kinetics of the harcourt eseen reaction (steve quarshie) part 1

Stephen AhiabahU6RJD

Kinetics of the Harcourt-Essen Reaction (Hydrogen peroxide variation)

Aims:

To find the order of reaction with respect to the hydrogen peroxide, Potassium iodide and the sulphuric acid by the use of a clock reaction

Calculate rate constant, mechanism and equation Find the effects of temperature on the rate of reaction The effects of using a catalyst on the rate of reaction Find the activation enthalpy of the reaction, with and without a catalyst.

Background knowledge

Hydrogen peroxide and potassium iodide equation:

Basic reaction without spectator ions

H2O2 + 2I– + 2H+ → I2 + 2H2O1

Iodine Clock reaction 2

The reaction whose rate varies according to the concentration of hydrogen peroxide (H2O2) is a slow reaction involving the hydrogen peroxide, hydrogen from the acid, and iodide ion from the potassium iodide:-

H2O2 + 2 H+ +2 I-→ I2 + H2O

Without any sodium thiosulphate Na2S2O3 present the colourless reactants would gradually form a brown colour when the iodine appears and it would be difficult to time the reaction.

When the thiosulphate ion is present there is a second very fast reaction between the iodine formed and the thiosulphate ion:-

I2 + 2S2O32- → 2I- + S4O6

2-.

This second reaction is so fast that the I2 is consumed instantaneously and no brown iodine colour appears until the S2O3

2-ions are used up. (The S2O32-as well as the S2O3

2-ion is colourless). Starch solution is added to the mix because it forms a black complex with Iodine I2.

The iodine clock reaction is a chemical reaction in which two colourless solutions are mixed and react together to form a brown/red colour. However, initially, the iodine will be of a small concentration, and will appear very light in colour, and therefore, the production of iodine will

1 Chemistry Individual Investigation (F336) - Starter

2 http://www.techknow.org.uk/wiki/index.php?title=Iodine_clock


be very hard to detect. So by the addition of starch, which instantaneously turns dark blue with the formation of iodine ions, gives a far more accurate time for the production of iodine ions and provide a representation of the rate of reaction. As the reaction demonstrates that reaction rates depend on the concentrations of the reagents involved in the overall reaction.

The time required to reach this point depends on the rates of the first two reactions, and consequently on the concentrations of all the reactants. Anything that increases the first reaction (e.g. Temperature) will shorten the time taken for the blue complex to appear. Therefore, increasing the concentration of iodide, hydrogen peroxide, or will accelerate the reaction. However, increasing the thiosulphate concentration will have the opposite effect; it will take longer for the blue colour to appear.

Rate of reactions and measurement of rates: 3

Rates of reaction are measured by monitoring the rate of change of an observable property. In this reaction, the rate is measured by adding a known amount of sodium thiosulfate solution and some starch indicator to the mixture and measuring the time it takes for the solution to turn blue.

Some methods of measuring the rate of reaction that are applicable to this experiment are as follows:

• Colorimetry measures the intensity of a colour in a reaction against time. I.e. in this clock reaction, it is timed until a colour change occurs when a certain amount of product is formed in the rate determining step. The intensity of the colour depends on the concentration of iodine. The absorption reading provides a measure of the colour intensity and therefore the concentration of iodine. In order to gain more accurate results I could use a calibration curve using solutions of a known concentration of iodine, then read off the exact concentrations of iodine that corresponds to particular absorbance reading from the curve. By taking absorbance reading s at different intervals I can obverse how the concentration of iodine changes as the reaction proceeds. And this is timed

• Titrimetric analysis of concentrations (quenching); titrations used to measure changing concentrations of a reactant or product. In this experiment, samples are taken from the reaction mixture at specific intervals. These samples are the then quenched by a hydrogen carbonate solution to neutralise the acid so the reaction stops. The concentration of iodine in these samples is measured by a titration wit sodium thiosulphate solution. And this is timed

The rate of reaction from both these methods are simply calculated by this reaction

Concentrations and rates of reactions

3 Salter Advanced chemistry module analysis TL 4.3 pg 54-55


Collision theory 4

This is the idea that reactions occur when particles of the reactant collide and collide with minimum kinetic energy. As the particles approach and collide, kinetic energy is converted into potential energy and the potential energy of the reactants rises. Existing bonds start to stretch and break and new bonds start to form. This minimum kinetic energy is known as activation enthalpy and the collision must overcome the activation enthalpy, if a reaction is going to result from the collision.

In relation to my investigation, to know how varying concentration of hydrogen peroxide, iodide and hydrogen ions will affect the rate of reaction, and therefore time taken for starch and iodine (formed as a result of the reaction between (H2O2 + 2I– + 2H+) to form a blue complex), application of collision theory is needed.

When hydrogen peroxide, iodide and hydrogen ions collide, they must collide with energy that exceeds the activation enthalpy. Or the collision will not result in a reaction

This then means that; increasing the concentration of the reactant substances of (potassium iodide, hydrogen peroxide and sulphuric acid) per unit of volume will increase the number of likely successful collisions per unit of time. This increase number of successful collisions will exceed the activation enthalpy of the reaction. Rate of reaction increases, as the concentration of reactants of increases.

Reactants used

Hydrogen peroxide5

Hydrogen peroxide is a clear liquid that is slightly more viscous than water. It is a powerful oxidising agent, and so, is a strong bleaching agent. It can be used as disinfectant and as a monopropellant in rockets. In this reaction, it is used as an oxidising agent.

The formula for hydrogen peroxide is H2O2 and it has a pH of 4.5.

Hydrogen peroxide often decomposes exothermically in the presence of light, and so, it needs to be stored in a cool environment in a dark coloured to ensure that that, in the experiment, the hydrogen peroxide solution will need to be freshly made up every day. If this doesn’t occur then the result of the data can be affected

It decomposes into water and oxygen spontaneously, as indicated by the following reaction:

2H2O2 2H2O + O2 + energy

The rate of decomposition is dependent on the temperature and the pH of chemicals present in the reaction. Hydrogen peroxide is incompatible with many substances which catalyse its decomposition, including most of the transition metals and their compounds. The decomposition of hydrogen peroxide is more likely in alkaline conditions, so, often; acid is

4 http://revisionworld.co.uk/a2-level-level-revision/chemistry/physical-inorganic-chemistry/kinetics/collision-theory5 http://en.wikipedia.org/wiki/Hydrogen_peroxide


added as a stabiliser. The sulphuric acid used in the reaction will mean that there is a much smaller chance for the hydrogen peroxide to decompose. Also, I will make sure that the hydrogen peroxide is newly made every day, and so, there is little chance for it to decompose.

Potassium Iodide6

Potassium iodide is a white, crystalline salt.

The formula for potassium iodide is KI

Potassium iodide acts as a simple ionic salt, K+I-. Since iodine is a mild reducing agent, potassium iodide can be easily oxidised by the hydrogen peroxide. Potassium iodine also forms the complex I3

- when combined with iodine. Potassium iodide can be used in photography, to prepare the silver (I) iodide. It can also be used in medicine, to protect the thyroid from radioactive iodine.

Sodium Thiosulphate7

Sodium thiosulphate is a colourless crystalline compound.

It has a molecular formula of Na2S2O3 and is more commonly found in its pentahydrate state (attached to water), Na2S2O3.5H2O (water of crystallisation).

In this experiment, I am provided with sodium thiosulphate in its pentahydrate state

Rate constants/equations: 8

6 http://en.wikipedia.org/wiki/Potassium_iodide7 http://en.wikipedia.org/wiki/Sodium_thiosulfate8 Chemical Ideas 10.3 pg.214http://www.chemguide.co.uk/physical/basicrates/arrhenius.html


Rate equations show the relationship between the rate of reaction and the powers to which the reactant concentrations are raised. If two reactants, A and B, react together then measurements of reaction rates with different concentrations of A and B

• K is the rate constant, which is constant at a particular temperature

• [A] and [B] are the concentrations of substances A (Iodide ion) and B( hydrogen peroxide)

• a is the order with respect to A (Iodide ion), and b is the order for B( hydrogen peroxide)

• The overall order of reaction is the sum of the individual orders, a+b

Order of reaction: 9

We can use results from such graph to then work out the:

9http://www.pearsonschoolsandfecolleges.co.uk/AssetsLibrary/SECTORS/Secondary/PDFs/Science/ EdexcelScience/ALevelRevisionGuides/EdexcelA2ChemistryRG_9781846905964_pg8-17_web.pdf

1st order reaction due to direct linear relationship

0 order reaction. Rate isn’t affected by concentration. Not all reactants are involved in the rate determining step

Either 2nd or 3rd order reaction, concentration values need to be squared or cubed to find out whether the relationship is linear or not

http://www.pearsonschoolsandfecolleges.co.uk/AssetsLibrary/SECTORS/Secondary/PDFs/Science/EdexcelScience/ALevelRevisionGuides/EdexcelA2ChemistryRG_9781846905964_pg8-17_web.pdf



Mechanisms10

When all the relevant molecules collide at once, they happen in steps. The slowest step controls how fast the overall reaction occurs – it is called the rate-determining step. Kinetic measurements establish the order of the reaction for each species. For a reaction between A, B and C the rate equation could be:

Rate = k [A] [B-] [C]

This rate equation demonstrates that:

• The reaction is first order with respect to A – so A is involved in the rate-determining step (Iodide ion)

• The reaction is first order with respect to substance B – so 1 mole of B are involved in the rate-determining step. (Hydrogen Peroxide)

. The reaction is first order with respect to substance C – so 1 mole of B are involved in the rate-determining step. (H+)

Effect of temperature on rates

When the temperature increases, the rate of a reaction increases too because the rate constant increases. The rate constant k is only a constant for a particular temperature. Changing the temperature changes the value of k because the proportion of molecules that have the required energy (greater than the activation energy) is increased and the colliding particles have a greater average energy. The Arrhenius equation shows the relationship between the rate constant k and the temperature T (in kelvin).

Where A is the frequency factor, k is the rate, Ea Is the activation energy, R is the gas constant (8.314 J K-1 mol-1), and T is temperature in Kelvin (K). The frequency factor has the same units as k.

10 http://www.lpscience.com/classes/apchemistry/baugher/printnotes/Reaction%20Mechanism.pdf

http://www.lpscience.com/classes/apchemistry/baugher/printnotes/Reaction%20Mechanism.pdf


By plotting ln(k) vs 1/T, Ea and A can be determined:

Where Ea = the activation energy

Comparing this to y =mx + c,

Y = ln(k) c = ln(A) x = 1/T m = -Ea/R

Therefore, the activation enthalpy can be found by:

Gradient =

-Gradient x R = Ea11

In relation to my investigation; as temperature increases, the reactant molecules of hydrogen peroxide, iodide and hydrogen ions gain more kinetic energy.

This increased movement as a result in in an increased kinetic energy will then increase the number of likely successful collisions that will exceed the activation enthalpy. So we can deduce that the rate of reaction is increased, as temperature increases.

11http://www.pearsonschoolsandfecolleges.co.uk/AssetsLibrary/SECTORS/Secondary/PDFs/Science/ EdexcelScience/ALevelRevisionGuides/EdexcelA2ChemistryRG_9781846905964_pg8-17_web.pdf




The Maxwell-Boltzmann Distribution can be used to show the distribution of energy of the reactant particles of hydrogen peroxide, iodide and hydrogen ions12

From the Maxwell Boltzmann Distribution Curve its seen that, as the temperature is increased, the number of the reactant molecules of hydrogen peroxide, iodide and hydrogen ions colliding exceeding the activation enthalpy will increase, meaning that fewer of these reactant molecules are colliding with energy less than the activation enthalpy and therefore, there are more successful collisions.

Catalysts

A substance that increases the rate of a chemical reaction without undergoing any permanent chemical change itself.

Catalysts increase the rate of a reaction by increasing the number of successful collisions. One method of doing this is to provide an alternative way for the reaction to happen which has a lower activation enthalpy.

12 http://www.webchem.net/notes/how_far/kinetics/maxwell_boltzmann.htm

http://www.webchem.net/notes/how_far/kinetics/maxwell_boltzmann.htm


The addition of a catalyst has exactly this effect on activation energy. A catalyst provides an alternative route for the reaction. This “alternative route” has a lower activation energy than the previous un catalysed route. This is represented on an energy profile as:

Homogenrous catalysts 13

Homogeneous catalysis occurs when the catalyst used in the same state (phase) as the reactants.

Homogeneous catalysis of the reaction between thiosoulpahte ions and iodide ions

Thiosulphate ions; S2O82- are strong oxidising agents. In contrast Iodide ions are very easily

oxidised to iodine. The reaction occurs very slowly in solution. This is because the reaction occurs between 2 negatively charged ions therefore the laws of attraction would deem that a high activation enthalpy is required for both these the ions to overcome the repulsion of their like charges to have a successful collision that will result in a reaction

The catalysed reaction however will prevent this repulsion from occurring. The catalyst I will be using is iron (II) chloride which will be added to the reaction solution. I am using this catalyst because transition metal compounds the ability to change oxidation state.

The thiosulphate ions oxidise the iron(II) ions to iron(III) ions. And in return the thiosulphate ions are reduced to sulphate ions.

The iron(III) ions are strong enough oxidising agents to oxidise iodide ions to iodine. In the process, they are reduced back to iron(II) ions again.

13 http://www.chemguide.co.uk/physical/catalysis/introduction.html (homogeneous catalysts)

http://www.chemguide.co.uk/physical/catalysis/introduction.html


Both of these individual stages in the overall reaction involve collision between positive and negative ions. This will be much more likely to be successful than collision between two negative ions in the uncatalysed reaction.

Outline Plan

Firstly add the same, fixed volume and concentration of sodium thiosulphate solution together with a starch solution 14 to Hydrogen peroxide and potassium iodide. The Iodide ions will be produced by the main reaction between hydrogen peroxide, potassium iodide and sulphuric acid immediately reacts with the thiosulphate ions:

I2 + 2S2O32– → S4O6

2– + 2I–

When all of the thiosulphate has been used up, the iodine accumulates in the solution and reacts with the starch to give a distinctive blue-black colour. The time from mixing the reactants to the appearance of the blue colour is therefore the time for a fixed concentration of iodine to be formed.

To find the concentration I2, I will change the volumes of the Hydrogen peroxide, Sulphuric acid, Potassium Iodide, in specific mixtures individually. When the volume of one reactant is changed then the other 2 volumes of reactants are kept constant. By altering this variable we can work out the concentration of I2 and ultimately the initial rate of reaction. This will be repeated 5 times to improve the reliability of the data and an average will be calculated.

Results from this will allow me to plot an “initial rate of reaction against the concentration of Iodine” graph, which will then allow me to calculate the order of reaction and rate equation.

To observe the effects of temperature and calculate the activation energy, the temperature will be varied 5 times and ensuring the concentration of the reactants remains constant. This will be repeated 5 times to improve the reliability of the data and an average will be calculated

I will also observe the effects of using a catalyst alongside varying the temperature of the reaction.

From this, the effects of temperature and the use of a catalyst on the rate of reaction can be observed.

The activation energy of the reaction with and without a catalyst can then be calculated, using the equation: Ae –Ea/RT as shown in the background knowledge.

14 November 1996 issue of Chemistry Review, Volume 6, Number 2, Pages 14 and 15.


Detailed Plan

List of Apparatus

Electronic Scale

Micropipette

Pipette

Beakers

Volumetric flasks

Burette

conical flasks

Droppers

Spatulas

Filter funnel

Electronic timer

Wash bottle

Hydrogen peroxide solution

Potassium iodide powder

Sodium thiosulphate powder

Distilled water

Starch solution

Water bath

Water bath

Thermometer

Stock solutions 15

Potassium iodide solution, 0.10 mol dm –3

Sulphuric acid, 1.0 mol dm–3

Sodium thiosulfate solution, 0.0050 mol dm–3, 6.0 cm3

Hydrogen peroxide solution, 0.10 mol dm–3

Starch solution, 1.0 cm3

Iron (II) chloride solution 1.0 mol dm–3 , 10 cm3

15 November 1996 issue of Chemistry Review, Volume 6, Number 2, Pages 14 and 15.


Preparations of solutions 16

Primarily I must prepare the solutions of both Potassium Iodide and sodium thiosulphate as they are made from solids that need to be dissolved in distilled water and also the hydrogen peroxide solutions

Mass required = volume required x concentration required x Mr

Preparation of 0.005 mol dm –3 of Sodium thiosulphate

1.24 g of solid sodium thiosulphate powder is needed.

a. This mass is calculated using; Moles=Mass/Mr

b. Moles required= 0.005, Mr of Sodium thiosulphate penhydrate = 248.18g/mol

c. Therefore d. Mass= Moles x Mre. Mass= 0.005mol x 248.18/molf. Mass= 1.24g (3sf)

90cm3 of sodium thiosulphate solution is needed as I am making up 5 mixtures in 3 repeats all with constant volume of 6cm3

As we are provided with volumetric flasks or 250cm3, I will make the solution up to 250cm3 also to compensate for errors.

1. Using a spatula 1.24g of sodium thiosulphate will be weighed in top pan balance on a weighing boat

2. Contents of the weighing boat will be transferred into a 150cm3 beaker3. Wash the weighing boat with the solvent (distilled water) in to the 150cm3 beaker4. Add approx. 100cm 3 of distilled water into the 150cm3 beaker5. Stir the solution with a glass rod to ensure the sodium thiosulphate has dissolved

into the distilled water.6. Wash the glass rod with distilled water into the 150cm3 beaker7. Using a funnel transfer the solution in the 150cm3 beaker into a 250cm3

volumetric flask.8. Wash both the contents of the funnel and the beaker into the 250cm3 volumetric

flask.9. Make up the solution 1 cm from the mark. Then using a petite pipette add distilled

water into he volumetric flask drop by drop up to the hark 10. Invert the volumetric flask 5 times to ensure all contents of the solution are

equally mixed

16 http://www.slideshare.net/wkkok1957/iodine-clock-reaction-experiment-using-potassium-iodide-and-hydrogen-peroxide

http://www.slideshare.net/wkkok1957/iodine-clock-reaction-experiment-using-potassium-iodide-and-hydrogen-peroxide



11. This forms 0.0500 mol dm-3 of sodium thiosulphate solution. 12. Transfer the solution into a burette that has been washed with distilled water that

has been set upright and clamped in a clamp stand

Preparation of 0.01 mol dm –3 potassium iodide

1.66 g of solid sodium potassium iodide is needed

a. This mass is calculated using; Moles=Mass/Mr

b. Moles required= 0.01, Mr of potassium iodide = 166.0g/molc. Therefore d. Mass= Moles x Mre. Mass= 0.015mol x 166g/molf. Mass= 1.66 (3sf)

375 cm3 of potassium is needed as I am making up 5 mixtures in 3 repeats with varying volumes between 5 cm3-25cm3

As we are provided with volumetric flasks of 500cm3, I will make the solution up to 500cm3 also to compensate for errors.

1. Using a spatula 1.66g of potassium iodide will be weighed in top pan balance on a weighing boat

2. Contents of the weighing boat will be transferred into a 500cm3 beaker3. Wash the weighing boat with the solvent (distilled water) in to the 500cm3 beaker4. Add approx. 400cm 3 of distilled water into the 500cm3 beaker5. Stir the solution with a glass rod to ensure the sodium thiosulphate has dissolved

into the distilled water.6. Wash the glass rod with distilled water into the 500cm3 beaker7. Using a funnel transfer the solution in the 500cm3 beaker into a 500cm3

volumetric flask.8. Wash both the contents of the funnel and the beaker into the 500cm3 volumetric

flask.9. Make up the solution 1 cm from the mark. Then using a petite pipette add distilled

water into the volumetric flask drop by drop up to the hark 10. Invert the volumetric flask 5 times to ensure all contents of the solution are

equally mixed11. This forms a 0.01 mol dm–3 potassium iodide solution12. Transfer the solution into a burette that has been washed with distilled water that

has been set upright and clamped in a clamp stand

Preparation of the Hydrogen Peroxide 0.1 mol dm –3 solutions


1. The volume of hydrogen peroxide required is 250 cm3 = 0.250 dm3 and the concentration required is 0.100 mol dm-3

Moles of hydrogen peroxide required = concentration required x volume required= 0.10 x 0.25= 0.0250 mol

2. The volume required would then equal the number of moles/concentration given. In this experiment, the concentration of hydrogen peroxide given is 1.76 mol dm-3. (20vol)17

3. Volume required = mol/given concentration= 0.00750/1.76= 0.14205 dm3 x 1000=14.21cm3

4. This will be obtained using a graduated pipette (with pipette filler) and transferred to a 250cm3 volumetric flask. The beaker will then be repeatedly washed using distilled water and these washings will be added to the volumetric flask. Further distilled water will be added until the solution is made up to 250 cm3

5. The hydrogen peroxide will have to be made up each day, as it decomposes in the presence of sunlight. It will need to be stored in a dark and cool place to kerb the possibility of decomposition and minimise it so that results are not affected.

Preparation of the Sulphuric acid 1.0 mol dm –3 solution

1. The volume of sulphuric acid required is 150 cm3 = 0.150 dm3 and the concentration required is 1.00 mol dm-3

Moles of sulphuric acid required = concentration required x volume required= 1.0x 0.150= 0.250 mol

2. The volume required would then equal the number of moles/concentration given. In this experiment, the concentration of sulphuric acid given is 1.00mol dm-3.

3. Volume required = mol/given concentration= 0.150/1.00= 0.150 dm3 x 1000=150cm3

4. This will be obtained using a graduated pipette (with pipette filler) and transferred to a 150cm3 volumetric flask.

Variables

17 Chemistry Individual Investigation (F336) – Pack


Variable Measured Method of controlling/measuring variable

Dependent Variable

Rate of reaction The rate of reaction will be calculated by taking the reciprocal of the average time taken for the complex to form.

Time taken for blue-black starch complex to form The time taken will be measured using a digital stopwatch. To reduce random errors, three readings will be taken and an average will be calculated.

Independent Variable

Concentration of hydrogen peroxide solution The solution of known concentration will be prepared. The volume used will be accurately measured using a pipette

Concentration of sulphuric acid solution added The solution of known concentration will be prepared. The volume used will be accurately measured using a pipette

Concentration of potassium iodide used The solution of known concentration will be prepared. The volume used will be accurately measured using a pipette


Controlled Variables

Volume of starch solution used 5 drops of starch will be added for each sample.

Temperature of reactants Temperature of the reactants was kept at a constant by conducting the experiment at room

Concentration and volume of sodium thiosulpahte The solution of known concentration will be prepared. The volume used will be accurately measured using a pipette

Methodology and Tables of volumes and concentrations 18 19

Changing H 2SO4 Volumes

Mixture Volume of H2SO4(aq)/ cm3

1.0 mol dm–3

Volume of starch solution/ cm3

Volume of H2O (aq)/ cm3

Volume of KI(aq)/ cm3

0.10 mol dm –3

Volume of Na2S2O3

(aq)/ cm3

0.0050 mol dm–3

Volume of H2O2 (aq)/ cm3

0.10 mol dm–3

1a 10.0 1.0 0.0 25.0 6.0 10.02a 8.0 1.0 2.0 25.0 6.0 10.03a 6.0 1.0 4.0 25.0 6.0 10.04a 4.0 1.0 6.0 25.0 6.0 10.05a 2.0 1.0 8.0 25.0 6.0 10.0

Formula for Concentration of sulphuric acid solution used:

18 Chemistry Individual Investigation (F336) - Starter Page 119 Salter Advanced chemistry module analysis TL 4.2 pg 52-53

http://en.wikipedia.org/wiki/Oxygen

http://en.wikipedia.org/wiki/Sulfur

http://en.wikipedia.org/wiki/Sodium

http://en.wikipedia.org/wiki/Sulfate

http://en.wikipedia.org/wiki/Hydrogen




Total stock volume = (10+1+25+6+10)/1000 = 0.052 dm3

Worked Example

Concentration of mixture 1a; When 10 cm3 of 1.0M H2SO4 used:

Concentrations of H 2SO4

Mixture Concentration (moldm-3)1a 1.92 x 10-1

2a 1.54 x 10-1

3a 1.15x 10-1

4a 7.70 x 10-2

5a 3.84 x 10-2

Varying concentrations of H 2SO4 and Iodine Clock reactions 20 21

Set up the 5 Burettes with the stock solutions of H2SO4(aq) 1.0 mol dm–3, H2O (aq), KI(aq) 0.10 mol dm –3, H2O2 (aq) 0.10 mol dm–3 and Na2S2O3(aq) 0.0050 mol dm–3in each, with every burette held upright in a clamp stand.

A) Measure the temperature of the room in order to keep it as constant as possible as, temperature is a control variable

B) In a beaker labelled X, add sulphuric acid and water, vary the concentration of sulphuric acid solution, by the use of the burette and the volumes of sulphuric acid and water stated in the specific table. Then add 5 drops (1cm3) of starch.

C) For the first mixture 1a, add 10.0 cm3 of sulphuric acid and 0.00 cm3 of water to beaker X, repeat for the other mixtures following the volumes shown on the table

D) Repeat this with the use of burette, for mixture 1a by adding, 6.0 cm3 Sodium thiosulfate solution, 0.0050 mol dm–3 , 10cm3 Hydrogen peroxide solution, 0.10 mol dm–3 and 25.0 cm3 Potassium iodide solution, 0.10 mol dm –3 and label this beaker Y. repeat for the other mixtures following the volumes shown on the table

E) Add the contents of Beaker X and Beaker Y into 1 beaker. Immediately start a stop clock

20 Salter Advanced chemistry module analysis TL 4.2 pg 52-5321 Chemistry Individual Investigation (F336) - Starter Page 1













F) Swirl contents of beaker steadily, ensuring a reaction occurs and that all contents are evenly mixed.

G) As soon as the main reaction has produced Iodine to react with the fixed volume of starch, a blue complex will form and stop the timer

H) Take 3 repeats for each concentration and calculate an average

I) As the time taken for the blue complex to appear is representative of the rate of reaction, we can now plot a “rate of reaction versus concentration” graph.

Changing KI Volumes


1.0 mol dm–3




0.10 mol dm –3

Volume of Na2S2O3 (aq)/ cm3

0.0050 mol dm–

3


0.10 mol dm–3

1b 10.0 1.0 0.00 25.0 6.0 10.02b 10.0 1.0 5.00 20.0 6.0 10.03b 10.0 1.0 10.0 15.0 6.0 10.04b 10.0 1.0 15.0 10.0 6.0 10.05b 10.0 1.0 20.0 5.00 6.0 10.0

Formula for Concentration of I- in potassium iodide solution used:

Total stock volume = (10+1+25+6+10)/1000 = 0.052 dm3

Worked Example:

Concentration of mixture 1B; when 25cm3 of 0.1M KI used:







Concentrations of KI

Mixture Concentration (moldm-3)1b 4.81 x 10-2

2b 3.84 x 10-2

3b 2.88 x 10-2

4b 1.92 x 10-3

5b 9.62x 10-3

Varying concentrations of KI and Iodine Clock reactions 22 23



B) In a beaker labelled X, add potassium iodide and water, vary the concentration of potassium iodide solution, by the use of the burette and the volumes of potassium iodide and water stated in the specific table. Then add 5 drops (1cm3) of starch.

C) For the first mixture 1b, add 25.0 cm3 of Potassium iodide and 0.00 cm3 of water to beaker X, repeat for the other mixtures following the volumes shown on the table

D) Repeat this with the use of burette, e.g. for mixture 1b by adding, 6.0 cm3 Sodium thiosulfate solution, 0.0050 mol dm–3 ,10cm3 Hydrogen peroxide solution, 0.10 mol dm–3 and 10 cm3 sulphuric acid solution, 1.0 mol dm –3 and label this beaker Y. repeat for the other mixtures following the volumes shown on the table

E) Add the contents of Beaker X and Beaker Y into 1 beaker. Immediately start a stop clock












Changing H 2O2 Volumes


1.0 mol dm–3




0.10 mol dm –3

Volume of Na2S2O3(aq)/ cm3

0.0050 mol dm–

3


0.10 mol dm–3

1c 10.0 1.0 0.0 25.0 6.0 10.02c 10.0 1.0 2.0 25.0 6.0 8.03c 10.0 1.0 4.0 25.0 6.0 6.04c 10.0 1.0 6.0 25.0 6.0 4.05c 10.0 1.0 8.0 25.0 6.0 2.0

Formula for Concentration hydrogen peroxide solution used:

Worked example:







Concentration of mixture 1B; when 20cm3 of 0.1M KI used:

Concentrations of H 2O2

Mixture Concentration (moldm-3)1c 1.92 x 10-2

2c 1.53 x 10-2

3c 1.15 x 10-2

4c 7.70 x 10-3

5c 3.84 x 10-3

Varying concentrations of H 2O2 and Iodine Clock reaction 24 25



B) In a beaker labelled X, add hydrogen peroxide and water, vary the concentration of hydrogen peroxide solution, by the use of the burette and the volumes of hydrogen peroxide and water stated in the specific table. Then add 5 drops (1cm3) of starch.

C) For the first mixture 1c, add 10 cm3 of hydrogen peroxide and 0.00 cm3 of water to beaker X, repeat for the other mixtures following the volumes shown on the table

D) Repeat this with the use of burette, e.g. for mixture 1c by adding, 6.0 cm3 Sodium thiosulfate solution, 0.0050 mol dm–3 ,10cm3 sulphuric acid solution, 1.0 mol dm–3

and 25 cm3 Potassium iodide solution, 0.10 mol dm –3 and label this beaker Y. repeat for the other mixtures following the volumes shown on the table

E) Add the contents of Beaker X and Beaker Y into 1 beaker. Immediately start a stop clock after












Finding the order of reaction

From the equation:

I2 + 2S2O32- → 2I- + S4O6

2-.

The volume and concentration of (thiosluphate) S2O32 is known as it is a controlled variable.

Moles of Na2S2O8 can be calculated as:

To work out the moles of Iodine used up we now refer to the equation:

I2 + 2S2O32- → 2I- + S4O6

2-.

In the reaction 2 moles thiosulphate ions react with 1 mole of iodine.

The reaction is 2:1 and the moles of iodine used in the clock reaction, is calculated to be simply:

To then work out the concentration iodine used we refer to the equation m=c x v and rearrange as shown below:


To calculate the initial rate for each mixture we use the equation

We now tabulate the results from these reactions which will look like this:

Mixture [I-](mol dm-3)

Clock time(s)

Rate(mol dm-3 s-1)

Temp(°C)

1x2x3x4x5x

We can now plot a graph of the initial rate of reaction against the concentration of Iodide ions for each mixture

We can now determine the order of reaction from the shape of the graph (see page 3).

Percentage progress calculations

Refer to the equations in background knowledge

Main reaction equation: H2O2 + 2I– + 2H+ → I2 + 2H2O

Clock reaction equation: I2 + 2 S2O32− → 2 I− + S4O6

2−

When varying the concentration and volume of sulphuric acid H2SO4




Reactant Volume(dm3) Moles (mol) Moles of I2 produced (mol)

H2SO4 Maximum= 0.01Minimum= 2.0x10–3

Maximum= 0.01Minimum= 2.0x10–3

Maximum= 5x10–3

Minimum= 1x10–3

KI 0.025 2.5x10–3 1.25x10–3

H2O2 0.01 1x10–3 1.10–3

Moles are calculated using the equation:

Moles (mol) = concentration (mol dm- 3) x volume (dm3 )

The mole ratio of this reaction 1:2:2

H2O2 has the lowest concentration out of all the reactants therefore this is the limiting reagent.

In the clock reaction;

Moles of Sodium Thiosulphate = concentration (mol dm- 3) x volume (dm3 )

= 0.005 mol dm- 3 x (6/1000) dm3

= 3x10–5 mols

Mole ratio in clock reaction is 1:2

Therefore mols of iodine = 3x10–5 /2

=1.5 x10–5

% Progress is given as

% Progress of maximum number of moles: % Progress of minimum number of moles:

When varying the

concentration and volume of Potassium Iodide KI





KI Maximum= 0.025Minimum= 5.0 x10–3

Maximum=2.5 x10–3

Minimum= 5x10–4Maximum= 1.25x10–3

Minimum= 2.5x10–4

H2SO4 0.01 0.01 5.0 x10–3

H2O2 0.01 1.0x10–3 1.10–3




KI has the lowest concentration out of all the reactants therefore this is the limiting reagent.



= 0.005 mol dm- 3 x (6/1000) dm3

= 3x10–5 mols


Therefore moles of iodine = 3x10–5 /2

=1.5 x10–5



When varying the concentration and volume of Hydrogen peroxide H2O2





H2O2 Maximum= 0.01Minimum= 2.0x10–3

Maximum=1.0x10–3

Minimum= 2.0x10–4Maximum= 1.0x10–3

Minimum= 2.0x10–4

H2SO4 0.01 0.01 5.0 x10–3

KI 0.025 2.5x10–3 1.25x10–3




H2O2 has the lowest concentration out of all the reactants therefore this is the limiting reagent.



= 0.005 mol dm- 3 x (6/1000) dm3

= 3x10–5 moles


Therefore moles of iodine = 3x10–5 /2

=1.5 x10–5



Varying the temperature of the reaction




Without a catalyst.

Set the apparatus in the same way when I was investigating the effect of changing in concentration of all reactants on rate of reaction.


In changing the temperature, I will use the mixture 1a from my previous hypothesis of changing the concentrations of reactants, when I change the temperature of the reaction


1.0 mol dm–3




0.10 mol dm –3

Volume of Na2S2O3

(aq)/ cm3

0.0050 mol dm–3


0.10 mol dm–3

1a 10.0 1.0 0.0 25.0 6.0 10.0

Concentrations

Reactant Concentration (mol dm–3)Sulphuric Acid (H2SO4(aq)) 1.92 x 10-1

Potassium Iodide (KI(aq)) 4.81 x 10-2

Hydrogen Peroxide (of H2O2 (aq)) 1.92 x 10-2

By using a water bath, the temperature will be altered

These amounts should be kept constant, to ensure the test is fair. The same volume of these solutions should always be used, as we are measuring the effect of temperature on rate of reaction.

A) For the mixture 1a, add 10 cm3 of sulphuric acid and 0.00 cm3 of water to beaker XThen add 5 drops (1cm3) of starch.

B) Repeat this with the use of burette, for mixture 1a by adding, 6.0 cm3 Sodium thiosulfate solution, 0.0050 mol dm–3 , 10cm3 Hydrogen peroxide solution, 0.10 mol dm–3 and 25 cm3 Potassium iodide solution, 0.10 mol dm –3 and label this beaker Y.

C) Place the beakers X and Y in water bath, so their temperature is increased to what its needed to be

D) Add the contents of Beaker X and Beaker Y into 1 beaker. Immediately start a stop clock

E) Swirl contents of beaker steadily, ensuring a reaction occurs and that all contents are evenly mixed.














F) As soon as the main reaction has produced Iodine to react with the fixed volume of starch, a blue complex will form and stop the timer. Place this new beaker above a white tile so the subjective error is reduced and the colour change is seen more promptly.

G) Ensure that the temperature remains constant. a thermometer can be used to monitor the temperature


With a catalyst.

Set the apparatus in the same way when I was investigating the effect of changing in concentration of all reactants on rate of reaction.


In changing the temperature, I will use the mixture 1a from my previous hypothesis of changing the concentrations of reactants, when I change the temperature of the reaction


1.0 mol dm–3




0.10 mol dm –3

Volume of Na2S2O3

(aq)/ cm3

0.0050 mol dm–3


0.10 mol dm–3

1a 10.0 1.0 0.0 25.0 6.0 10.0

Concentrations

Reactant Concentration (mol dm–3)Sulphuric Acid (H2SO4(aq)) 1.92 x 10-1

Potassium Iodide (KI(aq)) 4.81 x 10-2

Hydrogen Peroxide (of H2O2 (aq)) 1.92 x 10-2

As mentioned in my background knowledge the catalyst will be using is Iron(II) chloride solution.

By using a water bath, the temperature will be altered














These amounts should be kept constant, to ensure the test is fair. The same volume of these solutions should always be used, as we are measuring the effect of temperature on rate of reaction.

A) For the mixture 1a, add 10 cm3 of sulphuric acid, 0.00 cm3 of water and also using a petite pipette add 2 drops of Iron(II) chloride to beaker XThen add 5 drops (1cm3) of starch.

B) Repeat this with the use of burette, for mixture 1a by adding, 6.0 cm3 Sodium thiosulfate solution, 0.0050 mol dm–3 , 10cm3 Hydrogen peroxide solution, 0.10 mol dm–3 and 25 cm3 Potassium iodide solution, 0.10 mol dm –3 and label this beaker Y.

C) Place the beakers X and Y in water bath, so their temperature is increased to what its needed to be

D) Add the contents of Beaker X and Beaker Y into 1 beaker. Immediately start a stop clock

E) Swirl contents of beaker steadily, ensuring a reaction occurs and that all contents are evenly mixed.

F) As soon as the main reaction has produced Iodine to react with the fixed volume of starch, a blue complex will form and stop the timer. Place this new beaker above a white tile so the subjective error is reduced and the colour change is seen more promptly.

G) Ensure that the temperature remains constant. a thermometer can be used to monitor the temperature

H) Take 3 repeats for each temperature and calculate an average.

For both the catalysed and un-catalysed reaction a table will need to be drawn to show the effects of temperature and the catalyst on the rate of reaction.

Temperature of reactants(˚C)

Temperature of reactants (Kelvin) (T)

Clock times(s)(Test1-5)

Average Clock time(s)

Rate of Reaction (Ts-1)

20 293.1530 298.1540 303.1550 308.1560 313.5


Using Arrhenius equation, we draw up the following table to calculate the activation energy of the reactions: 26

Temperature of reactants(˚C)

Temperature of reactants (Kelvin) (K)

1/ Temperature(T-1)

Rate of Reaction (Ks-1)

Ln(K)

20 293.15 3.41 x 10-3

30 303.15 3.29 x 10-3

40 313.50 3.18 x 10-3

50 323.15 3.09 x 10-3

60 333.15 2.99 x 10-3

1. Take the natural logarithms of both sides of the Arrhenius equation. The Arrhenius equation is:

a. K = Ae –Ea/RT

b. This gives ln k = ln A - Ea/(RT).

c. Rearrange the equation into the formula of a straight line in the form y = mx + c. this therefore gives :

ln k = -Ea/(RT) + constant

Calculate K by using the rate equation and rearranging to isolate the rate constant k.

d. Use the rearranged equation to plot a graph of In(k) vs. 1/ to get

e. (- Ea/R) as the gradient

f. Divide the gradient by the universal gas constant to isolate the value (- Ea) from (- Ea/R).

g. Multiply (- Ea) by -1 to get Ea, the activation energy of your reaction. This forms the basis of the experimental determination of activation energy

26 http://www.ehow.com/how_7805881_activation-energy.html#ixzz2BRYMYgZN

http://www.ehow.com/how_7805881_activation-energy.html#ixzz2BRYMYgZN


Risk assessmentPrior to starting the investigation, it is imperative to consider the potential risks involved in the investigation and they would need to be removed from the experiment or minimised as much as possible.

Methods in which to minimise the possible hazards are shown below

1. Wear goggles to protect eyes

2. Tuck in loose ties, tie back long/curly hair

3. Button up loose cuffs on shirts

4. Keep work bench tidy and dry, ensuring no debris or paper in the working area

5. Clean up spills immediately with plenty of water

6. Keep loose coats and scarves away from work area

7. Open windows and keep room well ventilated

8. Wear a laboratory coat to protect clothing

9. If skin comes into contact with liquid wash immediately with plenty of water.

10. Have a fire extinguisher and a fire blanket present in the room in-case of fire

11. Ensure that the room is not too crowded

12. Keep bags well under the table to avoid tripping

13. Always have teachers supervision in-case of an extreme emergency

14. No eating or drinking in the laboratory

15. Iodine, hydrogen peroxide, sodium thiosulphate and potassium iodide are all irritants, and could irritate the skin. If they come into contact with the skin, wash with plenty of water.

16. Hydrogen peroxide is also a bleaching agent, so wash with plenty of water or sodium thiosulphate if spilt on clothes.

The hazards of each chemical are identified below (using HAZCARDS): 27 28

Sulphuric Acid 1.0 mol dm–3:

This can cause severe burns, and solutions between 0.5 mol dm-3 and 1.5 mol dm-3 should be labelled as corrosive. Mixing concentrated sulphuric acid to water is extremely dangerous when mixed with water, and dangerous reactions have been known to occur. Therefore, the

27 http://www.cleapss.org.uk/attachments/article/0/SSS35.pdf?Secondary/Science/Student%20Safety%20Sheets/

28 http://chemistry.slss.ie/resources/downloads/ph_sd_md_iron2chloride.pdf

http://www.cleapss.org.uk/attachments/article/0/SSS35.pdf?Secondary/Science/Student%20Safety%20Sheets/



concentrated sulphuric acid must always be slowly added to cold water, and never the reverse.

If swallowed: Wash out mouth and give a glass of water. Do not induce vomiting. Seek medical attention as soon as possible

If splashed in eye:

Flood the eye with gently running tap water for 10 minutes. Seek medical attention.

If spilt on skin or clothes:

Remove contaminated clothing and quickly wipe as much liquid as possible off the skin with a dry cloth before drenching the area with a large excess of water. If a large area is affected or blistering occurs, seek medical attention.

If spilt in laboratory:

Wear eye protection and gloves. Cover with mineral absorbent and scoop it up into a bucket. Add anhydrous sodium carbonate over the mixture and leave to react. Add lots of cold water. Rinse the area of the spill several times with water.

Hydrogen peroxide 0.10 mol dm–3:

If swallowed: Wash out mouth and give a glass of water. Seek medical attention as soon as possible

If liquid gets in eyes:



Flood affected area with water immediately. Seek medical attention if blistering occurs.


Wear eye protection and gloves. Cover with mineral absorbent and clear up into a bucket. Rinse several times. Add water to dilute at least ten times before washing the liquid down the foul-water drain.

Sodium thiosulphate 0.0050 mol dm–3 and potassium iodide 0.10 mol dm –3:

If swallowed: Give plenty of water. Seek medical attention as soon as possible




Flood affected area with water immediately. Seek medical attention if blistering occurs. Wash off skin with plenty of water.




Iron (II) chloride 1.0 mol dm–3

If swallowed: Give plenty of water. Seek medical attention as soon as possible


Immediately wash eyes with a copious amount of water for at least 20 minutes, then seek medical attention


Flood affected area with water immediately. May cause burns. Wash off skin with plenty of water. Remove any contaminated clothing



BIBLIOGRAPHY

1. Chemistry Individual Investigation (F336) - Starter Page 1

2. http://www.techknow.org.uk/wiki/index.php?title=Iodine_clock3. Salter Advanced chemistry module analysis TL 4.3 pg 54-554. http://revisionworld.co.uk/a2-level-level-revision/chemistry/physical-inorganic-chemistry/

kinetics/collision-theory5. http://en.wikipedia.org/wiki/Hydrogen_peroxide6. http://en.wikipedia.org/wiki/Potassium_iodide7. http://en.wikipedia.org/wiki/Sodium_thiosulfate8. Chemical Ideas 10.3 pg.2149. http://www.chemguide.co.uk/physical/basicrates/arrhenius.html10. http://www.pearsonschoolsandfecolleges.co.uk/AssetsLibrary/SECTORS/Secondary/PDFs/

Science/EdexcelScience/ALevelRevisionGuides/EdexcelA2ChemistryRG_9781846905964_pg8-17_web.pdf

11. http://www.lpscience.com/classes/apchemistry/baugher/printnotes/Reaction %20Mechanism.pdf

12. http://www.chemguide.co.uk/physical/catalysis/introduction.html

13. http://www.pearsonschoolsandfecolleges.co.uk/AssetsLibrary/SECTORS/Secondary/PDFs/ Science/EdexcelScience/ALevelRevisionGuides/EdexcelA2ChemistryRG_9781846905964_pg8-17_web.pdf




http://www.chemguide.co.uk/physical/catalysis/introduction.html







14. http://www.webchem.net/notes/how_far/kinetics/maxwell_boltzmann.htm

15. November 1996 issue of Chemistry Review, Volume 6, Number 2, Pages 14.16. November 1996 issue of Chemistry Review, Volume 6, Number 2, Pages 15.17. http://www.slideshare.net/wkkok1957/iodine-clock-reaction-experiment-using-potassium-

iodide-and-hydrogen-peroxide18. Salter Advanced chemistry module analysis TL 4.2 pg 52-5319. Chemistry Individual Investigation (F336) - Starter Page 1

20. Salter Advanced chemistry module analysis TL 4.2 pg 52-5321. Chemistry Individual Investigation (F336) - Starter Page 1

22. Chemistry Individual Investigation (F336) - Starter Page 123. Salter Advanced chemistry module analysis TL 4.2 pg 52-5324. Salter Advanced chemistry module analysis TL 4.2 pg 52-5325. Chemistry Individual Investigation (F336) - Starter Page 1

26. http://www.ehow.com/how_7805881_activation-energy.html#ixzz2BRYMYgZN

27. http://www.cleapss.org.uk/attachments/article/0/SSS35.pdf?Secondary/Science/Student %20Safety%20Sheets/

28. http://chemistry.slss.ie/resources/downloads/ph_sd_md_iron2chloride.pdf



http://www.ehow.com/how_7805881_activation-energy.html#ixzz2BRYMYgZN



http://www.webchem.net/notes/how_far/kinetics/maxwell_boltzmann.htm

kinetics of the harcourt eseen reaction (steve quarshie) part 1

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