Hydrogen Peroxide Iodine Clock

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Aims The aims of this investigation are: 1. To find the rate equation of the reaction of hydrogen peroxide and iodide ions. This will be achieved by using an iodine clock method and colorimetric analysis. 2. Draw a graph of rate against concentration for each reactant (Hydrogen peroxide, potassium iodide and H + ions). 3. Finding the order for each reactant 4. Finding the rate-determining step. 5. Proposing a mechanism for the reaction. 6. Using Arrhenius’ equation to find the activation enthalpy. Background The basic reaction for this can be illustrated with the following equation: 3I - (aq) + H 2 O 2(aq) + 2H + (aq) → I 3 - (aq) + 2H 2 O (aq) (1) 1 The half equations for this reaction can be written as follows: 3I - I 3 - + 2e - H 2 O 2 + 2H + + 2e - 2H 2 O This reaction demonstrates that iodide ions are oxidised by hydrogen peroxide to tri-iodide ions. This is stage one of a sequence of reactions, which continues below: I 3 - (aq) + 2S 2 O 3 2- (aq) 3I - (aq) + S 4 O 6 - (aq) (2) 2 1 http://antoine.frostburg.edu/chem/senese/101/kinetics/faq/mechanism- h2o2-iodide.shtml 2 http://antoine.frostburg.edu/chem/senese/101/kinetics/faq/mechanism- h2o2-iodide.shtml Planning 1

Transcript of Hydrogen Peroxide Iodine Clock

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AimsThe aims of this investigation are:1. To find the rate equation of the reaction of hydrogen peroxide and

iodide ions. This will be achieved by using an iodine clock method and colorimetric analysis.

2. Draw a graph of rate against concentration for each reactant (Hydrogen peroxide, potassium iodide and H+ ions).

3. Finding the order for each reactant 4. Finding the rate-determining step.5. Proposing a mechanism for the reaction. 6. Using Arrhenius’ equation to find the activation enthalpy.

BackgroundThe basic reaction for this can be illustrated with the following equation:

3I-(aq) + H2O2(aq) + 2H+

(aq) → I3-(aq) + 2H2O(aq) (1)1

The half equations for this reaction can be written as follows:

3I- I3- + 2e-

H2O2 + 2H+ + 2e- 2H2O

This reaction demonstrates that iodide ions are oxidised by hydrogen peroxide to tri-iodide ions.

This is stage one of a sequence of reactions, which continues below:

I3-(aq) + 2S2O3

2-(aq) 3I-

(aq) + S4O6-(aq) (2) 2

This shows that the tri-iodide ions are reduced back to iodide ions by the thiosulphate ions. Thus, the iodine that is formed in reaction (1) is immediately transformed into iodide ion and we do not see the blue-black colour of the starch-iodide complex until all of the thiosulphate ion has reacted with I 2(aq) and is exhausted.

I3-

(aq) + starch Starch-I5- complex + I-

(aq) 3

Once the thiosulphate ion has been exhausted, the tri-iodide ion can react with the starch, forming the Starch-I5

- complex, giving the blue-

1 http://antoine.frostburg.edu/chem/senese/101/kinetics/faq/mechanism-h2o2-iodide.shtml2 http://antoine.frostburg.edu/chem/senese/101/kinetics/faq/mechanism-h2o2-iodide.shtml3 http://antoine.frostburg.edu/chem/senese/101/kinetics/faq/mechanism-h2o2-iodide.shtml

Planning 1

Hasnain Tejani, 03/01/-1,
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black colour. When this occurs, we will then know the amount of hydrogen peroxide that has reacted and the time it took to react.

These equations will, thus, enable the slow step (rate-determining step) to be determined, which is another aim of this experiment.

Though details of the starch and iodine reaction are not yet fully known, it is thought that iodine fits inside the coils of amylose. The transfer of charge between the iodine and the starch and the spacing between the energy levels in the complex formed corresponds to the absorption spectrum, and so, the complementary colour, a blue-black solution, is observed. 4

Factors that affect the rates of reaction5,6

There are many factors that affect the rate of a reaction. These include surface area, concentration difference, presence of a catalyst, pressure and temperature.

Affect of surface area on the rate of reactionSurface area can also affect the rate of reaction. A reaction will happen more quicker if the solid is finely divided into a powder, rather than a lump of the same mass. This is because a reaction can only occur if the particles taking part in the reaction collide. A larger surface area provides a higher likelihood of collisions (and thus, a reaction) to take place. One example is called the “Bread and Butter Theory.” 7 If you take a loaf of bread and cut it into slices it, you have more surfaces to spread butter onto. Taking a more practical example,

4 en.wikipedia.org/wiki/Starch5 www.s-cool.co.uk6 http://www.chemguide.co.uk/physical/basicrates/orders.html, Jim Clark, 20027 www.purchon.com/chemistry/rates.htm

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Reactant 1

Reactant 2

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In the above example, reactant 1 can get to the outer atoms of reactant 2, but not the central atoms. This has a small surface area. However, if the surface area is increased,

Reactant 1 can get to all the atoms of reactant 2.Increasing the number of collisions per second increases the rate of reaction.

Affect of concentration on the rate of reactionFor many reactions, including this one, increasing the concentration of the reactants increases the rate of the reaction. This is because, for a reaction to occur, a collision must take place first. Increasing the concentration of the reactants will increase the frequency of the collisions between two reactants, as there are a higher number of reactants to collide with. From a probabilistic point of view, if there are a higher number of reactants (i.e. a higher concentration), the chance of a collision, and therefore, a reaction to take place, increases. For example, if we have the following situation:

supposing fixed positions and an equal probability of being hit, the probability of a green particle hitting a red particle is 1/3. If we increase the number of red particles to 2, the probability now of a green particle hitting a red particle is ½, which is thus, an increase by 1/6.

Although the temperature is being kept constant, however, the kinetic theory is applicable. This is because the molecules involved in the reaction have a range of energy levels. When colliding molecules have sufficient energy, a reaction takes place. If they do not, then a reaction cannot take place. This is because the temperature that is being measured is only the average temperature (and thus, kinetic energy, because T α Ek) of the substance. It is impossible for the kinetic energy of every atom in the substance to be the same, and so, the temperature is an average. The reason for this is that each

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molecule has a certain amount of kinetic energy and once it collides (perfectly elastically) with another molecule, it transfers its kinetic energy to the molecule it collided with, giving it a higher kinetic energy than the initial molecule. In most cases, when you increase the concentration, the rate of reaction also increases. In certain multi-step reactions, however, the reaction happens in a series of small steps. Suppose the reaction happens as so:

The speed at which A splits into X and Y dictates the rate of the reaction. This is also known as the rate-determining step. If you increase the concentration of A, the chances of the first step happening also increase, due to the increase in the number of molecules of A. Increasing the concentration of B undoubtedly speeds up the second step, but makes little difference to the overall rate.

Affect of the presence of a catalyst on the rate of reactionA catalyst is a substance that speeds up a reaction by providing an alternative pathway with a lower activation enthalpy, and is chemically unchanged at the end of a reaction. Reactions can only take place if the two reactants collide with enough energy to initiate the reaction (i.e., to begin breaking the bonds). Majority of the molecules do not have enough energy, and simply bounce apart after collisions. One way of speeding up a reaction is to provide an alternative pathway for the reaction to occur with a lower activation energy. In other words, the activation energy on the Maxwell-Boltzmann Distribution graph should look like this:

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A catalyst does this, and this can be shown on an enthalpy profile diagram:

A catalyst works in two ways. One of them is adsorption, and this is where the molecules are attached to the surface of the catalyst due to the weak interactions (typically Van Der Waal’s forces) between the surface and the reactants. Initially, the bonds in the reactants weaken

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Now all these particles have enough energy to react as well

Diagram obtained from http://cwx.prenhall.com/bookbind/pubbooks/hillchem3/medialib/media_portfolio/13.html

ΔH

Uncatalysed activation enthalpy

Catalysed activation enthalpy

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and break. Bonds are then formed between the reactants, forming the products, and they then diffuse away from the surface of the catalyst. Another way is the formation of the intermediate compounds, and many catalysts, including all enzymes, work by forming intermediates. The reactants involved in the reaction combine with the catalyst making an intermediate compound, but this is very unstable. When this intermediate breaks down, it releases the new compound and the original catalyst.

Affect of the pressure on the rate of reactionIncreasing the pressure on a reaction involving gases increases the rate of reaction. This does not happen with reactions involving solids or liquids. Increasing the pressure works in the same way as increasing the concentration; if the pressure of a given mass is increased, it is just squashed into a smaller volume. Having the same number of particles in a smaller volume works in the same way as increasing the concentration. The ideal gas equation illustrates this (As liquids act similarly to gases, the ideal gas equation can give a fair demonstration of how liquids would act):

pV = nRT

Where p = pressure,V = volumen = number of molesR = molar gas constantT = Temperature (in Kelvin)

This can be re-arranged to read:

p = n/v x RT

n/v is the number of moles divided by the volume, which is the concentration.RT is constant at a constant temperature. Because this is just a constant, it can be shown that p = k(n/v)or p α (n/v)(the pressure is proportional to the concentration). Thus, if you double one, the other will also be doubled.

Whether you are considering a reaction where collisions between two different particles or two of the same particles occur, the same law applies: for any reaction to occur, collisions must happen first. This is true when both particles or one of the two sets of particles are in the

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gaseous state. If the pressure is higher, the chance for collisions to occur also increases.

If the reaction involves a particle splitting, the reacting particles must have enough energy to react. Supposing that one in a hundred particles have enough energy to react - if you had two hundred particles, two of them would react, and so forth, so if you double the pressure, the rate of reaction also doubles.

The main variable to be tested in this experiment will be the concentration of hydrogen peroxide and the concentration of iodide ions. These will be varied in two separate experiments, thus enabling a fair test to be attained. If both variables are altered at once in an experiment, it would be very difficult to say which variable has had more effect on the rate of reaction. Therefore, in the first investigation the concentration of hydrogen peroxide will be varied, keeping all other variables constant and in the second investigation the concentration of iodide ions will be altered, keeping all the other variables constant. The results of this investigation will enable me to draw a rate graph for the two investigations. This will later enable me to combine these two rates to form an overall rate equation.

The iodine clock reaction is a chemical reaction in which two colourless solutions are mixed and react together to form a red/brown 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 be very hard to detect. This is, thus, enhanced by the addition of starch, which instantaneously turns dark blue/black with the formation of iodine ions, giving a more accurate time for the production of iodine ions.

A colorimeter is an instrument that measures the concentration of a substance by comparing its colour with the standard (i.e. distilled water). In this experiment, the complementary colour to the orange of

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Low pressure High pressure

Diagram referenced to www.chemguide.co.uk

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iodine is blue-green, so a blue-green filter (470 nm) will be used in method two on the colorimeter.

This 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 accelerates the first reaction (e.g., iron catalysis or temperature) will shorten the time. Thus, increasing the concentration of iodide, hydrogen peroxide, or acid (it neutralises the

hydroxide ion) will accelerate the reaction. On the other hand, increasing the thiosulphate concentration will have the opposite effect; it will take longer for the blue colour to appear. The rate of reaction is a measure of how fast the reaction occurs. The graph that can be drawn from the results is time against volume (concentration) of the variable solution. A rate of reaction can then be obtained. Average rates are not very good comparisons, because the reaction may finish before the designated time interval. The fairest way to measure the rate is at the start, as a fair test is attained. Thus, in this experiment, the initial rate of the reaction will be measured. This is to allow the reaction to progress to about 10-15%, which enables a fair comparison of the reactions to take place. To measure the initial rate, a number of readings will take place within a fairly short space of time. Drawing a tangent to the curve where it crosses the origin will help to measure the initial rate of the reaction. This is indicated below:dy/dx then gives the gradient of the curve at that particular point, or the initial rate.

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dy

dx

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All the initial rates will be put onto a graph and the initial rate will be plotted against the concentration. This will give the order of rate for one reactant. When the rate orders for all the reactants are found, a rate equation can be formed.

It will generally look like this:Rate = k.[H2O2]a[I-]b[H+]c

The quantities in brackets are concentrations in moles of each reactant and k is the rate constant for the reaction. This stays the same throughout the experiment, no matter what the concentration of the reactants are. This rate constant is temperature dependent. The quantities a, b, and c are called the reaction order for H2O2, I- and H+; they will be determined in the experiment.

To accomplish this, two principles from kinetic studies will be applied. The first principle is to hold two of the reactants constant while varying only one component. If the concentrations of I- and H+ are held constant we may write the rate as:Rate = d[I-]/dt = k.[H2O2]a = [S2O3

2-]/2Δt

Where:● Δt = time to reach an observable blue colour● [I-]b and [H+]c have been absorbed into the constant k.

If rate = k.[H2O2]a,

Ln(Rate) = ln(k) + ln([H2O2])a

= Ln(Rate) = ln(k) + a.ln([H2O2]).

By plotting the ln(Rate) versus ln([H2O2]), a linear relationship will be formed:

Comparing this with y = mx + c enables us to identify that the result will be a graph with a gradient of a, and a y-intercept of ln(k). This enables us to find a. Thus, a, b and c can all be found by looking at their reactions.

The second principle is called “The Method of Initial Rates”. In this experiment, the concentration of thiosulphate is much smaller than the other reactants. The end point colour appears after all of the thiosulphate is used up, allowing I3

- to react with the starch to form the blue complex (reactions 2 and 3). The amounts of reactants used up in causing this to take place are small, so the reactant concentrations remain essentially constant throughout the time of reaction. Also, the amount of reactants used up at the time of the

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endpoint is a constant because the amount (moles) of thiosulphate present is constant at the beginning of all the reactions.

Affect of temperature on the rate of reactionThe collision theory states that when two chemicals react, their molecules must have sufficient energy during the collisions for a reaction to take place. The two molecules will only react with each other if there is enough activation energy for the reaction to begin. Activation energy is required to initiate the breaking of the bonds in the molecules. This requires energy, and so, it follows that, if there is not enough energy to break the bonds, no new bonds will be able to form, and thus, no reactions take place. If the two molecules collide with each other with less than the required activation energy, they will just bounce apart elastically. By heating the mixture, the amount of kinetic energy being given to the molecules is raised. The kinetic theory then states that increasing the kinetic (heat) energy of a molecule means that the molecules begin to move faster. The faster they move, the more chance there is of a collision, and thus, a reaction to take place. Due to the importance of activation energy in reactions, it is very useful to know which proportions of the particles present have enough energy to react. In any system, the particles will have a wide range of energies. For gases, the Maxwell-Boltzmann Distribution can represent this, which is basically a plot of the number of particles against energy.

The area under the graph represents the number of molecules present. Some interesting points to note from this distribution curve

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Obtained from www.webchem.net

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are that the kinetic energy can never be zero, there are only a few molecules with a high energy, and there is no defined maximum energy value (and so, it can theoretically continue to infinity).8

If we take the example of a single particle inside a box,9

If we imagine that the particle moves from one wall to the other, and then rebounds back (perfectly elastically), the momentum of the particle intially moving from the right to the left would be p = mu

Using vector quantities, once the particle bounces off the wall and heads in the opposite direction, the momentum is p = - mu

The change in momentum is therefore, Δp = mu - (-mu) = 2mu

Speed = distance/time taken= 2l/u

Thus, the rate of change of momentum = 2mu/(2l/u) = mu2/l

Newton's second law of motion states that force = rate of change of

momentum, and so the force on particle = m (u12 + u2

2 +… + uN2)

Since pressure = force per unit area, we can say that pressure, p, exterted on a wall of area l2

is given by:

p = m (u12 + u2

2 +… + uN2)

If <u2> represents the mean value of the squares of all the velocities, <u2> = (u1

2 + u22 +… + uN

2)/N

8 Chemistry: AS level and A level (International), Ratcliffe, Eccles, Raffan, Nicholson and Johnson, 20049 Oxford: Succes at AQA Physics B A2, Ken Price and Gerard Kelly, 2001

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l

l3

l

u

v

c

w

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and N<u2> = u1

2 + u22 +… + uN

2

Therefore, p = Nm<u2>/l3

For any molecule, pythagoras’ theorem can be applied to give:

c2 = u2 + v2 + w2

This will also be true for all the mean square values,

<c2> = <u2> + <v2> + <w2>

However, since N is large and the particles move randomly, it follows that the mean values for u2, v2 and w2 are equal.

Thus, <c2> = 3<u2>

Therefore, <c2>/3 = <u2>

Hence, p = Nm<c2>/3l3

l3 = volume of the cube = V, so

pV= 1/3 (Nm<c2>)

N = L = Avogadro’s constant, giving:

pV = 1/3 (Lm<c2>)

This can be written as:

pV = 2/3 L(½m<c2>)

The ideal gas equation for a mole is pV = RT,

Where R is the molar gas constant. Combining the last two equations gives:

2/3 L(½m<c2>) = RT

Therefore, (½m<c2>) = 3/2 (R/L)T

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Now ½m<c2> is the average kinetic energy for each molecule, and from the above equation, we know R and L are constants. The ration R/L is called Boltzmann’s constant = k = 1.38 x 10-23 J K-1

Hence, we can say:

½ m<c2> = 3/2 kT. This will be particularly useful when working out the kinetic energy of each particle in the analysis, and for the extension task.

Transition Theory10

Another factor that could affect a reaction taking place is summarised in the Transition Theory. This states that a collision between reactant molecules may or may not yield a product. The factors that decide if a reaction takes place or not are the relative kinetic energies of the molecules involved, the orientation, and the internal energy of the molecules. If the reactants form an activated complex (an unstable grouping of atoms that break apart to form products) they are not guaranteed to go on and form products - they could just fall apart back into reactants.

Arrhenius’ equationAs mentioned earlier, k is temperature dependent. The rate constant is related to temperature by the Arrhenius equation, which predicts the rate of a chemical reaction at a certain temperature, given the activation energy and chance of successful collision of molecules:

k = A.e-Ea/RT

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

Ln(k) = ln(A) – Ea/RT,

Where Ea = the activation energy

Comparing this to y =mx + c,

Y = ln(k)

10 http://www.kobold.demon.co.uk/kinetics/collisio.htm

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c = ln(A)x = 1/Tm = -Ea/R

Therefore, the activation enthalpy can be found by:

Gradient = -Ea/R

-Gradient x R = Ea

How can the rate be calculated?Measuring volumeThe volume of product produced could be measured, either using an inverted burette, or a gas syringe. An example of where this is applicable is in the reaction of hydrogen peroxide and catalase. The oxygen evolved from the boiling tube passes through the rubber tubing, and displaces water in an inverted burette, or displaces the stopper in a gas syringe. This enables the volume of product formed to be measured. A graph can then be drawn for each, and the initial rate of reaction would be measured as the gradient of volume of product produced against time graph. The initial rates for each value would then be plotted against concentration, to give a graph to find the order of a reaction.

Measuring pressureWhen two substances react with each other, and the product is a gas, the effect of varying pressure can be detected. Allowing a reaction to take place under normal room pressure and comparing this with a reaction under compression, such as placing a bung on a test tube, allows one to plot a graph of pressure (which can be calculated as nRT/V = p) against time. The gradient at the start of the experiment of this will give the initial rate of the reaction. These can then be plotted on a graph to give the overall order of a reaction.

Titrimetric analysisThis is a standard method of chemical analysis, and it can be used to determine the concentration of a known reactant. Using a burette, it is possible to determine the exact amount of a reactant before the endpoint is reached. With a reaction, small portions must be taken out at suitable time intervals, and the reaction quenched (halted). One way of quenching is to put the samples into a conical flask, immersed into a ice/salt mixture. This allows the concentration of the substance

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to be measured by titration. The measurements made are the concentrations of the reactant at the moment the sample was taken out. This can then be plotted on a graph of concentration against time, and the gradient at the start of the graph will give the initial rate of reaction, which can be determined by drawing on a tangent. The collection of these rates plotted together against concentration will give the overall order of the reaction.

Colorimetric analysis11

The size of the filter chosen for the colorimeter is extremely important, as the wavelength of light that is transmitted by the colorimeter has to be same as that absorbed by the substance. This is the complementary colour for the colour of the substance. The percentage absorbency can be measured at set intervals throughout the experiment. The percentage absorbency can then be plotted against time (with units of mol s-1). A calibration curve can also be plotted, with absorbance being plotted against time.

Coductimetric analysisThis is the measurement of the conductance of a solution. Ions conduct electricity, and as they will be present while the reaction is occurring, a measure of the electrical conductivity could be used to measure the rate of the reaction. For this, the conductivity of the substance would be plotted against the respective concentration, and the initial part of the curve gives the initial rate of reaction.

Clock methodThis works by reacting small amounts of one reactant with another, and converting this to another intermediate substance. An indicator is added to enhance the production of the final product. An additional substance is also added which blocks the reactant from being produced by means of a chemical reaction, and so, until the additional substance is used up, the reactant will not react with the indicator to form a coloured complex. For example, if a small known volume of sodium thiosulphate is added to the reaction producing iodine, until all the thiosulphate ions have been consumed, the iodine will not build up.

Rate ordersThe rate orders that I can expect for my results are one of three; Zero order, first order or second order.

11 Facts and Practice for A-level: Chemistry, Max Parsonage, 2001

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Zero order

Rate laws are differential equations because the rate of a reaction is the rate of change of a reaction with time. Most reactions are of either first or second order. Rate laws of zero order are not common. Except for zero-order rate laws for which the rate is independent of concentration, the rate of a reaction will change as the reaction proceeds because the concentrations of reactants and products change as the reaction proceeds.

The following graph is the graph of a zero order reaction, and it shows that the rate of reaction with zero order remains constant throughout the reaction. The gradient is constant, producing a straight line. It can be seen that the half-life is always decreasing with the decreasing

concentration.A zero order graph indicates that the rate of reaction is not affected by the reactant. If a reactant is zero order then it is not included in the rate equation as its not affecting the rate of the reaction.

Planning 16

Initial rate of Reaction

Concentration of reactant

Half-life decreases with decreasing concentration

Graph obtained from www.chem.purdue.edu

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First OrderThe following graph is the graph of a first order reaction, and it shows how the rate of reaction with first order varies throughout the reaction. The gradient is constantly changing, producing a curve (like the left-hand side of an x2 graph). It can be seen that the half-life is always decreasing with the decreasing concentration.

A first order graph indicates that the rate of reaction is affected by the reactant. If a reactant is first order then it is included in the rate equation as it is directly affecting the rate of the reaction.

Planning 17

Initial rate of Reaction

Concentration of reactant

Half-life decreases with decreasing concentration

Graph obtained from wps.prenhall.com

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As the concentration increases, the rate of reaction also increases. The evidence that shows that this graph is of first order is the direct proportionality between the concentration of the reactant and the initial rate of reaction, as it has a straight line and passes through the origin.

Second order

The following graph is the graph of a second order reaction, and it shows how the rate of reaction with second order varies throughout the reaction. The gradient is constantly varying, producing an exponential curve. It can be seen that the half-life is always decreasing with the decreasing concentration.

Planning 18

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It is apparent upon looking at this graph that the initial rate does not have a linear relationship with the concentration of the reactant. It has an exponential relationship. This means that if the concentration doubles, the rate quadruples. It can be concluded that the initial rate is directly proportional to the concentration of the reactant squared (it has a squared relationship), i.e. Rate α [x2]

This graph is a second order graph.

Background of Hydrogen Peroxide12,13

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.

12 Hydrogen Peroxide, in Kirk-Othmer Encyclopaedia of Chemical Technology13 http://en.wikipedia.org/wiki/Hydrogen_peroxide

Planning 19

Initial rate of Reaction

Concentration of reactant

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Hydrogen peroxide often decomposes exothermically in the presence of light, and so, it needs to be stored in a cool environment in a brown bottle. This will mean that, in the experiment, the hydrogen peroxide solution will need to be freshly made up everyday.

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 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 everyday, and so, there is little chance for it to decompose.

Background of Potassium Iodide14

Potassium iodide is a white, crystalline salt, with a formula of KI. It is often used as an iodide source, because it is less hygroscopic (attracting and retaining water) than sodium iodide, which enables it to be easier to work with. 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.

14 Dictionary of Inorganic Compounds, J.E. Macintyre, 1992

Planning 20

Structural formula obtained from: http://en.wikipedia.org/wiki/Hydrogen_peroxide

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The structural formula for potassium iodide is shown below:

.

Background of Sodium Thiosulphate

Sodium thiosulphate is a colourless crystalline compound, and is more commonly found in its pentahydrate state, Na2S2O3.5H2O. In this experiment, the sodium thiosulphate is in the pentahydrate state. It is often used in photography as a fixer of film, and in the tanning of leather in chemical manufacture.

The thiosulphate anion reacts with iodine, reducing it to iodide as it is oxidised tetrathionate: I3

-(aq) + 2S2O3

2-(aq) 3I-

(aq) + S4O6-(aq).

The structural formula for sodium thiosulphate is as follows:

Planning 21

Structural formula obtained from: http://www.webelements.com/webelements/compounds/text/K/I1K1-7681110.html

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ExperimentVariables

When a set of results are obtained, it is very difficult to determine exactly which variable has the greatest effect on the result. Therefore, there can be no more than one variable. The presence of only one variable is called a “fair test.”

The variables that will need to be kept the same in every reading for this experiment are:

Use of the same volume of starch solution:In each experiment, 1.0cm3 of starch will be used. It is essential that the same volume of starch is present before each experiment. The reason for this is that for a higher volume of starch, there is a higher concentration, and so, the time taken for the blue-black colour to appear would lessen by a few milliseconds each time. 1.0 cm3 is a reasonable volume of starch to use, so it does not affect the experiment too much. If there was a lower volume of starch solution, a small inaccuracy in the reading of the starch solution could have a negligible effect on the experiment, and thus, vary the results of the experiment. This enables a fair test to be attained each time.

Use of the same volume of potassium iodide solution:In each experiment, 2.0cm3 of potassium iodide will be used. It is essential that the same volume of potassium iodide is present before each experiment. The reason for this is that the potassium iodide is one of the factors that effects the rate of reaction, and the

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Structural formula obtained from:http://msds.pcd.go.th/images/Formula_Chain/10102-17-7.gif

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order of the rate, and so, it must be kept constant if the affect of varying hydrogen peroxide is to be measured. 2.0 cm3 is a reasonable volume of potassium iodide to use, as it can be measured quite accurately with a graduated pipette, with a relatively small uncertainty. If the volume of potassium iodide were varied, then it would be very hard to measure the effect of varying hydrogen peroxide on the rate of reaction. This enables a fair test to be attained each time.

Use of the same volume of sulphuric acid:In each experiment, 4.0cm3 of 0.300 mol dm-3 sulphuric acid will be used. It is essential that the same volume of sulphuric acid is present before each experiment. The reason for this is that the H+ ions have an effect on the reaction and the order of the rate, and so, must be kept the same of the affect of varying hydrogen peroxide concentration is to be measured. Also, hydrogen peroxide is less likely to decompose in acidic conditions. 4.0 cm3 is a reasonable volume of buffer solution to use, as it can be measured quite accurately with a graduated pipette, with a relatively small uncertainty. This enables a fair test to be attained each time.

Use of universal indicator paperIn each experiment, the pH of each reaction will be measured using universal indicator paper. It is essential that the pH is kept the same each time. The reason for this is that varying the pH could cause the results to become slanted and flawed, and so, to stop the results obtained from the experiment to be nullified, the pH must be fixed. This enables a fair test to be attained each time.

Use of a balance:To make up each solution, a balance will be required to help measure out the weight of the solid chemical, and then made into a liquid. It is important to keep the same balance each time, because if there is a small calibration problem, then this will be carried all the way through the experiment, and so, the values obtained will be relative. This would mean a fair test is attained each time.

Use of the same volume of sodium thiosulphate solution:In each experiment, 2.0cm3 of sodium thiosulphate will be used. It is essential that the same volume of sodium thiosulphate is present before each experiment. The reason for this is that the sodium thiosulphate is one of the factors that affects the rate of this reaction, as it destroys the iodide ions. 2.0 cm3 is a reasonable volume of sodium thiosulphate to use, as it can be measured quite accurately with a graduated pipette, with a relatively small uncertainty. If the volume of sodium thiosulphate were varied, then

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it would be very hard to measure the effect of varying hydrogen peroxide on the rate of reaction. This is because the time measured each time would vary, as the amount of iodine destroyed would vary each time. Keeping the same volume of sodium thiosulphate each time enables a fair test to be attained.

Keeping the total volume constant:In each experiment, the total volume will be 14.0cm3. It is essential that the total volume should stay the same for each experiment. The reason for this is that if the total volume changes, the concentration also changes and the results obtained from the experiment vary. If the volume is kept the same, when you double the volume, the concentration will also double. This enables a fair test to be attained each time.

Use of the same colorimeter:The colorimeter in every experiment will be the same, and the same blue-green filter (470 nm) will be used. This is due to the fact that different colorimeters may be calibrated differently at the place of manufacturing, and therefore, give slightly different absorption readings. This would hinder the final outcome of the results, as there would be fluctuations in the results obtained. The colorimeter will also have to be calibrated back to zero absorption each time with distilled water to keep a fair test, as a slight change in calibration can hinder the final results.

Cleaning the equipment with distilled water:The equipment must be cleaned after each experiment to remove all the chemicals from the equipment. This ensures a fair test and makes sure that the results of the experiment are not flawed.

Use of the same graduated pipette each time:If there is some error in the graduated pipette that makes it unable to read the volume accurately, then, as it will be used in every experiment, the errors will be relative to one another and it will not effect the final outcome greatly. Furthermore, due to human error, if the volume is not measured accurately, then as these errors will be continued in every experiment, the errors will be relative to one another and it will not effect the final outcome greatly. This ensures a fair test.

Use of the same thermometerIf there is some error in the thermometer that makes it unable to read the temperature accurately, then, as it will be used in every experiment, the errors will be relative to one another and it will not effect the final outcome greatly. This ensures a fair test.

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Use of a volumetric flaskIt is important to use a volumetric flask to measure solutions, because the calibration line accurately shows the volume of a liquid at room temperature. This ensures that the concentration of each solution is precise each time, meaning a fair test.

Measuring below the meniscus:I will measure from below the meniscus to provide the most accurate reading possible. This is because the bottom of the meniscus provides the actual reading of the volume inside the measuring cylinder. This ensures a fair test each time.

Use of a water bath:A water bath will be used to make sure that the temperature is kept constant at 25oC each time. As temperature is also a factor that affects the rate of a reaction, keeping this constant each time will allow the effect of changing concentration to be measured accurately. The starch solution, hydrogen peroxide, sodium thiosulphate, sulphuric acid and water will all be placed in a boiling tube, which will then be kept in a water bath, before the potassium iodide is added. Just to be doubly sure that the water bath is at the stated temperature, the mercury thermometer will be immersed into the water bath, and the temperature checked. This enables the reaction to take place at the same temperature each time, ensuring a constant temperature, and thus, enabling a fair test.

The same stopwatch each time:If there is some error in the calibration of the stopwatch that makes it unable to read the time accurately, then, as it will be used in every experiment, the errors will be relative to one another and it will not effect the final outcome greatly. Furthermore, due to human error, if the time is not measured accurately, then as these errors will be carried through in every experiment, the errors will be relative to one another and it will not effect the final outcome greatly. This ensures a fair test.

Use of a volumetric flask

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It is bad practice to put solids into a volumetric flask. Therefore, before making up a solution, the solids must be added to a beaker, and about half the required distilled water added. The solution should then be stirred with a stirring rod. This solution should then be transferred to a volumetric flask using a glass funnel. The glass funnel and stirring rod should be washed repeatedly, and these “washings” added to the volumetric flask. Distilled water should then be added until the solution is 1.00 cm below the calibration line. The distilled water should then be added drop by drop until the bottom of the

meniscus is just touching the calibration line. A stopper should then be placed on top of the volumetric flask, and the volumetric flask inverted a few times.

It is important to use a volumetric flask to measure solutions, because the calibration line accurately shows the volume at room temperature. Also, to keep precision, all utensils used in the makeup of the solution should be washed repeatedly, and these “washings” added to the solution. This ensures a high level of precision, and a fair test overall.

Makeup of solutionsFor a solid,

Mol = mass/Mr,

Where Mol is the number of moles of the substance, the mass is measured in grams using a balance, and the Mr is the relative molecular mass of the substance

For a liquid,

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This is the meniscus. You must always measure volumes from below the meniscus.

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Mol = volume x concentration.

Equating these two equations gives

Mass/Mr = volume x concentration.

Therefore, mass required = volume required x concentration required x Mr

Mass of sodium thiosulphate required

The formula for sodium thiosulphate in its pentahyrdrate state is Na2S2O3.5H2O. The Mr of this is 248. In this experiment, the volume of sodium thiosulphate required is 100 cm3 = 0.100 dm3 and the concentration required is 0.00500 mol dm-3

Therefore, Mass required = volume required x concentration required x Mr,

= 0.100 x 0.00500 x 248= 0.124g

However, this is a very small value, and so, there is a high likelihood for uncertainties to arise. Therefore, 1.24g of sodium thiosulphate will be added to a beaker, and then about 50.0cm3 of water added. This solution will then be transferred to a 100 cm3 volumetric flask using a glass funnel, and the beaker, glass funnel and stirring rod will then be repeatedly washed using distilled water. These washings will then be added to the volumetric flask, until the solution is made up to 100 cm3. This forms 0.0500 mol dm-3 of sodium thiosulphate solution, so a further ten-fold dilution will be made, by taking 10.0 cm3 out of the 0.0500 mol dm-3 sodium thiosulphate and putting this inside another 100 cm3 volumetric flask. This solution will then be made up to 100cm3 using distilled water.

Mass of potassium iodide required

The formula for potassium iodide is KI. The Mr of this is 166. In this experiment, the volume of potassium iodide required is 250 cm3 = 0.250 dm3 and the concentration required is 0.0200 mol dm-3

Therefore, Mass required = volume required x concentration required x Mr,

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= 0.2 x 0.25 x 166= 0.830g

However, this is a very small value, and so, there is a high likelihood for uncertainties to arise. Therefore, 8.30g of potassium iodide will be added to a beaker, and then about 150 cm3 of water added. This solution will then be transferred to a 250cm3 volumetric flask using a glass funnel, and the beaker, glass funnel and stirring rod will then be repeatedly washed using distilled water. These washings will then be added to the volumetric flask, until the solution is made up to 250 cm3. This forms 0.200 mol dm-3 of potassium iodide solution, so a further ten-fold dilution will be made, by taking 25.0 cm3 out of the 0.200 mol dm-3 potassium iodide solution and putting this inside another 250 cm3 volumetric flask. This solution will then be made up to 250cm3 using distilled water.

Mass of starch granules required

The starch solution can be made up using a standard amount of starch granules. In this case, the standard is 0.33g of starch granules added to 0.250 dm3 of water.

The method to make it up is as follows:

Place starch granules into a beaker, and add about two thirds of the required volume of distilled water. In this case, the volume required is 0.250 dm3 of water, and so, I will add about 80.0 cm3 of water into the volumetric flask. Heat and swirl this over a bunsen burner until it starts to bubble. When it bubbles, add another 100 cm3 of water and remove the bunsen burner from under the beaker. Keep swirling for about a minute, and then place down onto a heatproof mat, leaving it to cool. When it has cooled, the starch should be added to the volumetric flask using a glass funnel. The beaker and glass funnel should then be repeatedly washed, and the washings added to the solution. The distilled water should then be added drop-by-drop until the bottom of the meniscus is just touching the calibration line. It is important to let the starch solution cool, as heat causes the volume of water to rise slightly above the calibration line (as T α V) and the volume would be inaccurate, and thus, the test would not be fair.

The same starch solution cannot be kept throughout the duration of the experiment, as starch can decompose, and so, a fresh solution of starch should be made-up each day. This will also need to be stored in

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a cool place before use, as the starch-iodine complex becomes unstable above 35oC.

Volume of hydrogen peroxide required

As hydrogen peroxide does not come as a solid, but as a liquid, a slightly different approach will be needed to find the right amount of hydrogen peroxide required in this experiment.

In this experiment, the volume of hydrogen peroxide required is 250 cm3 = 0.250 dm3 and the concentration required is 0.0300 mol dm-3

Mols of hydrogen peroxide required = concentration required x volume required

= 0.03 x 0.25= 0.00750 mol

The volume required would then equal the number of moles/concentration given. In this experiment, the concentration of hydrogen peroxide given is 1.67 mol dm-3.

Volume required = mol/given concentration= 0.00750/1.67= 4.50 cm3

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

The hydrogen peroxide will have to be made up each day, as it decomposes in the presence of sunlight. Also, during its use, it will need to be stored in a dark and cool place to minimise decomposition.

Sulphuric acid

In this experiment, 0.300 mol dm-3 of sulphuring acid will be required. The recipe card states that, for small concentration values such as these, a ten-fold dilution will need to take place from 3.00 mol dm-3. As acid is very corrosive, gloves and eye protection (including a face shield if possible) should always be worn. Acid must always be added to water, and so, initially, the 500 cm3 volumetric flask should be half-

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filled with distilled water. Next, 81.0 cm3 of concentrated sulphuric acid will be added to the 500 cm3 volumetric flask, and the solution then made up to 500 cm3. 50cm3 of this 3.00 mol dm-3 sulphuric acid solution will be taken out and placed in another 500 cm3 volumetric flask half-filled with water. This will then be made up to 500cm3 to give sulphuric acid of 0.300 mol dm-3 concentration.

Risk assessmentBefore commencement of any experiment, the potential risks involved must be considered and eliminated in the experiment (or minimised as much as possible).

Ways in which to minimise the potential chance of accidents happening are as follows:

1. Wear goggles to protect eyes

2. Tuck in loose ties

3. Button up loose cuffs on shirts

4. Keep work bench tidy and organised, making sure there are no loose bits of paper on desks

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. Tie hair back in necessary

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

12. Ensure that the room is not too crowded

13. Keep bags well under the table to avoid tripping

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

15. No eating or drinking in the laboratory

16. Do not wear hair gel/mousse when working with an open flame because the compounds contained in these substances may be flammable

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

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

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

Sulphuric Acid: 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 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.

Iodine: It is harmful by inhalation and by contact with the skin. Avoid contact with the eyes. The solid has a corrosive action on the skin, causing burns if left for some time.

If swallowed:

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

If vapour inhaled:

Remove victim to fresh air. Seek medical attention if breathing is slightly affected.

If vapour affects eyes:

Bathe eyes. If discomfort persists, seek medical attention.

If solid gets in eyes:

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

If spilt on skin or clothes:

Brush off solid immediately. Flood affected area with water immediately, or bathe in sodium thiosulphate solution. Seek medical attention if blistering occurs.

If spilt in Wear eye protection and gloves. Ventilate area of spill.

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laboratory: Spread sodium thiosulphate over area of spill, leave for an hour, and then mop up and rinse.

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Hydrogen peroxide:

If swallowed:

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

If liquid gets in eyes:

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

If spilt on skin or clothes:

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

If spilt in laboratory:

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 and potassium iodide:

If swallowed:

Give plenty of water. Seek medical attention as soon as possible

If liquid gets in eyes:

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

If spilt on skin or clothes:

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

If spilt in laboratory:

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.

I have created a guideline as to how to carry out the practical, which I will comply with throughout the experiment. These rules are in place to ensure that the experiment is carried out safely, and to avoid any injuries or contamination. It also tries to ensure that I am safe, and that the people around me are also as safe as possible.

General list of apparatus required throughout the experimentIt is necessary to know what apparatus is being used before doing the experiment. The following apparatus will be used for this experiment:

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A graduated pipette:

This helps measure the volume of solution required to a high degree of accuracy. This will help in trying to make sure that the data collected is reliable by indirectly making sure that the total volume is kept the same each time (i.e. 14cm3).

Universal indicator paper:In each experiment, the pH of each reaction will be measured using universal indicator paper. It is essential that the pH is kept the same each time. The reason for this is that varying the pH could cause the results to become slanted and flawed, and so, to stop the results obtained from the experiment to be nullified, the pH must be fixed. This enables a fair test to be attained each time.

A mercury thermometer:A mercury thermometer is essential to this experiment, as it will determine when the reading should be taken. In the experiment, an analogue mercury thermometer will be used. This will be used because it measures to an accurate degree of accuracy (+ or – 0.5oC). It will make sure that the data that is collected is reliable as mercury thermometers provide an accurate measurement of the temperature. As temperature is also a factor that affects the rate of a reaction, measuring, and therefore keeping, this constant each time will allow the effect of changing concentration to be measured accurately. This enables a fair test.

Test tube rackBoiling tubes can be stored in the test tube rack to prevent accidents

A colorimeterA standard colorimeter on the blue-green filter (470 nm) will be used. This is because it provides an accurate reading of the percentage absorbency. The results obtained will be reliable each time, because, as the colorimeter is the same, the calibration from the manufacturers will also be the same, causing reliable results to be attained.

15 boiling tubesThese will ensure that there is enough space for the reaction to proceed in without the fear of the solutions overflowing the boiling tube. Also, the boiling tube is easy to use and can easily be placed on a test tube rack.

A wash bottle of distilled water

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This will be used to dilute the solutions to the required concentration. They will also be used to clean the equipment each time, resulting in a fair test each time as no contamination will occur.

Water bathA water bath will be used to make sure that the solution is at the stated temperature for the extension task. The starch solution, hydrogen peroxide, sodium thiosulphate and water will all be placed in a boiling tube, which will then be kept in a water bath, before the potassium iodide is added. Just to be doubly sure that the water bath is at the stated temperature, the mercury thermometer will be immersed into the water bath, and the temperature checked. This enables the reaction to take place at the stated temperature each time, thus, enabling a fair test.

5 volumetric flasks:1 x 0.100 dm3 flasks,3 x 0.250 dm3 flask1 x 0.500 dm3 flasksThe solutions will be made up in the volumetric flasks, to ensure accuracy, and to make sure no cross-contamination occurs. If contamination does occur, then the results obtained would be flawed, and thus, unreliable. The use of separate volumetric flasks ensures a fair test.

5 beakersTo avoid cross-contamination, the chemicals will be placed inside a beaker, and not be returned to the volumetric flask. This makes sure that no contamination of equipment can occur. If contamination does occur, then the results obtained would be flawed, and thus, unreliable. The use of separate beakers ensures a fair test.

StopwatchIf there is some error in the calibration of the stopwatch that makes it unable to read the time accurately, then, as it will be used in every experiment, the errors will be relative to one another and it will not effect the final outcome greatly. Furthermore, due to human error, if the time is not measured accurately, then as these errors will be carried through in every experiment, the errors will be relative to one another and it will not effect the final outcome greatly. This ensures a fair test.

2 cuvettes

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For use in the colorimeter.

Before running the experiment, a preliminary experiment will be performed so that I can become accustomed to the equipment and carry out the experiment without any flaws. This also enables my efficiency and confidence with the equipment to increase.

Method 115 This method was initially proposed in Shakhashiri's Chemical Demonstrations, Vol. 4, pages 42-43. However, for this experiment, instead of using a buffer comprised of ethanoic acid and sodium ethanoate, I will use 0.3 mol dm-3 sulphuric acid. This is because the reaction requires H+ ions, which cannot be provided by this buffer solution quick enough for this reaction. I will measure the pH (and therefore, keep it constant each time) by using universal indicator paper.

15 Shakhashiri's Chemical Demonstrations, Vol. 4, pages 42-43

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Test tube containing the solutions

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Chemicals required (these will be made up as previously discussed):

150 cm3 of 0.0300 mol dm-3 hydrogen peroxide (H2O2) 50 cm3 of Starch Solution 100 cm3 0.00500 mol dm-3 sodium thiosulphate (Na2S2O3) 150 cm3 of 0.0200 mol dm-3 potassium iodide (KI) 150 cm3 of 0.300 mol dm-3 sulphuric acid (H2SO4)

Apparatus

A graduated pipette:This helps measure the volume of solution required to a high degree of accuracy. This will help in trying to make sure that the data collected is reliable by indirectly making sure that the total volume is kept the same each time (i.e. 14cm3).

Universal indicator paper:In each experiment, the pH of each reaction will be measured using universal indicator paper. It is essential that the pH is kept the same each time. The reason for this is that varying the pH could cause the results to become slanted and flawed, and so, to stop the results obtained from the experiment to be nullified, the pH must be fixed. This enables a fair test to be attained each time.

Test tube rackBoiling tubes can be stored in the test tube rack to prevent accidents

A mercury thermometer:A mercury thermometer is essential to this experiment, as it will determine when the reading should be taken. In the experiment, an analogue mercury thermometer will be used. This will be used because it measures to an accurate degree of accuracy (+ or – 0.5oC). It will make sure that the data that is collected is reliable as mercury thermometers provide an accurate measurement of the temperature.

15 boiling tubesThese will ensure that there is enough space for the reaction to proceed in without the fear of the solutions overflowing the boiling tube. Also, the boiling tube is easy to use and can easily be placed on a test tube rack.

A wash bottle of distilled water

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This will be used to dilute the solutions to the required concentration. They will also be used to clean the equipment each time, resulting in a fair test each time as no contamination will occur.

5 beakersTo avoid cross-contamination, the chemicals will be placed inside a beaker, and not be returned to the stock bottle. This makes sure that no contamination of equipment can occur. If contamination does occur, then the results obtained would be flawed, and thus, unreliable.

StopwatchIf there is some error in the calibration of the stopwatch that makes it unable to read the time accurately, then, as it will be used in every experiment, the errors will be relative to one another and it will not effect the final outcome greatly. Furthermore, due to human error, if the time is not measured accurately, then as these errors will be carried through in every experiment, the errors will be relative to one another and it will not effect the final outcome greatly. This ensures a fair test.

Method1. First place the solutions of hydrogen peroxide, starch solution,

sodium thiosulphate, sulphuric acid and water into a boiling tube (A) according to Table 1 (note: keep the potassium iodide separate; for Table 1, please refer to the appendices). The only concentrations that change are the volumes of hydrogen peroxide and water. The total volume stays the same.

2. The boiling tube will then be placed in a test tube rack. 3. Measure the temperature each time, and record this down. The

temperature must be kept constant in every experiment to keep a fair test.

4. Next, pipette 2cm3 of potassium iodide into a different boiling tube (B) (using table 1)

5. Pour this (B) into the other (A) and immediately start the timing from the time the solution is added.

6. Stir.7. Stop the time when the first blue colour appears and record this in

the table of results. 8. After the reaction finishes, check the pH and record it down to

ensure that the pH is kept constant in each experiment. This ensures a fair test.

9. Repeat the reaction twice, giving a total of three experiments per concentration of hydrogen peroxide. This helps to reduce

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anomalies and gives an accurate set of results, with the experiment being a fair test. If there are is an anomalous result, repeat the reading.

10. Next, plot a graph for the time against the concentration for each set of (averaged) results. A tangent will then be drawn on the first part of the graph. The gradient of this tangent will give the initial rate of the reaction. Doing this for each set of results will give initial rates for all of the experiments.

11. These would then be plotted against concentration, allowing the order of reaction to be found.

For the next set of results, the volume of potassium iodide will be varied, with the volume of hydrogen peroxide and sulphuric acid kept constant each time.

1. First place the solutions of hydrogen peroxide, starch solution, sodium thiosulphate, sulphuric acid and water into a boiling tube (A) according to Table 2 (note: keep the potassium iodide separate; for Table 2, please refer to the appendices). The only concentrations that change are the volumes of the potassium iodide and water. The total volume stays the same.

2. Place the boiling tube into a test tube rack. 3. Measure the temperature each time, and record this down. The

temperature must be kept constant in every experiment to keep a fair test.

4. Next, pipette the set volume of potassium iodide into a different boiling tube (B) (using Table 2).

5. Pour this (B) into the other boiling tube (A) and immediately start the timing from the time the solution is added.

6. Stir7. Stop the time when the first blue colour appears and record this in

the table of results. 8. After the reaction finishes, check the pH and record it down to

ensure that the pH is kept constant in each experiment. This ensures a fair test.

9. Repeat the reaction twice, giving a total of three experiments per concentration of potassium iodide. This helps to reduce anomalies and gives an accurate set of results, with the experiment being a fair test. If there are is an anomalous result, repeat the reading.

10. Next, plot a graph for the time against the concentration for each set of (averaged) results. A tangent will then be drawn on the first part of the graph. The gradient of this tangent will give the initial rate of the reaction. Doing this for each set of results will give initial rates for all of the experiments.

11. These would then be plotted against concentration, allowing the order of reaction to be found.

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For the next set of results, the volume of sulphuric acid will be varied, with the volume of hydrogen peroxide and potassium iodide kept constant each time.

1. First place the solutions of hydrogen peroxide, starch solution, sodium thiosulphate, sulphuric acid and water into a boiling tube (A) according to Table 3 (note: keep the potassium iodide separate; for Table 3, please refer to the appendices). The only concentrations that change are the volumes of the potassium iodide and water. The total volume stays the same.

2. Place this in the test tube rack. 3. Measure the temperature each time, and record this down. The

temperature must be kept constant in every experiment to keep a fair test.

4. Next, pipette the set volume of potassium iodide into a different boiling tube (B) (using Table 3).

5. Pour this (B) into the other boiling tube (A) and immediately start the timing from the time the solution is added.

6. Stir7. Stop the time when the first blue colour appears and record this in

the table of results. 8. After the reaction finishes, check the pH and record it down to

ensure that the pH is kept constant in each experiment. This ensures a fair test.

9. Repeat the reaction twice, giving a total of three experiments per concentration of sulphuric acid. This helps to reduce anomalies and gives an accurate set of results, with the experiment being a fair test. If there are is an anomalous result, repeat the reading.

10. Next, plot a graph for the time against the concentration for each set of (averaged) results. A tangent will then be drawn on the first part of the graph. The gradient of this tangent will give the initial rate of the reaction. Doing this for each set of results will give initial rates for all of the experiments.

11. These would then be plotted against concentration, allowing the order of reaction to be found.

Method 2

Planning 41

Colorimeter

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To corroborate the results from method 1, the measurement of the absorbency of the solution will now be measured. This absorbency will be plotted against time, and a tangent from the graph taken to give the initial rate. These initial rates will be plotted against their corresponding concentrations. A calibration curve will also be obtained. In this experiment, a 470nm filter will be used each time.

Chemicals required (these will be made up as previously discussed):

150 cm3 of 0.0300 mol dm-3 Hydrogen Peroxide (H2O2) 150 cm3 of 0.0200 mol dm-3 Potassium Iodide (KI) 150 cm3 of 0.300 mol dm-3 of sulphuric acid (H2SO4)

Apparatus

A plastic pipette:This helps to add each solution drop by drop, to a high degree of accuracy.

Universal indicator paper:In each experiment, the pH of each reaction will be measured using universal indicator paper. It is essential that the pH is kept the same each time. The reason for this is that varying the pH could cause the results to become slanted and flawed, and so, to stop the results obtained from the experiment to be nullified, the pH must be fixed. This enables a fair test to be attained each time.

Test tube rackBoiling tubes can be stored in the test tube rack to prevent accidents

A mercury thermometer:A mercury thermometer is essential to this experiment, as it will determine when the reading should be taken. In the experiment, an analogue mercury thermometer will be used. This will be used because it measures to an accurate degree of accuracy (+ or – 0.5oC). It will make sure that the data that is collected is reliable as mercury thermometers provide an accurate measurement of the temperature.

Planning 42

Cuvette containing solution

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A colorimeterA standard colorimeter with a blue-green filter will be used. This is because it provides an accurate reading of the percentage absorbency, as the light emitted by the colorimeter is the same as that absorbed by the substance. The results obtained will be reliable each time, because, as the colorimeter is the same, the calibration from the manufacturers will also be the same, causing reliable results to be attained.

A wash bottle of distilled waterThis will be used to dilute the solutions to the required concentration. They will also be used to clean the equipment each time, resulting in a fair test each time as no contamination will occur.

StopwatchIf there is some error in the calibration of the stopwatch that makes it unable to read the time accurately, then, as it will be used in every experiment, the errors will be relative to one another and it will not effect the final outcome greatly. Furthermore, due to human error, if the time is not measured accurately, then as these errors will be carried through in every experiment, the errors will be relative to one another and it will not effect the final outcome greatly. This ensures a fair test.

2 cuvettesFor use in the colorimeter.

Method

1. First place the solutions of hydrogen peroxide, sodium thiosulphate, sulphuric acid and water into a boiling tube (A) according to Table 4 (note: keep the potassium iodide separate; for Table 4, please refer to the appendices). The only concentrations that change are the volumes of the hydrogen peroxide and water. The total volume stays the same.

2. Place this in the test tube rack. 3. Measure the temperature each time, and record this down. The

temperature must be kept constant in every experiment to keep a fair test.

4. Next, pipette the set volume of potassium iodide into a different boiling tube (B) (using Table 4).

5. Calibrate and reset the colorimeter using a cuvette filled with distilled water.

6. Place a new cuvette into the colorimeter.

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7. Pour the solution from one boiling tube (B) into the other boiling tube (A) and immediately start the timing from the time the solution is added.

8. Quickly transfer some of the solution from the boiling tube to the cuvette in the colorimeter

9. Every ten seconds measure the percentage absorbency of the solution, and record this in a table.

10. Stop after 120 seconds. 11. After the reaction finishes, check the pH and record it down to

ensure that the pH is kept constant in each experiment. This ensures a fair test.

12. Repeat the reaction twice, giving a total of three experiments per concentration of hydrogen peroxide. This helps to reduce anomalies and gives an accurate set of results, with the experiment being a fair test. If there are is an anomalous result, repeat the reading.

13. Next, plot a graph for the time against the % absorbency for each set of (averaged) results. A tangent will then be drawn on the first part of the graph. The gradient of this tangent will give the initial rate of the reaction. Doing this for each set of results will give initial rates for all of the experiments.

14. These would then be plotted against concentration, allowing the order of reaction to be found.

For the next set of results, the volume of potassium iodide will be varied, with the volume of hydrogen peroxide kept constant each time.

1. First place the solutions of hydrogen peroxide, sodium thiosulphate, sulphuric acid and water into a boiling tube (A) according to Table 5 (note: keep the potassium iodide separate; for Table 5, please refer to the appendices). The only concentrations that change are the volumes of the hydrogen peroxide and water. The total volume stays the same.

2. Place this in the test tube rack. 3. Measure the temperature each time, and record this down. The

temperature must be kept constant in every experiment to keep a fair test.

4. Next, pipette the set volume of potassium iodide into a different boiling tube (B) (using Table 5).

5. Calibrate and reset the colorimeter using a cuvette filled with distilled water.

6. Place a new cuvette into the colorimeter.7. Pour the solution from one boiling tube (B) into the other boiling

tube (A) and immediately start the timing from the time the solution is added.

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8. Quickly transfer some of the solution from the boiling tube to the cuvette in the colorimeter

9. Every ten seconds measure the percentage absorbency of the solution, and record this in a table.

10. Stop after 120 seconds. 11. After the reaction finishes, check the pH and record it down to

ensure that the pH is kept constant in each experiment. This ensures a fair test.

12. Repeat the reaction twice, giving a total of three experiments per concentration of hydrogen peroxide. This helps to reduce anomalies and gives an accurate set of results, with the experiment being a fair test. If there are is an anomalous result, repeat the reading.

13. Next, plot a graph for the time against the % absorbency for each set of (averaged) results. A tangent will then be drawn on the first part of the graph. The gradient of this tangent will give the initial rate of the reaction. Doing this for each set of results will give initial rates for all of the experiments.

14. These would then be plotted against concentration, allowing the order of reaction to be found.

For the next set of results, the volume of potassium iodide will be varied, with the volume of hydrogen peroxide kept constant each time.

1. First place the solutions of hydrogen peroxide, sodium thiosulphate, sulphuric acid and water into a boiling tube (A) according to Table 6 (note: keep the potassium iodide separate; for Table 6, please refer to the appendices). The only concentrations that change are the volumes of the hydrogen peroxide and water. The total volume stays the same.

2. Place this in the test tube rack. 3. Measure the temperature each time, and record this down. The

temperature must be kept constant in every experiment to keep a fair test.

4. Next, pipette the set volume of potassium iodide into a different boiling tube (B) (using Table 6).

5. Calibrate and reset the colorimeter using a cuvette filled with distilled water.

6. Place a new cuvette into the colorimeter.7. Pour the solution from one boiling tube (B) into the other boiling

tube (A) and immediately start the timing from the time the solution is added.

8. Quickly transfer some of the solution from the boiling tube to the cuvette in the colorimeter

Planning 45

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9. Every ten seconds measure the percentage absorbency of the solution, and record this in a table.

10. Stop after 120 seconds. 11. After the reaction finishes, check the pH and record it down to

ensure that the pH is kept constant in each experiment. This ensures a fair test.

12. Repeat the reaction twice, giving a total of three experiments per concentration of hydrogen peroxide. This helps to reduce anomalies and gives an accurate set of results, with the experiment being a fair test. If there are is an anomalous result, repeat the reading.

13. Next, plot a graph for the time against the % absorbency for each set of (averaged) results. A tangent will then be drawn on the first part of the graph. The gradient of this tangent will give the initial rate of the reaction. Doing this for each set of results will give initial rates for all of the experiments.

14. These would then be plotted against concentration, allowing the order of reaction to be found.

General pointsFor both methods, it is important to store the hydrogen peroxide in a brown stock bottle to stop it decomposing with light. For both sets of results, draw graphs of time versus concentration (for method 1), and percentage absorbency against concentration (method 2).

Obtaining a calibration curveA calibration curve is necessary to corroborate the results of the results obtained by the colorimeter work. This ensures that, for a gradual increase in concentration of a coloured solution (in this case, iodine), the absorbency also gradually increases. That is, concentration α % absorbency. The results for each will be recorded into a table (Appendix 10). A graph will then be plotted to show the relationship. 15 measurements will be taken, with the first reading at 0.000100 mol dm-3 and the last reading at 0.001500 mol dm-3, at 0.000100 mol dm-3 intervals. The reason for this is that iodine solutions of higher concentration are too concentrated and too dark in colour, and so, give absorbencies of 2.00%. To make the iodine solution, I will need to follow the recipe sheet. Therefore, according to that, I will need to make up the iodine solution in the following way:

1. Measure out 3.00g of potassium iodide and place into a beaker

Planning 46

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2. Add about 65.0 cm3 of water3. Measure out 2.54g of iodine granules and add this to the solution

in the beaker.4. Stir using a magnetic stirrer for several minutes.5. Pour the solution into a 100 cm3 volumetric flask and make up to

100 cm3

6. This produces a 0.1 mol dm-3 solution of iodine.7. Take out 1.5 cm3 of the iodine solution using a graduated pipette,

and add this to another 100cm3 volumetric flask.8. Make up to the 100 cm3 mark.9. This makes a 0.001500 mol dm-3 iodine solution

The method to perform the experiment is below:

1. Reset the colorimeter by putting a cuvette filled with distilled water into the colorimeter and press reset.

2. Put a fresh cuvette into the colorimeter3. Make up the required solution as table 8 (please refer to

appendices) in a 50 cm3 beaker4. Pipette some into a cuvette5. Measure the absorbency and record into a table (Appendix 10)6. Draw a graph of absorbency against concentration

Extension task - 1For this the affect of temperature will be looked at. Method 1, the clock collection method will be used for this. This is because it will be much easier for the boiling tubes to be kept in a water bath, compared to putting a cuvette and colorimeter in a water bath. The step-by-step method outlined in method 1 should be followed again. The concentration of hydrogen peroxide and potassium iodide will be kept constant each time (see table 7). The only variable will be the temperature that the solution is in. This will be changed using a water bath. The time will start when the solutions are mixed, and will be stopped once the solution turns blue-black. Each experiment will be repeated at least 3 times.

Extension task - 2For this, I will experiment by using a different filter (490nm) to the one I will use throughout the main part of the experiments (470nm) and test the absorbency (using table 4) of the set solutions. I will then compare this with the results obtained using the 470nm filter.

Planning 47

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Synoptic table

Concept Unit met Module YearRate orders Engineering

proteins2849 A2

Redox reactions (reduction and oxidation)

From Minerals to Elements

2848 AS

Absorption Spectrum Elements of Life 2850 ASHalf-equations Steel Story 2849 A2Rate-concentration graphs

Engineering Proteins

2849 A2

Rate-determining step

Engineering Proteins

2849 A2

Rate-mechanisms Engineering Proteins

2849 A2

Half-life Engineering Proteins

2849 A2

Intermediates The Atmosphere 2848 ASReactions incorporating colour changes

Steel Story 2849 A2

Ionic lattices From Minerals to Elements

2848 AS

Affect of concentrations on rate

Minerals to elements

2848 AS

The mole Elements of life 2850 ASKinetic theory Engineering

proteins2849 A2

Boltzmann Distribution

Engineering proteins

2849 A2

Affect of pressure on rate

The atmosphere 2848 AS

Factors affecting rate

The atmosphere 2848 AS

Affect of temperature on rate

The atmosphere 2848 AS

Affect of catalysts on rate

Developing fuels and the atmosphere

2850 + 2848 AS

Van der Waals forces Polymer revolution

2848 AS

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Making up a standard solution

Elements of Life 2850 AS

Using colorimeter Steel Story 2849 A2Measuring the time taken for a reaction to reach a particular stage

Aspects of Agriculture

2854 A2

Arrhenius equation Aspects of agriculture

2854 A2

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Preliminary ExperimentTo ensure that all my equipment works without glitches and to ensure that my results will be consistent, I will perform a preliminary experiment. This helps me to, not only identify areas of improvement for my experiment, but also assists me to familiarise myself with the equipment.

I will run this experiment with the following concentrations of substance:

Hydrogen peroxide: 5.00 cm3

Potassium iodide: 2.00 cm3

Starch solution: 2.00 cm3

Sulphuric acid: 4.00 cm3

Sodium thiosulphate: 1.00 cm3

I will follow method 1 from my plan for this experiment.

Results

Time for reaction to happen (seconds)

1st 2nd 3rd Average

60.67 59.79

58.71

59.72

pH 1.0 1.0 1.0 1.0Temperature (K)

294 294 294 294

Planning 50

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Changes after the ExperimentFurther to my preliminary experiment, I was happy with the way the experiment was running. However, I need to make the following changes to my experiment:

Hydrogen peroxideThe reaction was too slow, and therefore, I had to increase the concentration of the hydrogen peroxide to 0.100 mol dm-3. To do this, I needed:

Mols of hydrogen peroxide required = concentration required x volume required

= 0.1 x 0.25= 0.0250 mol

The volume required would then equal the number of moles/concentration given. In this experiment, the concentration of hydrogen peroxide given is 1.67 mol dm-3.

Volume required = mol/given concentration= 0.0250/1.67= 15.0 cm3

Potassium iodideFurthermore, the concentration of the potassium iodide was also too weak, and so, this also needed to be upgraded to 0.200 mol dm-3. To do this:

Mass required = volume required x concentration required x Mr,

= 0.200 x 0.250 x 166= 8.30g

Therefore, 8.30g of potassium iodide will be added to a beaker, and then about 150 cm3 of water added. This solution will then be transferred to a 250cm3 volumetric flask using a glass funnel, and the beaker, glass funnel and stirring rod will then be repeatedly washed using distilled water. These washings will then be added to the volumetric flask, until the solution is made up to 250 cm3. This forms 0.200 mol dm-3 of potassium iodide solution, so a further ten-fold dilution will be made, by taking 25.0 cm3 out of the 0.200 mol dm-3 potassium iodide solution and putting this inside another 250 cm3

Planning 51

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volumetric flask. This solution will then be made up to 250cm3 using distilled water.

Planning 52

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Implementing - Method 1

Changing the concentration of hydrogen peroxideFor the first set of experiments, the following volumes of solutions were used:

Hydrogen peroxide: 5.00 cm3

Potassium iodide: 2.00 cm3

Starch solution: 2.00 cm3

Sulphuric acid: 4.00 cm3

Sodium thiosulphate: 1.00 cm3

Time for reaction to happen (seconds)

1st 2nd 3rd Average53.13 56.12 58.59 55.95

pH 1.0 1.0 1.0 1.0Temperature (K)

294.0 294.0 294.0 294.0

When I initially conducted this experiment, I got the following anomalous results:

Time for reaction to happen (seconds)

1st 2nd 3rd Average59.15 63.21 62.63 61.66

pH 1.0 1.0 1.0 1.0Temperature (K)

294.0 294.0 294.0 294.0

These were anomalous, because the rate for this set of results was off the line of best fit. I put these anomalies down to teething errors with the beginning of the experiment.

For the next set of experiments, the following volumes of solutions were used:

Hydrogen peroxide: 4.00 cm3

Potassium iodide: 2.00 cm3

Starch solution: 2.00 cm3

Implementing 1

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Sulphuric acid: 4.00 cm3

Sodium thiosulphate: 1.00 cm3

Water: 1.00 cm3

Time for reaction to happen (seconds)

1st 2nd 3rd Average70.68 72.25 76.81 73.25

pH 1.0 1.0 1.0 1.0Temperature (K)

294.0 294.0 294.0 294.0

Implementing 2

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For the next set of experiments, the following volumes of solutions were used:

Hydrogen peroxide: 3.00 cm3

Potassium iodide: 2.00 cm3

Starch solution: 2.00 cm3

Sulphuric acid: 4.00 cm3

Sodium thiosulphate: 1.00 cm3

Water: 2.00 cm3

Time for reaction to happen (seconds)

1st 2nd 3rd Average92.22 93.22 92.09 92.51

pH 1.0 1.0 1.0 1.0Temperature (K)

294.0 294.0 294.0 294.0

For the next set of experiments, the following volumes of solutions were used:

Hydrogen peroxide: 2.00 cm3

Potassium iodide: 2.00 cm3

Starch solution: 2.00 cm3

Sulphuric acid: 4.00 cm3

Sodium thiosulphate: 1.00 cm3

Water: 3.00 cm3

Time for reaction to happen (seconds)

1st 2nd 3rd Average141.25 142.94 139.15 141.11

pH 1.0 1.0 1.0 1.0Temperature (K)

294.0 294.0 294.0 294.0

For the next set of experiments, the following volumes of solutions were used:

Hydrogen peroxide: 1.00 cm3

Potassium iodide: 2.00 cm3

Starch solution: 2.00 cm3

Sulphuric acid: 4.00 cm3

Implementing 3

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Sodium thiosulphate: 1.00 cm3

Water: 4.00 cm3

Time for reaction to happen (seconds)

1st 2nd 3rd Average297.31 301.14 300.17 299.54

pH 1.0 1.0 1.0 1.0Temperature (K)

294.0 294.0 294.0 294.0

The table for the average set of results for each concentration for hydrogen peroxide is below:

Concentration of hydrogen peroxide (mol dm-3)

Average time taken (seconds)

0.357 55.950.286 73.250.214 92.510.143 141.110.0714 299.54

(In the table above, the concentration of hydrogen peroxide is worked out by dividing the volume of hydrogen peroxide in the solution by the total volume).

Working out the rate of reaction for each concentration

The rate for a clock reaction can be worked out by dividing the concentration of iodine produced by the average time taken for the reaction to occur. Looking at the equation:

I3-(aq) + 2S2O3

2-(aq) 3I-

(aq) + S4O6-(aq)

The stoichiometry shows that, for every one mole of iodine produced, two moles of S2O3

2- ions are produced. Therefore, to work out the concentration of iodine produced, the following analysis can be made:

[S2O32-]/2Δt,

The thiosulphate concentration for each is 0.00500 mol dm-3, and therefore, the concentration of the iodine produced is 0.00250 mol dm-3.

Implementing 4

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Using this equation and the above information, the following table can be constructed:

Concentration of hydrogen peroxide (mol dm-3)

Average time taken (seconds)

Calculation for rate

Rate (mol s-1)

0.0357 55.95 0.0025/55.95 4.47 x 10-5

0.0286 73.25 0.0025/73.25 3.41 x 10-5

0.0214 92.51 0.0025/92.51 2.70 x 10-5

0.0143 141.11 0.0025/141.11 1.77 x 10-5

0.00714 299.54 0.0025/299.54 8.35 x 10-6

The graphs for hydrogen peroxide are shown overleaf.

Implementing 5

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Changing the concentration of potassium iodideFor the first set of experiments, the following volumes of solutions were used:

Potassium iodide: 5.00 cm3

Hydrogen peroxide: 2.00 cm3

Starch solution: 2.00 cm3

Sulphuric acid: 4.00 cm3

Sodium thiosulphate: 1.00 cm3

Time for reaction to happen (seconds)

1st 2nd 3rd Average64.75 67.37 66.10 66.07

pH 1.0 1.0 1.0 1.0Temperature (K)

294.0 294.0 294.0 294.0

For the next set of experiments, the following volumes of solutions were used:

Potassium iodide: 4.00 cm3

Hydrogen peroxide: 2.00 cm3

Starch solution: 2.00 cm3

Sulphuric acid: 4.00 cm3

Sodium thiosulphate: 1.00 cm3

Water: 1.00 cm3

Time for reaction to happen (seconds)

1st 2nd 3rd Average80.38 81.97 83.56 81.97

PH 1.0 1.0 1.0 1.0Temperature (K)

294.0 294.0 294.0 294.0

Implementing 6

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For the next set of experiments, the following volumes of solutions were used:

Potassium iodide: 3.00 cm3

Hydrogen peroxide: 2.00 cm3

Starch solution: 2.00 cm3

Sulphuric acid: 4.00 cm3

Sodium thiosulphate: 1.00 cm3

Water: 2.00 cm3

Time for reaction to happen (seconds)

1st 2nd 3rd Average108.16 108.63 108.01 108.27

PH 1.0 1.0 1.0 1.0Temperature (K)

294.0 294.0 294.0 294.0

For the next set of experiments, the following volumes of solutions were used:

Potassium iodide: 2.00 cm3

Hydrogen peroxide: 2.00 cm3

Starch solution: 2.00 cm3

Sulphuric acid: 4.00 cm3

Sodium thiosulphate: 1.00 cm3

Water: 3.00 cm3

Time for reaction to happen (seconds)

1st 2nd 3rd Average143.32 143.02 141.71 142.68

PH 1.0 1.0 1.0 1.0Temperature (K)

294.0 294.0 294.0 294.0

For the next set of experiments, the following volumes of solutions were used:

Potassium iodide: 1.00 cm3

Hydrogen peroxide: 2.00 cm3

Starch solution: 2.00 cm3

Sulphuric acid: 4.00 cm3

Implementing 7

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Sodium thiosulphate: 1.00 cm3

Water: 4.00 cm3

Time for reaction to happen (seconds)

1st 2nd 3rd Average303.94 308.87 297.06 303.29

pH 1.0 1.0 1.0 1.0Temperature (K)

294.0 294.0 294.0 294.0

The table for the average set of results for each concentration for hydrogen peroxide is below:

Concentration of potassium iodide (mol dm-3)

Average time taken (seconds)

0.0714 66.070.0571 81.970.0429 108.270.0286 142.680.0143 303.29

(In the table above, the concentration of hydrogen peroxide is worked out by dividing the volume of hydrogen peroxide in the solution by the total volume).

Working out the rate of reaction for each concentration

The rate for a clock reaction can be worked out by dividing the concentration of iodine produced by the average time taken for the reaction to occur. Looking at the equation:

I3-(aq) + 2S2O3

2-(aq) 3I-

(aq) + S4O6-(aq)

The stoichiometry shows that, for every one mole of iodine produced, two moles of S2O3

2- ions are produced. Therefore, to work out the concentration of iodine produced, the following analysis can be made:

[S2O32-]/2Δt,

The thiosulphate concentration for each is 0.00500 mol dm-3, and therefore, the concentration of the iodine produced is 0.00250 mol dm-3.

Implementing 8

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Using this equation and the above information, the following table can be constructed:

Concentration of potassium iodide (mol dm-3)

Average time taken (seconds)

Calculation for rate

Rate (mol s-1)

0.0714 66.07 0.0025/55.95 3.78 x 10-5

0.0571 81.97 0.0025/73.25 3.05 x 10-5

0.0429 108.27 0.0025/92.51 2.31 x 10-5

0.0286 142.68 0.0025/141.11 1.75 x 10-5

0.0143 303.29 0.0025/299.54 8.24 x 10-6

The graphs for potassium iodide are shown overleaf.

Implementing 9

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Changing the concentration of H+ ionsFor the first set of experiments, the following volumes of solutions were used:

Sulphuric acid: 5.00 cm3

Hydrogen peroxide: 3.00 cm3

Potassium iodide: 3.00 cm3

Starch solution: 2.00 cm3

Sodium thiosulphate: 1.00 cm3

Time for reaction to happen (seconds)

1st 2nd 3rd Average95.15 96.37 96.93 96.15

pH 1.0 1.0 1.0 1.0Temperature (K)

294.0 294.0 294.0 294.0

When I initially conducted this experiment, I got the following anomalous results:

Time for reaction to happen (seconds)

1st 2nd 3rd Average65.03 63.21 62.63 63.62

pH 1.0 1.0 1.0 1.0Temperature (K)

295.0 293.5 292.0 293.5

These were anomalous, because the rate for this set of results was off the line of best fit. I put these anomalies down to the fluctuation of the temperature in each experiment.

For the next set of experiments, the following volumes of solutions were used:

Sulphuric acid: 4.00 cm3

Hydrogen peroxide: 3.00 cm3

Potassium iodide: 3.00 cm3

Starch solution: 2.00 cm3

Sodium thiosulphate: 1.00 cm3

Water: 1.00 cm3

Implementing 11

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Time for reaction to happen (seconds)

1st 2nd 3rd Average108.29 106.00 109.27 107.85

PH 1.0 1.0 1.0 1.0Temperature (K)

294.0 294.0 294.0 294.0

Implementing 12

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For the next set of experiments, the following volumes of solutions were used:

Sulphuric acid: 3.00 cm3

Hydrogen peroxide: 3.00 cm3

Potassium iodide: 3.00 cm3

Starch solution: 2.00 cm3

Sodium thiosulphate: 1.00 cm3

Water: 2.00 cm3

Time for reaction to happen (seconds)

1st 2nd 3rd Average121.32 122.49 119.21 121.01

PH 1.0 1.0 1.0 1.0Temperature (K)

294.0 294.0 294.0 294.0

For the next set of experiments, the following volumes of solutions were used:

Sulphuric acid: 2.00 cm3

Hydrogen peroxide: 3.00 cm3

Potassium iodide: 3.00 cm3

Starch solution: 2.00 cm3

Sodium thiosulphate: 1.00 cm3

Water: 3.00 cm3

Time for reaction to happen (seconds)

1st 2nd 3rd Average137.78 136.08 138.96 137.61

PH 1.0 1.0 1.0 1.0Temperature (K)

294.0 294.0 294.0 294.0

For the next set of experiments, the following volumes of solutions were used:

Sulphuric acid: 1.00 cm3

Hydrogen peroxide: 3.00 cm3

Potassium iodide: 3.00 cm3

Starch solution: 2.00 cm3

Implementing 13

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Sodium thiosulphate: 1.00 cm3

Water: 4.00 cm3

Time for reaction to happen (seconds)

1st 2nd 3rd Average172.04 172.19 173.00 172.41

pH 1.0 1.0 1.0 1.0Temperature (K)

294.0 294.0 294.0 294.0

When I first conducted this experiment, I obtained these anomalies:

Time for reaction to happen (seconds)

1st 2nd 3rd Average208.22 208.38 210.64 209.08

pH 1.0 1.0 1.0 1.0Temperature (K)

294.5 293.0 292.0 293.0

I put this down to the fluctuations in the temperature.

The table for the average set of results for each concentration for hydrogen peroxide is below:

Concentration of sulphuric acid (mol dm-3)

Average time taken (seconds)

0.179 96.150.143 107.850.107 121.010.0714 137.610.0357 172.41

(In the table above, the concentration of hydrogen peroxide is worked out by dividing the volume of hydrogen peroxide in the solution by the total volume).

Working out the rate of reaction for each concentration

Implementing 14

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The rate for a clock reaction can be worked out by dividing the concentration of iodine produced by the average time taken for the reaction to occur. Looking at the equation:

I3-(aq) + 2S2O3

2-(aq) 3I-

(aq) + S4O6-(aq)

The stoichiometry shows that, for every one mole of iodine produced, two moles of S2O3

2- ions are produced. Therefore, to work out the concentration of iodine produced, the following analysis can be made:

[S2O32-]/2Δt,

The thiosulphate concentration for each is 0.00500 mol dm-3, and therefore, the concentration of the iodine produced is 0.00250 mol dm-3.

Implementing 15

Page 67: Hydrogen Peroxide Iodine Clock

Using this equation and the above information, the following table can be constructed:

Concentration of sulphuric acid (mol dm-

3)

Average time taken (seconds)

Calculation for rate

Rate (mol s-1)

0.179 96.15 0.0025/96.15 2.60 x 10-5

0.143 107.85 0.0025/107.85 2.32 x 10-5

0.107 121.01 0.0025/121.01 2.07 x 10-5

0.0714 137.61 0.0025/137.61 1.82 x 10-5

0.0357 172.41 0.0025/172.41 1.45 x 10-5

The graphs for H+ ions are shown overleaf.

Implementing 16

Page 68: Hydrogen Peroxide Iodine Clock

Implementing - Method 2

Changing the concentration of hydrogen peroxide

For the first set of experiments, the following volumes of solutions were used:

Hydrogen peroxide: 5.00 cm3

Potassium iodide: 2.00 cm3

Sulphuric acid: 4.00 cm3

Starch solution: 2.00 cm3

Sodium thiosulphate: 1.00 cm3

The concentration of hydrogen peroxide used for this experiment was 0.0357 mol dm-3

% AbsorbencyTime (seconds) Run 1 Run 2 Run 3 Average

absorbency10 0.14 0.14 0.13 0.1420 0.32 0.27 0.28 0.2930 0.40 0.41 0.39 0.4040 0.58 0.58 0.57 0.5850 0.73 0.72 0.74 0.7360 0.89 0.87 0.88 0.8870 1.01 1.01 1.02 1.0180 1.10 1.18 1.17 1.1590 1.24 1.22 1.23 1.23100 1.28 1.29 1.28 1.28110 1.32 1.30 1.31 1.31120 1.35 1.35 1.35 1.35Temperature (K)

294 294 294 294

pH 1.0 1.0 1.0 1.0

Implementing 17

Page 69: Hydrogen Peroxide Iodine Clock

When I initially conducted this experiment, I got the following anomalous result:

Time (seconds) Result10 0.1820 0.4230 0.5640 0.7850 0.9960 1.1070 1.2180 1.3190 1.41100 1.50110 1.58120 1.57Temperature (K)

294

pH 1.0

As this does not fit the general trend of my results, I put this anomalous result down to the initial teething errors of beginning the experiment and getting used to it.

The graph of % absorbency against time for this concentration of hydrogen peroxide is shown overleaf

Implementing 18

Page 70: Hydrogen Peroxide Iodine Clock

For the next set of experiments, the following volumes of solutions were used:

Hydrogen peroxide: 4.00 cm3

Potassium iodide: 2.00 cm3

Sulphuric acid: 4.00 cm3

Starch solution: 2.00 cm3

Sodium thiosulphate: 1.00 cm3

Distilled water: 1.00 cm3

The concentration of hydrogen peroxide used for this experiment was 0.0286 mol dm-3

% AbsorbencyTime (seconds) Run 1 Run 2 Run 3 Average

absorbency10 0.11 0.11 0.12 0.1120 0.22 0.21 0.26 0.2330 0.35 0.34 0.35 0.3540 0.45 0.48 0.45 0.4650 0.56 0.60 0.56 0.5760 0.72 0.71 0.70 0.7070 0.81 0.81 0.82 0.8180 0.91 0.91 0.92 0.9190 0.98 1.00 0.99 0.99100 1.03 1.03 1.02 1.03110 1.07 1.08 1.07 1.07120 1.12 1.14 1.10 1.12Temperature (K)

294 294 294 294

pH 1.0 1.0 1.0 1.0

The graph of % absorbency against time for this concentration of hydrogen peroxide is shown overleaf

Implementing 20

Page 71: Hydrogen Peroxide Iodine Clock

For the next set of experiments, the following volumes of solutions were used:

Hydrogen peroxide: 3.00 cm3

Potassium iodide: 2.00 cm3

Sulphuric acid: 4.00 cm3

Starch solution: 2.00 cm3

Sodium thiosulphate: 1.00 cm3

Distilled water: 2.00 cm3

The concentration of hydrogen peroxide used for this experiment was 0.0214 mol dm-3

% AbsorbencyTime (seconds) Run 1 Run 2 Run 3 Average

absorbency10 0.13 0.12 0.11 0.1220 0.21 0.21 0.19 0.2030 0.29 0.29 0.27 0.2840 0.37 0.39 0.33 0.3650 0.45 0.47 0.43 0.4560 0.52 0.55 0.50 0.5270 0.60 0.63 0.58 0.6080 0.66 0.69 0.64 0.6690 0.72 0.76 0.70 0.73100 0.78 0.83 0.76 0.79110 0.84 0.89 0.82 0.85120 0.90 0.95 0.88 0.91Temperature (K)

294 294 294 294

pH 1.0 1.0 1.0 1.0

The graph of % absorbency against time for this concentration of hydrogen peroxide is shown overleaf

Implementing 22

Page 72: Hydrogen Peroxide Iodine Clock

For the next set of experiments, the following volumes of solutions were used:

Hydrogen peroxide: 2.00 cm3

Potassium iodide: 2.00 cm3

Sulphuric acid: 4.00 cm3

Starch solution: 2.00 cm3

Sodium thiosulphate: 1.00 cm3

Distilled water: 3.00 cm3

The concentration of hydrogen peroxide used for this experiment was 0.0143 mol dm-3

% AbsorbencyTime (seconds) Run 1 Run 2 Run 3 Average

absorbency10 0.08 0.1 0.09 0.0920 0.17 0.18 0.17 0.1730 0.26 0.25 0.25 0.2540 0.36 0.35 0.35 0.3550 0.43 0.46 0.41 0.4360 0.51 0.52 0.5 0.5170 0.61 0.6 0.59 0.680 0.68 0.68 0.67 0.6890 0.77 0.73 0.75 0.75100 0.82 0.81 0.81 0.81110 0.85 0.85 0.84 0.85120 0.88 0.9 0.89 0.89Temperature (K)

294 294 294 294

pH 1.0 1.0 1.0 1.0

The graph of % absorbency against time for this concentration of hydrogen peroxide is shown overleaf

Implementing 24

Page 73: Hydrogen Peroxide Iodine Clock

For the next set of experiments, the following volumes of solutions were used:

Hydrogen peroxide: 1.00 cm3

Potassium iodide: 2.00 cm3

Sulphuric acid: 4.00 cm3

Starch solution: 2.00 cm3

Sodium thiosulphate: 1.00 cm3

Distilled water: 4.00 cm3

The concentration of hydrogen peroxide used for this experiment was 0.00714 mol dm-3

% AbsorbencyTime (seconds) Run 1 Run 2 Run 3 Average

absorbency10 0.06 0.06 0.07 0.0620 0.13 0.11 0.12 0.1230 0.18 0.17 0.19 0.1840 0.24 0.23 0.22 0.2350 0.29 0.30 0.30 0.3060 0.35 0.35 0.36 0.3570 0.40 0.42 0.41 0.4180 0.47 0.48 0.46 0.4790 0.52 0.50 0.51 0.51100 0.54 0.55 0.56 0.54110 0.59 0.60 0.58 0.59120 0.63 0.62 0.64 0.63Temperature (K)

294 294 294 294

pH 1.0 1.0 1.0 1.0

The graph of % absorbency against time for this concentration of hydrogen peroxide is shown overleaf.

Implementing 26

Page 74: Hydrogen Peroxide Iodine Clock

Changing the concentration of potassium iodide

For the first set of experiments, the following volumes of solutions were used:

Potassium iodide: 5.00 cm3

Hydrogen peroxide: 2.00 cm3

Sulphuric acid: 4.00 cm3

Starch solution: 2.00 cm3

Sodium thiosulphate: 1.00 cm3

The concentration of potassium iodide used for this experiment was 0.0714 mol dm-3

% AbsorbencyTime (seconds) Run 1 Run 2 Run 3 Average

absorbency10 0.16 0.14 0.15 0.1520 0.28 0.29 0.27 0.2830 0.43 0.42 0.38 0.4140 0.54 0.55 0.48 0.5250 0.65 0.66 0.68 0.6660 0.75 0.76 0.76 0.7670 0.85 0.88 0.79 0.8480 0.93 1.00 0.90 0.9490 1.01 0.98 0.99 0.99100 1.09 1.08 1.07 1.08110 1.12 1.13 1.12 1.12120 1.17 1.18 1.16 1.17Temperature (K)

294 294 294 294

PH 1.0 1.0 1.0 1.0

The graph of % absorbency against time for this concentration of potassium iodide is shown overleaf

Implementing 28

Page 75: Hydrogen Peroxide Iodine Clock

For the next set of experiments, the following volumes of solutions were used:

Potassium iodide: 4.00 cm3

Hydrogen peroxide: 2.00 cm3

Sulphuric acid: 4.00 cm3

Starch solution: 2.00 cm3

Sodium thiosulphate: 1.00 cm3

Distilled water: 1.00 cm3

The concentration of potassium iodide used for this experiment was 0.0571 mol dm-3

% AbsorbencyTime (seconds) Run 1 Run 2 Run 3 Average

absorbency10 0.13 0.16 0.12 0.1420 0.24 0.29 0.20 0.2430 0.34 0.34 0.29 0.3240 0.44 0.48 0.45 0.4650 0.52 0.56 0.48 0.5260 0.61 0.68 0.63 0.6470 0.71 0.74 0.69 0.7180 0.81 0.81 0.82 0.8190 0.90 0.90 0.89 0.90100 0.90 0.96 0.99 0.95110 0.96 0.98 0.97 0.97120 1.00 1.01 1.00 1.00Temperature (K)

294 294 294 294

pH 1.0 1.0 1.0 1.0

Implementing 30

Page 76: Hydrogen Peroxide Iodine Clock

When I initially conducted this experiment, I got the following anomalous result:

Time (seconds) Result10 0.1720 0.3530 0.5740 0.8450 1.0160 1.1970 1.3480 1.4590 1.54100 1.60110 1.65120 1.69Temperature (K)

294

pH 1.0

As this does not fit the general trend of my results, I put this anomalous result down to the fact that I had to move the solutions quickly, introducing errors into the experiment.

The graph of % absorbency against time for this concentration of potassium iodide is shown overleaf

Implementing 31

Page 77: Hydrogen Peroxide Iodine Clock

For the next set of experiments, the following volumes of solutions were used:

Potassium iodide: 3.00 cm3

Hydrogen peroxide: 2.00 cm3

Sulphuric acid: 4.00 cm3

Starch solution: 2.00 cm3

Sodium thiosulphate: 1.00 cm3

Distilled water: 2.00 cm3

The concentration of potassium iodide used for this experiment was 0.0429 mol dm-3

% AbsorbencyTime (seconds) Run 1 Run 2 Run 3 Average

absorbency10 0.11 0.10 0.08 0.1020 0.19 0.20 0.15 0.1830 0.26 0.29 0.21 0.2540 0.33 0.34 0.28 0.3250 0.41 0.38 0.44 0.4160 0.48 0.47 0.40 0.4570 0.54 0.55 0.54 0.5480 0.61 0.62 0.61 0.6190 0.67 0.69 0.68 0.68100 0.73 0.76 0.71 0.73110 0.72 0.82 0.76 0.77120 0.78 0.79 0.78 0.78Temperature (K)

294 294 294 294

pH 1.0 1.0 1.0 1.0

The graph of % absorbency against time for this concentration of potassium iodide is shown overleaf

Implementing 33

Page 78: Hydrogen Peroxide Iodine Clock

For the next set of experiments, the following volumes of solutions were used:

Potassium iodide: 2.00 cm3

Hydrogen peroxide: 2.00 cm3

Sulphuric acid: 4.00 cm3

Starch solution: 2.00 cm3

Sodium thiosulphate: 1.00 cm3

Distilled water: 3.00 cm3

The concentration of potassium iodide used for this experiment was 0.0286 mol dm-3

% AbsorbencyTime (seconds) Run 1 Run 2 Run 3 Average

absorbency10 0.06 0.07 0.04 0.0620 0.11 0.12 0.09 0.1130 0.17 0.18 0.18 0.1840 0.22 0.24 0.18 0.2150 0.27 0.29 0.22 0.2660 0.32 0.34 0.36 0.3470 0.37 0.39 0.38 0.3880 0.41 0.44 0.44 0.4390 0.46 0.49 0.48 0.48100 0.50 0.53 0.52 0.52110 0.54 0.57 0.54 0.55120 0.58 0.62 0.59 0.60Temperature (K)

294 294 294 294

pH 1.0 1.0 1.0 1.0

The graph of % absorbency against time for this concentration of potassium iodide is shown overleaf

Implementing 35

Page 79: Hydrogen Peroxide Iodine Clock

For the next set of experiments, the following volumes of solutions were used:

Potassium iodide: 1.00 cm3

Hydrogen peroxide: 2.00 cm3

Sulphuric acid: 4.00 cm3

Starch solution: 2.00 cm3

Sodium thiosulphate: 1.00 cm3

Distilled water: 4.00 cm3

The concentration of potassium iodide used for this experiment was 0.0143 mol dm-3

% AbsorbencyTime (seconds) Run 1 Run 2 Run 3 Average

absorbency10 0.02 0.02 0.03 0.0220 0.03 0.05 0.05 0.0430 0.05 0.07 0.08 0.0740 0.07 0.10 0.11 0.0950 0.09 0.09 0.14 0.1160 0.13 0.14 0.12 0.1370 0.16 0.15 0.15 0.1580 0.17 0.18 0.17 0.1790 0.19 0.20 0.18 0.19100 0.20 0.20 0.21 0.20110 0.22 0.20 0.20 0.21120 0.22 0.23 0.21 0.22Temperature (K)

294 294 294 294

pH 1.0 1.0 1.0 1.0

The graph of % absorbency against time for this concentration of potassium iodide is shown overleaf.

Implementing 37

Page 80: Hydrogen Peroxide Iodine Clock

Changing the concentration of H+ ions

For the first set of experiments, the following volumes of solutions were used:

Sulphuric acid: 5.00 cm3

Hydrogen peroxide: 3.00 cm3

Potassium iodide: 3.00 cm3

Starch solution: 2.00 cm3

Sodium thiosulphate: 1.00 cm3

The concentration of sulphuric acid used for this experiment was 0.179 mol dm-3

% AbsorbencyTime (seconds) Run 1 Run 2 Run 3 Average

absorbency10 0.10 0.10 0.11 0.1020 0.18 0.17 0.18 0.1830 0.29 0.28 0.28 0.2840 0.37 0.38 0.36 0.3750 0.48 0.48 0.48 0.4860 0.55 0.55 0.56 0.5570 0.63 0.62 0.63 0.6380 0.70 0.69 0.70 0.7090 0.76 0.76 0.77 0.76100 0.82 0.82 0.83 0.82110 0.87 0.86 0.85 0.86120 0.90 0.89 0.89 0.89Temperature (K)

294 294 294 294

PH 1.0 1.0 1.0 1.0

The graph of % absorbency against time for this concentration of H+ ions is shown overleaf

Implementing 39

Page 81: Hydrogen Peroxide Iodine Clock

For the next set of experiments, the following volumes of solutions were used:

Sulphuric acid: 4.00 cm3

Hydrogen peroxide: 3.00 cm3

Potassium iodide: 3.00 cm3

Starch solution: 2.00 cm3

Sodium thiosulphate: 1.00 cm3

Distilled water: 1.00 cm3

The concentration of sulphuric acid used for this experiment was 0.143 mol dm-3

% AbsorbencyTime (seconds) Run 1 Run 2 Run 3 Average

absorbency10 0.08 0.07 0.08 0.0820 0.16 0.18 0.17 0.1730 0.25 0.26 0.25 0.2540 0.34 0.33 0.33 0.3350 0.41 0.41 0.42 0.4160 0.50 0.50 0.49 0.5070 0.58 0.58 0.57 0.5880 0.67 0.67 0.66 0.6790 0.73 0.73 0.72 0.73100 0.78 0.78 0.79 0.78110 0.81 0.81 0.82 0.81120 0.84 0.84 0.83 0.84Temperature (K)

294 294 294 294

pH 1.0 1.0 1.0 1.0

Implementing 41

Page 82: Hydrogen Peroxide Iodine Clock

When I initially conducted this experiment, I got the following anomalous result:

Time (seconds) Result10 0.1120 0.1930 0.2640 0.3450 0.4160 0.4870 0.5580 0.6190 0.67100 0.73110 0.78120 0.83Temperature (K)

294

pH 1.0

As this does not fit the general trend of my results, I put this anomalous result down to the fact that I had to move the solutions quickly, introducing errors into the experiment.

The graph of % absorbency against time for this concentration of H+ ions is shown overleaf

Implementing 42

Page 83: Hydrogen Peroxide Iodine Clock

For the next set of experiments, the following volumes of solutions were used:

Sulphuric acid: 3.00 cm3

Hydrogen peroxide: 3.00 cm3

Potassium iodide: 3.00 cm3

Starch solution: 2.00 cm3

Sodium thiosulphate: 1.00 cm3

Distilled water: 2.00 cm3

The concentration of sulphuric acid used for this experiment was 0.107 mol dm-3

% AbsorbencyTime (seconds) Run 1 Run 2 Run 3 Average

absorbency10 0.07 0.07 0.06 0.0720 0.15 0.14 0.14 0.1430 0.21 0.21 0.22 0.2140 0.28 0.29 0.28 0.2850 0.35 0.36 0.35 0.3560 0.42 0.41 0.42 0.4270 0.50 0.50 0.49 0.5080 0.56 0.57 0.56 0.5690 0.60 0.60 0.61 0.60100 0.66 0.65 0.63 0.65110 0.70 0.70 0.71 0.70120 0.72 0.73 0.73 0.73Temperature (K)

294 294 294 294

pH 1.0 1.0 1.0 1.0

The graph of % absorbency against time for this concentration of H+ ions is shown overleaf

Implementing 44

Page 84: Hydrogen Peroxide Iodine Clock

When I initially conducted this experiment, I got the following anomalous result:

Time (seconds) Result10 0.1020 0.2130 0.2940 0.3650 0.4360 0.5070 0.5380 0.5890 0.64100 0.71110 0.77120 0.82Temperature (K)

294

pH 1.0

As this does not fit the general trend of my results, I put this anomalous result down to the fact that I had to move the solutions quickly, introducing errors into the experiment.

The graph of % absorbency against time for this concentration of H+ ions is shown overleaf

Implementing 45

Page 85: Hydrogen Peroxide Iodine Clock

For the next set of experiments, the following volumes of solutions were used:

Sulphuric acid: 2.00 cm3

Hydrogen peroxide: 3.00 cm3

Potassium iodide: 3.00 cm3

Starch solution: 2.00 cm3

Sodium thiosulphate: 1.00 cm3

Distilled water: 3.00 cm3

The concentration of sulphuric acid used for this experiment was 0.0714 mol dm-3

% AbsorbencyTime (seconds) Run 1 Run 2 Run 3 Average

absorbency10 0.06 0.06 0.07 0.0620 0.12 0.13 0.12 0.1230 0.17 0.16 0.18 0.1740 0.24 0.23 0.23 0.2350 0.29 0.30 0.28 0.2960 0.35 0.35 0.34 0.3570 0.39 0.41 0.39 0.4080 0.44 0.46 0.43 0.4490 0.49 0.50 0.48 0.49100 0.53 0.55 0.53 0.54110 0.57 0.59 0.57 0.58120 0.62 0.63 0.61 0.62Temperature (K)

294 294 294 294

pH 1.0 1.0 1.0 1.0

The graph of % absorbency against time for this concentration of H+ ions is shown overleaf

Implementing 47

Page 86: Hydrogen Peroxide Iodine Clock

For the next set of experiments, the following volumes of solutions were used:

Sulphuric acid: 1.00 cm3

Hydrogen peroxide: 3.00 cm3

Potassium iodide: 3.00 cm3

Starch solution: 2.00 cm3

Sodium thiosulphate: 1.00 cm3

Distilled water: 4.00 cm3

The concentration of sulphuric acid used for this experiment was 0.0357 mol dm-3

% AbsorbencyTime (seconds) Run 1 Run 2 Run 3 Average

absorbency10 0.04 0.05 0.04 0.0420 0.09 0.10 0.10 0.1030 0.13 0.14 0.14 0.1440 0.16 0.18 0.19 0.1850 0.21 0.22 0.23 0.2260 0.25 0.26 0.27 0.2670 0.29 0.30 0.31 0.3080 0.32 0.34 0.35 0.3490 0.36 0.38 0.39 0.38100 0.39 0.41 0.43 0.41110 0.44 0.46 0.42 0.44120 0.45 0.48 0.46 0.46Temperature (K)

294 294 294 294

pH 1.0 1.0 1.0 1.0

The graph of % absorbency against time for this concentration of H+ ions is shown overleaf.

Implementing 49

Page 87: Hydrogen Peroxide Iodine Clock

Calibration curve

The colorimeter was tested using different known concentrations of iodine at specific increments. The percentage absorbency was then measured for each. The results are below:

Percentage absorbency (%)Concentration (mol dm-3)

Run 1 Run 2 Run 3 Average

0.00010 0.04 0.06 0.06 0.050.00020 0.10 0.11 0.10 0.100.00030 0.15 0.15 0.15 0.150.00040 0.21 0.20 0.20 0.200.00050 0.27 0.26 0.24 0.260.00060 0.30 0.29 0.32 0.300.00070 0.38 0.34 0.34 0.350.00080 0.44 0.41 0.39 0.410.00090 0.49 0.48 0.47 0.480.0010 0.52 0.50 0.49 0.500.0011 0.56 0.57 0.55 0.560.0012 0.63 0.60 0.60 0.610.0013 0.65 0.64 0.63 0.640.0014 0.70 0.70 0.69 0.700.0015 0.76 0.81 0.75 0.77

The graph for this is shown overleaf.

Implementing 49

Page 88: Hydrogen Peroxide Iodine Clock

Extension task: 1 - Varying the temperature

An experiment was conducted by varying the temperature. This enabled a graph to be drawn, and the activation enthalpy to be found using Arrhenius’ Equation. This was done using method 1 (the clock reaction). The results are shown below:

Temperature (oC) Temperature (K) Time (seconds)11.0 284.0 645.222.0 295.0 414.631.0 304.0 237.340.5 313.5 156.651.5 324.5 77.962.0 335.0 39.469.5 342.5 22.179.5 352.5 9.6

Implementing 51

Page 89: Hydrogen Peroxide Iodine Clock

Extension task: 2 - Investigating a reaction using a different filter on the colorimeter

This was done using method 2 (the colorimeter). This is to identify how one filter varies from another filter. In this case, the 490nm filter is being compared with the standard 470nm filter I used in this experiment.

In this experiment, the only variable was the volume of hydrogen peroxide and water. All other solutions were kept constant at:

Potassium iodide: 2.00 cm3

Sulphuric acid: 4.00 cm3

Starch solution: 2.00 cm3

Sodium thiosulphate: 1.00 cm3

Percentage absorbency (%)Time (seconds)

Hydrogen peroxide = 5 cm3, distilled water = 0 cm3

Hydrogen peroxide = 4 cm3, distilled water = 1 cm3

Hydrogen peroxide = 3 cm3, distilled water = 2 cm3

Hydrogen peroxide = 2 cm3, distilled water = 3 cm3

Hydrogen peroxide = 1 cm3, distilled water = 4 cm3

10 0.08 0.05 0.06 0.05 0.0220 0.16 0.13 0.13 0.09 0.0330 0.24 0.19 0.18 0.14 0.0540 0.31 0.24 0.24 0.18 0.0750 0.38 0.30 0.29 0.22 0.0860 0.44 0.35 0.34 0.25 0.0970 0.50 0.40 0.30 0.29 0.1180 0.56 0.44 0.43 0.32 0.1290 0.61 0.49 0.47 0.36 0.14100 0.66 0.53 0.52 0.39 0.15110 0.71 0.58 0.56 0.42 0.16120 0.76 0.62 0.60 0.45 0.18Temperature (K)

294 294 294 294 294

pH 1.0 1.0 1.0 1.0 1.0

Implementing 52

Page 90: Hydrogen Peroxide Iodine Clock

The graphs for each of these concentrations are shown overleaf.

Implementing 53

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Analysis

Implementing 51