do not allow the acid or a solution of it to come into contact with your skin

When diluting acid always wear eye protection, and ALWAYS add acid to water (not the reverse) slowly and with great care.

Use constant stirring (sulfuric acid is much denser than water, and if you do not stir when adding acid to water, a layer of concentrated acid may form at the bottom of the beaker, creating a substantial temperature gradient where acid and water meet)


9

Reaction of propanone with iodine: Hazard/Risk

The product (iodopropane) is strongly irritant to the eyes. Control Measures

Dispose of it immediately by pouring it down a foul-water drain.

Wear eye protection Apparatus Used Glass Pipettes Hazard/Risk

Rick of breakage and injury to hand when inserting into pipette fillers Control Measures

Insert carefully and hold the pipette as close as possible to the pipette filler Thermometer Hazard/Risk

The mercury inside the thermometer is highly toxic Control Measures

The thermometer must be handled with absolute care to prevent breakage Water Bath Hazard/Risk

There is a rick as the plug may get in contact with water Control measures

Insure the bath isn’t overfilled and everything is kept dry

References 1. Salters Advanced Chemistry: Chemical Ideas, eds. D.Denby, C. Otter and K. Stephenson, Heineman, Oxford, Third

Edition, 2008, p.210-215 2. Salters Advanced Chemistry: Chemical Ideas, eds. D.Denby, C. Otter and K. Stephenson, Heineman, Oxford, Third

Edition, 2008, p.217-220 3. Kenneth Connors, Chemical Kinetics, 1990, VCH Publishers, pg. 14 4. Salters Advanced Chemistry: Chemical Ideas, eds. D.Denby, C. Otter and K. Stephenson, Heineman, Oxford, Third

Edition, 2008, p.372 5. Laidler, K. J. (1997) Chemical Kinetics,Third Edition, Benjamin-Cummings

6. http://old.iupac.org/goldbook/A00446.pdf 7. http://cartwright.chem.ox.ac.uk/hsci/chemicals/hsci_chemicals_list.html 8. Hazcards 9. http://www.chem.fsu.edu/chemlab/chm1046course/activation.html


10

Analysis Making and using the calibration curve Below are the results I gained for constructing a calibration curve. The –log was used to turn the standard log into positive figure making the graph easier to follow.

Conc of Iodine Conc Log Conc -log[-I2] / moldm-3 Absorbance

0.009 -2.04576 2.045757491 1.99

0.005 -2.30103 2.301029996 1.44

0.001 -3 3 0.42

0.0005 -3.30103 3.301029996 0.25

0.00025 -3.60206 3.602059991 0.14

0.0001 -4 4 0.09

0.00005 -4.30103 4.301029996 0.03

The graph created from these results gave a good exponential calibration curve. By plotting the graph in excel, it also gave the equation of the curve, which was . Yet by rearranging the equation in terms of x, the concentration of iodine can be calculated with a

given absorbance figure. In terms of x, the equation is

. Yet this would only give the

–log of the concentration of Iodine in the solution, meaning that [I2] . This gives an overall formula for the concentration of iodine in a solution as:

[I2]

Where Abs is the absorption reading obtained from the colorimeter with a 290nm filter. By using this equation, the concentration of Iodine in the solution can be calculated accurately and without human error by trying to read it from a graph. It can also be quickly calculated in a spreadsheet making it a lot quicker to calculate multiple concentrations instead of doing each one individually. The spreadsheet formula to use is =10^-(((LN(E5))-(LN(81.929)))/(-1.772)) where E5 is the absorption reading obtained from the colorimeter.

y = 81.929e-1.772x

0

0.5

1

1.5

2

2.5

2 2.5 3 3.5 4 4.5

Ab

sorb

ance

Conc -log[-I2] / moldm-3

Calibration Curve Absorbance


11

Deducing orders of reactions The aim of the investigation was to deduce the overall order of the reaction as well as the order of reaction to each individual reactant involved. Later, this would help in calculating the rate constant. The rate equation before any orders of reaction are calculated is:

(the H

+ ions coming from the sulfuric acid)

Deducing the order of reaction with respect to Iodine


-3

iodine solution (cm

3)


-3


3)


3)



3)

1 3.0 1.5 0.5 0 5.0

2 2.5 1.5 0.5 0.5 5.0

3 2.0 1.5 0.5 1.0 5.0

4 1.5 1.5 0.5 1.5 5.0

5 1.0 1.5 0.5 2.0 5.0

6 0.5 1.5 0.5 2.5 5.0

The Iodine graphs correspond to the experiments listed in the table above and the results from the results section. The gradient of the line equals the rate of the reaction because rate=conc/time and the gradient of the line gives conc/time. Below gives the found gradient and therefore the rate for each experiment.

Run1 Experiment Rate(moldm-3s-1)


1 0.0000242 1 0.0000276

2 0.0000175 2 0.000025

3 0.0000163 3 0.0000246

4 0.0000217 4 0.000025

5 0.0000205 5 0.000025 6 0.0000188 6 0.0000214

Each experiment number corresponds to a concentration of Iodine. This can be calculated by: Overall concentration of iodine =conc x (volume of iodine/total volume) The concentration used was 0.001 moldm-3 and the total volume was cm-3

Experiment Concentration of Iodine

1 0.006

2 0.005

3 0.004

4 0.003

5 0.002 6 0.001

From this, a graph of concentration of iodine against rate of reaction can be drawn.


12

Run1

Run2

The two graphs show a relatively flat line, hence showing that the order of reaction with respect to iodine is 0 order as when the concentration of Iodine increases in the reaction, the rate remains unchanged at an average of 0.0000198 moldm-3s-1 for Run 1 and 0.0000248 for Run 2. Therefore, Iodine is not included in the rate equation as no matter what the concentration of iodine is used, the rate of reaction remains unchanged.

0

0.000005

0.00001

0.000015

0.00002

0.000025

0.00003

0 0.001 0.002 0.003 0.004 0.005 0.006 0.007

Rat

e/m

old

m-3

s-3

Concentration of Iodine/moldm-3

0

0.000005

0.00001

0.000015

0.00002

0.000025

0.00003

0 0.001 0.002 0.003 0.004 0.005 0.006 0.007

Rat

e/m

old

m-3

s-3

Concentration of Iodine/moldm-3


13

Deducing the order of reaction with respect to Sulfuric Acid


-3

iodine solution (cm

3)


-3


3)


3)



3)

1 0.5 3.0 1.5 0 5.0

2 0.5 2.5 1.5 0.5 5.0

3 0.5 2.0 1.5 1.0 5.0

4 0.5 1.5 1.5 1.5 5.0

5 0.5 1.0 1.5 2.0 5.0

6 0.5 0.5 1.5 2.5 5.0

The Sulfuric Acid graphs correspond to the experiments listed in the table above and the results from the results section. The gradient of the line equals the rate of the reaction because rate=conc/time and the gradient of the line gives conc/time. Below gives the found gradient and therefore the rate for each experiment.



1 0.0001875 1 0.0002045

2 0.0001731 2 0.0001731

3 0.0001286 3 0.000125

4 0.0001047 4 0.00009375

5 0.00007031 5 0.00006429 6 0.00003774 6 0.00003636

Each experiment number corresponds to a concentration of Sulfuric Acid. This can be calculated by: Overall concentration of Sulfuric Acid =conc x (volume of Sulfuric Acid/total volume) The concentration used was 2.00 moldm-3 and the total volume was cm-3

Experiment Concentration of Sulfuric Acid

1 1.2

2 1.0

3 0.8

4 0.6

5 0.4 6 0.2

From this, a graph of rate of reaction against concentration of iodine can be drawn.


14

Run 1

Run 2

From these graphs which show proportionality between the concentration of sulfuric acid and the rate. Hence showing that the order of reaction with respect to Sulfuric Acid is 1

st order as when the

concentration is increased, the rate increases proportionally. Therefore, sulfuric acid is involved in the rate equation.

y = 0.000165x

0

0.00005

0.0001

0.00015

0.0002

0.00025

0 0.2 0.4 0.6 0.8 1 1.2 1.4

Rat

e/m

old

m-3

s-1

Concentration of Sulfuric Acid /moldm-3

y = 0.000167x

0

0.00005

0.0001

0.00015

0.0002

0.00025

0 0.2 0.4 0.6 0.8 1 1.2 1.4

Rat

e/m

old

m-3

s-1

Concentration of Sulfuric Acid /moldm-3


15

The order of reaction with respect to propanone is also 1st

order meaning that it is also included in the rate equation. This leaves an overall equation of the reaction to be:

Meaning that the rate constant, k, can be calculated by rearranging the rate equation to:


16

Deducing the rate constant of the reaction One of my aims was to calculate a rate constant for the reaction between acid catalysed reaction between iodine and propanone. I have all ready worked out the rate equation, , and in

terms of k,

. In order to calculate the constant as accurately as

possible, I can use the rate against concentration graph from Sulfuric Acid. The gradient of the line

gives

, therefore by combining this into the main rate equation,

, where g is

the gradient of the rate against concentration for sulfuric acid and [propanone] is the concentration of propanone used. Calculating the concentration of propanone used As I used neat propanone in my experiments, I need to be able to calculate the concentration used before I can work out the rate constant of the reaction. The density of liquid propanone is 0.7925 g/cm

3, meaning that in one dm

3 there is 792.5g. The

amount of moles can then be calculated:

Molecule formula= C3H6O Mr= 58.08 g mol

−1

Mass=792.5g

Moles in 1dm

3= 13.64mol

Therefore the concentration of propanone I used 13.64moldm

-3.

Since I used 1.5cm3 of neat propanone, the concentration used in each experiment:

=conc x (volume of Propanone/total volume) =13.64moldm

-3 x (1.5cm

3/5cm

3)

=4.092moldm-3

Deducing the rate constant The overall rate constant can now be calculated. The average gradient from the rate against concentration of sulfuric acid from both run 1 and run 2 is 0.000166. The concentration of the propanone used is 4.092 moldm

-3, these figures can be used to calculate the rate.

k=4.057 10-5 So the rate constant, k, for the acid catalysed reaction between iodine and propanone is 4.057 10

-5.


17

Deducing the activation enthalpy, Ea, of the reaction Experiment Volume of

0.01 moldm

-3

iodine solution (cm

3)


-3


3)


3)


3)

1 3.0 1.5 0.5 5.0

These mixtures were repeated at temperature intervals of 5

o. From the results obtained, graphs of

concentration of iodine against time can be drawn, from them, using the initial rates method; the rate of reaction for each experiment at different temperatures can be calculated. Below is a table of the rate of reaction at different temperatures. The temperatures have to be converted from Celsius to Kelvin so they work when used in an Arrhenius graph. The gradient of the graphs gave these results:

Temperature (C

o) Temperature (K) Rate (moldm

-3s

-1)

30 303.15 0.000021

35 308.15 0.000033

40 313.15 0.00004

45 318.15 0.0000741

50 323.15 0.0001013 From these rates, the rate constant for each temperature can be calculated. This done with the equation:

Where the concentration of both sulfuric acid = 0.6moldm-3

and propanone =0.682moldm-3

.

Temperature, T, (K) Constant calculation Constant, k, value

303.15

0.00005132

308.15

0.00008065

313.15

0.00009775

318.15

0.00018109

323.15

0.00024756

Now T has to converted to 1/T (reciprocal of T) and k to lnk (the natural log of k)

T k 1/T lnk

303.15 0.00005132 0.0032987 -9.877

308.15 0.00008065 0.0032452 -9.425

313.15 0.00009775 0.0031934 -9.233

318.15 0.00018109 0.0031432 -8.617

323.15 0.00024756 0.0030945 -8.304

From the constant values, I can now draw a graph of against

.


18

Below shows the graph of lnk against 1/T.

The gradient of the line

hence , where g is the gradient and R is ideal gas constant.

By using , the Ea can be calculated: G=-7739.7 R= 8.314 Jmol-1K-1 Jmol-1 KJmol-1 The activation enthalpy of the reaction was found to be 64.348 KJmol

-1. Yet it is also possible to

propose the full Arrhenius expression of the reaction by calculating a value for the pre-exponential factor, A, in the Arrhenius equation. The x intercept on the Arrhenius graph is equal to the natural log

of A. The x intercept is 15.637 so . Therefore, . This gives the value of A to be 6.18x10

7. The overall Arrhenius expression is therefore:

A=6.18x10

7

Jmol-1 R= 8.314 Jmol-1K-1

(the overall Arrhenius expression)

y = -7739.7x

-10

-9.8

-9.6

-9.4

-9.2

-9

-8.8

-8.6

-8.4

-8.2

0.003050.00310.003150.00320.003250.00330.00335

lnk

1/T


19

Proposing a possible mechanism Below shows a possible mechanism of the reaction:

After the analysis of the reaction, I have the order of reaction with respect to each reactant and so can prove whether the mechanism is correct or incorrect. As Iodine is zero order, it is not involved in the rate determining step as it has to effect on the rate, therefore the rate determining step must include propanone and a H

+ (from S2HO4) as those are the

only two reactant which affect the rate. The rate determining step also has to be the slowest step, which on the diagrams, is the first. Therefore, the rate determining step is:

This is the rate determining step because is both the slowest and contains the only reactants which can affect the rate. Although there is a significant possibility that this is what does happens in the reaction, it is simply a proposition of the mechanisms which best describes what actually happens.


20

Evaluation General limitation of practical procedure For my experiment there are several limitations to consider. The most troubling limitation would be transferring the propanone to the small test tube. The problem is that as soon as the propanone is added the mixture, the reaction is already taking place. This therefore affects the accuracy of the initial rate. Another unavoidable limitation of the practical procedure is that the mixture has to be quickly shaken before put into the colorimeter to ensure the mixture is thoroughly mixed. This affects the consistency of the experiments as some may be mixed more than others by human error. Regarding the practical procedure in use for the water bath experiments, the limitation is that there is a lot to do in a short space of time, so again consistency of the experiment may not be very good. Limitation of practical equipment It is impossible to have a 100% accuracy for any practical equipment meaning that there are uncertainties involved. Yet the overall uncertainty can be reduced by selecting the appropriate equipment when necessary. Below shows all the equipment which I used and the limitations involved when using it.

Equipment Uncertainty Limitations

Volumetric flask 100cm

3 0.08

Human error involved when assuring the bottom of the meniscus is level with the 100cm

3 line

Grad Pipette 5cm3 0.05

Liquid solution can drip out and the human error involved when assuring the bottom of the meniscus is level with the line





Colorimeter 0.05

Same colorimeter has to be used throughout to insure your calibration curve works. The amount of light in the room could also potentially affect the absorbance reading.

Thermometer 0.5 Few limitations apart from the glass tends to steam up making it hard to read the temperature.

Water Bath The only limitation would be that it takes a long time for the solutions to reach the correct desired temperature

Pipette 10cm3 0.04

Simple to use, yet when taking some liquid, extra has to be taken to compensate for some lost when taking the pipette out of the solution.


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Percentage Uncertainties To gain a good indication of the errors involved in the practical procedure, percentage uncertainties can be calculated. Iodine concentration uncertainty

As I had to dilute the concentration of iodine to 0.01moldm-3

, there would be uncertainties involved in this made concentration. This involved using a 20cm

3 pipette with an uncertainty of 0.06 and also

using a 100cm3volumetric flask with an uncertainty of 0.08. Therefore, the uncertainty of the

0.01moldm-3

concentration is:

=2.4x10

-6

So the uncertainty is very small. Absorbance experiment uncertainty As the colorimeter has an uncertainty of 0.05, there is also an uncertainty involved in the. Below shows an example of what the uncertainty is. Time Abs % Uncertainty

0 10 1.47 3.401

20 1.43 3.497

30 1.35 3.704

40 1.31 3.817

50 1.25 4.000

60 1.19 4.202

70 1.15 4.348

80 1.09 4.587

90 1.04 4.808

100 1.00 5.000

110 0.93 5.376

120 0.89 5.618

It is clear to see that when the absorbance is low, the uncertainty is higher. Fortunately, i avoided using experiments which included a low absorbance so that the percentage uncertainty would not be too high therefore improving the accuracy of my practical.


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Method Overall, I feel that my method worked out well as I was able to achieve all my aims which I stated in my plan to a high level of accuracy. Finding various rates of reaction All worked well and I was able to determine the rate and order of reaction with respect to each reactant. Fining the Activation Enthalpy This worked well and gave a good result of KJmol-1. Although it was difficult to do, it turned out good. Alterations that I would change If I was to change anything I would probably try and use a data logger to create accurate and on time readings. This would reduce human error. Yet, there is not many things that I would actually change, as the method and practical procedure all went to plan. Reliability I feel that my results are reliable to use, as I repeated experiments until a match was found. From this, a good average could be attained. I also had to experiment to get the “perfect combination” for the mixture. The graphs showed very few anomalies hence further proving that my results are reliable. Accuracy Accuracy is a measurement of how close the result values are to the true value. The two values which I could compare to true values would be the rate equation and the Ea. The true rate equation is Rate=k[propanone] [H

+]. Mine was found to be the same, showing my results were accurate.

Secondly, my activation enthalpy of the reaction. The true value was found to be 68.2kJmol-1

and my value was KJmol-, so only being 5.6% off the true value, it is fair to say my results are highly accurate. The reason that my value may just be a little bit to low was probably because my temperatures were said to be lower than they actually are. Validity Overall, I feel that my results re valid because they have a relatively low percentage uncertainty. Also, my values and results match what they should be.


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Conclusion

Aims of the investigation By the end of this investigation I hope to have worked out the rate equation and the rate constant for the above reaction. From the rate equation I will than propose a possible mechanism of the reaction. Lastly, with all information and data already gathered, I will calculate the activation enthalpy of the reaction. The general steps that I will take to achieve my aim are:

1. With use of a colorimeter, I will measure the absorbance of known concentrations of iodine solutions to plot a calibration curve of absorbance against concentration of iodine

2. To deduce the orders of reaction with respect to Iodine, Sulfuric Acid and Propanone a graph of concentration of solution against time would be created again with use of a colorimeter. I can then work out the rate equation and rate constant. (via initial rates method)

3. Propose a possible mechanism which is consistent with the rate equation 4. Finally, I will vary the temperature of the reaction and record new rate constants in order to

calculate the activation enthalpy by using the Arrhenius Equation. Looking back at my aims I am happy to say that my investigation went very well. I deduced the order of reaction with respect to iodine and propanone, proposed a rate equation and therefore a rate constant, calculated the activation enthalpy of the reaction and lastly proposed a mechanism. Below shows all of the aims of the investigation worked out:

k=4.057 10-5 KJmol-1

kinetics of the acid catalysed reaction between iodine and propanone final lol

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