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) + H2O2(aq) + 2H+(aq) → I3-(aq) + 2H2O(aq) (1)1
The half equations for this reaction can be written as follows:
3I-  I3- + 2eH2O2 + 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) + 2S2O32-(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-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.
1
http://antoine.frostburg.edu/chem/senese/101/kinetics/faq/mechanism-h2o2-iodide.shtml
http://antoine.frostburg.edu/chem/senese/101/kinetics/faq/mechanism-h2o2-iodide.shtml
3
http://antoine.frostburg.edu/chem/senese/101/kinetics/faq/mechanism-h2o2-iodide.shtml
2
Planning
1
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 reaction
Surface 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,
Reactant 1
Reactant 2
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 reaction
4
en.wikipedia.org/wiki/Starch
www.s-cool.co.uk
6
http://www.chemguide.co.uk/physical/basicrates/orders.html, Jim Clark, 2002
7
www.purchon.com/chemistry/rates.htm
5
Planning
2
For 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 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 reaction
A 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.
Planning
3
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 MaxwellBoltzmann Distribution graph should look like this:
Now all these particles have
enough energy to react as well
A catalyst does this, and this can be shown on an enthalpy profile diagram:
Uncatalysed activation
enthalpy
Catalysed activation
enthalpy
Diagram obtained from
http://cwx.prenhall.com
/bookbind/pubbooks/hill
chem3/medialib/media_
portfolio/13.html
ΔH
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
Planning
4
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 reaction
Increasing 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 = volume
n = number of moles
R = molar gas constant
T = 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 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.
Planning
5
Low pressure
High pressure
Diagram referenced to
www.chemguide.co.uk
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 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
Planning
6
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
dy/dx then gives the gradient of the curve at that particular point, or the initial rate.
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 = [S2O32-]/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
Planning
7
= 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 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 reaction
The 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.
Obtained from
www.webchem.net
Planning
8
The area under the graph represents the number of molecules present. Some interesting
points to note from this distribution curve 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
w
l
c
u
v
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 + u22 +… + uN2)
l
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 + u22 +… + uN2)
l3
If <u2> represents the mean value of the squares of all the velocities,
<u2> = (u12 + u22 +… + uN2)/N
8
9
Chemistry: AS level and A level (International), Ratcliffe, Eccles, Raffan, Nicholson and Johnson, 2004
Oxford: Succes at AQA Physics B A2, Ken Price and Gerard Kelly, 2001
Planning
9
and
N<u2> = u12 + u22 +… + uN2
Therefore, p = Nm<u2>/l3
For any molecule, pythagoras’ theorem can be applied to give:
c 2 = u2 + v 2 + w 2
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
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:
Planning 10
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’ equation
As 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)
c = ln(A)
x = 1/T
m = -Ea/R
Therefore, the activation enthalpy can be found by:
Gradient = -Ea/R
 -Gradient x R = Ea
10
http://www.kobold.demon.co.uk/kinetics/collisio.htm
Planning 11
How can the rate be calculated?
Measuring volume
The 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 pressure
When 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 analysis
This 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 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 analysis
This 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
11
Facts and Practice for A-level: Chemistry, Max Parsonage, 2001
Planning 12
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 method
This 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 orders
The rate orders that I can expect for my results are one of three; Zero order, first order or
second order.
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.
Graph obtained from
www.chem.purdue.edu
Half-life
decreases with
decreasing
concentration
Planning 13
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.
Initial
rate of
Reaction
Concentration of reactant
Planning 14
First Order
The 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.
Graph obtained
from
wps.prenhall.com
Half-life
decreases with
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.
Initial rate
of Reaction
Concentration of reactant
Planning 15
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 16
Initial
rate of
Reaction
Concentration of reactant
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.
Structural formula obtained from:
http://en.wikipedia.org/wiki/Hydrogen_peroxide
12
13
Hydrogen Peroxide, in Kirk-Othmer Encyclopaedia of Chemical Technology
http://en.wikipedia.org/wiki/Hydrogen_peroxide
Planning 17
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.
The structural formula for potassium iodide is shown below:
Structural formula obtained
from:
http://www.webelements.com/w
ebelements/compounds/text/K/I
1K1-7681110.html
.
14
Dictionary of Inorganic Compounds, J.E. Macintyre, 1992
Planning 18
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) + 2S2O32-(aq)  3I-(aq) + S4O6-(aq).
The structural formula for sodium thiosulphate is as follows:
Structural formula obtained from:
http://msds.pcd.go.th/images/Form
ula_Chain/10102-17-7.gif
Experiment
Variables
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
Planning 19
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
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 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.
 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 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
Planning 20
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 thermometer
If 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.
 Use of a volumetric flask
It 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.
Planning 21
 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
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.
This is the meniscus. You
must always measure
volumes from below the
meniscus.
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
Planning 22
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 solutions
For 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,
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.
Planning 23
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,
= 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 a cool place before use, as the starch-iodine complex becomes
unstable above 35oC.
Planning 24
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-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 assessment
Before commencement of any experiment, the potential risks involved must be considered
and eliminated in the experiment (or minimised as much as possible).
Planning 25
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
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.
Planning 26
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:
If splashed in
eye:
If spilt on skin
or clothes:
If spilt in
laboratory:
Wash out mouth and give a glass of water. Do not induce vomiting. Seek
medical attention as soon as possible
Flood the eye with gently running tap water for 10 minutes. Seek medical
attention.
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.
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:
If vapour
inhaled:
If vapour
affects eyes:
If solid gets in
eyes:
If spilt on skin
or clothes:
If spilt in
laboratory:
Wash out mouth and give a glass of water. Seek medical attention as soon as
possible
Remove victim to fresh air. Seek medical attention if breathing is slightly
affected.
Bathe eyes. If discomfort persists, seek medical attention.
Flood the eye with gently running tap water for 10 minutes. Seek medical
attention.
Brush off solid immediately. Flood affected area with water immediately, or
bathe in sodium thiosulphate solution. Seek medical attention if blistering
occurs.
Wear eye protection and gloves. Ventilate area of spill. Spread sodium
thiosulphate over area of spill, leave for an hour, and then mop up and
rinse.
Planning 27
Hydrogen peroxide:
If swallowed:
If liquid gets in
eyes:
If spilt on skin
or clothes:
If spilt in
laboratory:
Wash out mouth and give a glass of water. Seek medical attention as soon as
possible
Flood the eye with gently running tap water for 10 minutes. Seek medical
attention.
Flood affected area with water immediately. Seek medical attention if
blistering occurs.
Wear eye protection and gloves. Cover with mineral absorbent and clear up
into a bucket. Rinse several times. Add water to dilute at least ten times
before washing the liquid down the foul-water drain.
Sodium thiosulphate and potassium iodide:
If swallowed:
If liquid gets in
eyes:
If spilt on skin
or clothes:
If spilt in
laboratory:
Give plenty of water. Seek medical attention as soon as possible
Flood the eye with gently running tap water for 10 minutes. Seek medical
attention.
Flood affected area with water immediately. Seek medical attention if
blistering occurs. Wash off skin with plenty of water.
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 experiment
It is necessary to know what apparatus is being used before doing the experiment. The
following apparatus will be used for this experiment:
 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
Planning 28
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 rack
Boiling tubes can be stored in the test tube rack to prevent accidents
 A colorimeter
A 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 tubes
These 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
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 bath
A 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 flask
1 x 0.500 dm3 flasks
The 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
Planning 29
obtained would be flawed, and thus, unreliable. The use of separate volumetric flasks
ensures a fair test.
 5 beakers
To 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.
 Stopwatch
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.
 2 cuvettes
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.
Test tube
containing the
solutions
15
Shakhashiri's Chemical Demonstrations, Vol. 4, pages 42-43
Planning 30
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 rack
Boiling 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 tubes
These 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
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 beakers
To 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
Planning 31
occur. If contamination does occur, then the results obtained would be flawed, and thus,
unreliable.
 Stopwatch
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.
Method
1. 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 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.
Planning 32
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. 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
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.
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. 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
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.
Planning 33
11. These would then be plotted against concentration, allowing the order of reaction to be
found.
Method 2
Cuvette containing
solution
Colorimeter
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 rack
Boiling 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 34
 A colorimeter
A 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 water
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.
 Stopwatch
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.
 2 cuvettes
For 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.
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.
Planning 35
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.
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.
Planning 36
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
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 points
For 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 curve
A 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
Planning 37
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 dm3
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.
2.
3.
4.
5.
6.
7.
Measure out 3.00g of potassium iodide and place into a beaker
Add about 65.0 cm3 of water
Measure out 2.54g of iodine granules and add this to the solution in the beaker.
Stir using a magnetic stirrer for several minutes.
Pour the solution into a 100 cm3 volumetric flask and make up to 100 cm3
This produces a 0.1 mol dm-3 solution of iodine.
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 colorimeter
3. Make up the required solution as table 8 (please refer to appendices) in a 50 cm3 beaker
4. Pipette some into a cuvette
5. Measure the absorbency and record into a table (Appendix 10)
6. Draw a graph of absorbency against concentration
Extension task - 1
For 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-bystep 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 - 2
For 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 38
Synoptic table
Concept
Unit met
Module
Year
Rate orders
Redox reactions (reduction
and oxidation)
Absorption Spectrum
Half-equations
Rate-concentration graphs
Rate-determining step
Rate-mechanisms
Half-life
Intermediates
Reactions incorporating
colour changes
Ionic lattices
Engineering proteins
From Minerals to
Elements
Elements of Life
Steel Story
Engineering Proteins
Engineering Proteins
Engineering Proteins
Engineering Proteins
The Atmosphere
Steel Story
2849
2848
A2
AS
2850
2849
2849
2849
2849
2849
2848
2849
AS
A2
A2
A2
A2
A2
AS
A2
From Minerals to
Elements
Minerals to elements
2848
AS
2848
AS
Elements of life
Engineering proteins
Engineering proteins
The atmosphere
The atmosphere
The atmosphere
2850
2849
2849
2848
2848
2848
AS
A2
A2
AS
AS
AS
Developing fuels and
the atmosphere
Polymer revolution
Elements of Life
2850 + 2848
AS
2848
2850
AS
AS
Steel Story
Aspects of
Agriculture
2849
2854
A2
A2
Aspects of agriculture
2854
A2
Affect of concentrations on
rate
The mole
Kinetic theory
Boltzmann Distribution
Affect of pressure on rate
Factors affecting rate
Affect of temperature on
rate
Affect of catalysts on rate
Van der Waals forces
Making up a standard
solution
Using colorimeter
Measuring the time taken
for a reaction to reach a
particular stage
Arrhenius equation
Planning 39
Preliminary Experiment
To 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)
1
2nd
3rd
Average
60.67
59.79 58.71 59.72
1.0
1.0
1.0
1.0
294
294
294
294
st
pH
Temperature (K)
Planning 40
Changes after the Experiment
Further 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 peroxide
The 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 iodide
Furthermore, 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 volumetric flask. This solution will
then be made up to 250cm3 using distilled water.
Planning 41
Implementing - Method 1
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
Starch solution: 2.00 cm3
Sulphuric acid: 4.00 cm3
Sodium thiosulphate: 1.00 cm3
pH
Temperature (K)
1st
53.13
1.0
294.0
Time for reaction to happen (seconds)
2nd
3rd
Average
56.12
58.59
55.95
1.0
1.0
1.0
294.0
294.0
294.0
When I initially conducted this experiment, I got the following anomalous results:
Time for reaction to happen (seconds)
1
2nd
3rd
Average
59.15
63.21
62.63
61.66
1.0
1.0
1.0
1.0
294.0
294.0
294.0
294.0
st
pH
Temperature (K)
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
Sulphuric acid: 4.00 cm3
Sodium thiosulphate: 1.00 cm3
Water: 1.00 cm3
pH
Temperature (K)
1st
70.68
1.0
294.0
Time for reaction to happen (seconds)
2nd
3rd
Average
72.25
76.81
73.25
1.0
1.0
1.0
294.0
294.0
294.0
Implementing
1
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
pH
Temperature (K)
1st
92.22
1.0
294.0
Time for reaction to happen (seconds)
2nd
3rd
Average
93.22
92.09
92.51
1.0
1.0
1.0
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)
1
2nd
3rd
Average
141.25
142.94
139.15
141.11
1.0
1.0
1.0
1.0
294.0
294.0
294.0
294.0
st
pH
Temperature (K)
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
Sodium thiosulphate: 1.00 cm3
Water: 4.00 cm3
Time for reaction to happen (seconds)
1
2nd
3rd
Average
297.31
301.14
300.17
299.54
1.0
1.0
1.0
1.0
294.0
294.0
294.0
294.0
st
pH
Temperature (K)
Implementing
2
The table for the average set of results for each concentration for hydrogen peroxide is
below:
Concentration of hydrogen peroxide (mol
dm-3)
0.357
0.286
0.214
0.143
0.0714
Average time taken (seconds)
55.95
73.25
92.51
141.11
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) + 2S2O32-(aq)  3I-(aq) + S4O6-(aq)
The stoichiometry shows that, for every one mole of iodine produced, two moles of S2O32ions 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.
Using this equation and the above information, the following table can be constructed:
Concentration of
hydrogen peroxide
(mol dm-3)
0.0357
0.0286
0.0214
0.0143
0.00714
Average time taken
(seconds)
Calculation for rate
Rate (mol s-1)
55.95
73.25
92.51
141.11
299.54
0.0025/55.95
0.0025/73.25
0.0025/92.51
0.0025/141.11
0.0025/299.54
4.47 x 10-5
3.41 x 10-5
2.70 x 10-5
1.77 x 10-5
8.35 x 10-6
The graphs for hydrogen peroxide are shown overleaf.
Implementing
3
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
Starch solution: 2.00 cm3
Sulphuric acid: 4.00 cm3
Sodium thiosulphate: 1.00 cm3
pH
Temperature (K)
1st
64.75
1.0
294.0
Time for reaction to happen (seconds)
2nd
3rd
Average
67.37
66.10
66.07
1.0
1.0
1.0
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)
1
2nd
3rd
Average
80.38
81.97
83.56
81.97
1.0
1.0
1.0
1.0
294.0
294.0
294.0
294.0
st
PH
Temperature (K)
Implementing
6
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
PH
Temperature (K)
Time for reaction to happen (seconds)
1st
2nd
3rd
Average
108.16
108.63
108.01
108.27
1.0
1.0
1.0
1.0
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)
1
2nd
3rd
Average
143.32
143.02
141.71
142.68
1.0
1.0
1.0
1.0
294.0
294.0
294.0
294.0
st
PH
Temperature (K)
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
Sodium thiosulphate: 1.00 cm3
Water: 4.00 cm3
Time for reaction to happen (seconds)
1
2nd
3rd
Average
303.94
308.87
297.06
303.29
1.0
1.0
1.0
1.0
294.0
294.0
294.0
294.0
st
pH
Temperature (K)
Implementing
7
The table for the average set of results for each concentration for hydrogen peroxide is
below:
Concentration of potassium iodide (mol
dm-3)
0.0714
0.0571
0.0429
0.0286
0.0143
Average time taken (seconds)
66.07
81.97
108.27
142.68
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) + 2S2O32-(aq)  3I-(aq) + S4O6-(aq)
The stoichiometry shows that, for every one mole of iodine produced, two moles of S2O32ions 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.
Using this equation and the above information, the following table can be constructed:
Concentration of
potassium iodide
(mol dm-3)
0.0714
0.0571
0.0429
0.0286
0.0143
Average time taken
(seconds)
Calculation for rate
Rate (mol s-1)
66.07
81.97
108.27
142.68
303.29
0.0025/55.95
0.0025/73.25
0.0025/92.51
0.0025/141.11
0.0025/299.54
3.78 x 10-5
3.05 x 10-5
2.31 x 10-5
1.75 x 10-5
8.24 x 10-6
The graphs for potassium iodide are shown overleaf.
Implementing
8
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
pH
Temperature (K)
1st
95.15
1.0
294.0
Time for reaction to happen (seconds)
2nd
3rd
Average
96.37
96.93
96.15
1.0
1.0
1.0
294.0
294.0
294.0
When I initially conducted this experiment, I got the following anomalous results:
pH
Temperature (K)
1st
65.03
1.0
295.0
Time for reaction to happen (seconds)
2nd
3rd
Average
63.21
62.63
63.62
1.0
1.0
1.0
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
PH
Temperature (K)
Time for reaction to happen (seconds)
1st
2nd
3rd
Average
108.29
106.00
109.27
107.85
1.0
1.0
1.0
1.0
294.0
294.0
294.0
294.0
Implementing 11
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
PH
Temperature (K)
Time for reaction to happen (seconds)
1st
2nd
3rd
Average
121.32
122.49
119.21
121.01
1.0
1.0
1.0
1.0
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)
1
2nd
3rd
Average
137.78
136.08
138.96
137.61
1.0
1.0
1.0
1.0
294.0
294.0
294.0
294.0
st
PH
Temperature (K)
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
Water: 4.00 cm3
Time for reaction to happen (seconds)
1
2nd
3rd
Average
172.04
172.19
173.00
172.41
1.0
1.0
1.0
1.0
294.0
294.0
294.0
294.0
st
pH
Temperature (K)
Implementing 12
When I first conducted this experiment, I obtained these anomalies:
Time for reaction to happen (seconds)
1
2nd
3rd
Average
208.22
208.38
210.64
209.08
1.0
1.0
1.0
1.0
294.5
293.0
292.0
293.0
st
pH
Temperature (K)
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)
0.179
0.143
0.107
0.0714
0.0357
Average time taken (seconds)
96.15
107.85
121.01
137.61
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
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) + 2S2O32-(aq)  3I-(aq) + S4O6-(aq)
The stoichiometry shows that, for every one mole of iodine produced, two moles of S2O32ions 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 13
Using this equation and the above information, the following table can be constructed:
Concentration of
sulphuric acid (mol
dm-3)
0.179
0.143
0.107
0.0714
0.0357
Average time taken
(seconds)
Calculation for rate
Rate (mol s-1)
96.15
107.85
121.01
137.61
172.41
0.0025/96.15
0.0025/107.85
0.0025/121.01
0.0025/137.61
0.0025/172.41
2.60 x 10-5
2.32 x 10-5
2.07 x 10-5
1.82 x 10-5
1.45 x 10-5
The graphs for H+ ions are shown overleaf.
Implementing 14
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
% Absorbency
Run 2
Run 3
0.14
0.32
0.40
0.58
0.73
0.89
1.01
1.10
1.24
1.28
1.32
1.35
294
0.14
0.27
0.41
0.58
0.72
0.87
1.01
1.18
1.22
1.29
1.30
1.35
294
0.13
0.28
0.39
0.57
0.74
0.88
1.02
1.17
1.23
1.28
1.31
1.35
294
Average
absorbency
0.14
0.29
0.40
0.58
0.73
0.88
1.01
1.15
1.23
1.28
1.31
1.35
294
1.0
1.0
1.0
1.0
Time (seconds) Run 1
10
20
30
40
50
60
70
80
90
100
110
120
Temperature
(K)
pH
Implementing 17
When I initially conducted this experiment, I got the following anomalous result:
Time (seconds)
10
20
30
40
50
60
70
80
90
100
110
120
Temperature
(K)
pH
Result
0.18
0.42
0.56
0.78
0.99
1.10
1.21
1.31
1.41
1.50
1.58
1.57
294
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
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
% Absorbency
Run 2
Run 3
0.11
0.22
0.35
0.45
0.56
0.72
0.81
0.91
0.98
1.03
1.07
1.12
294
0.11
0.21
0.34
0.48
0.60
0.71
0.81
0.91
1.00
1.03
1.08
1.14
294
0.12
0.26
0.35
0.45
0.56
0.70
0.82
0.92
0.99
1.02
1.07
1.10
294
Average
absorbency
0.11
0.23
0.35
0.46
0.57
0.70
0.81
0.91
0.99
1.03
1.07
1.12
294
1.0
1.0
1.0
1.0
Time (seconds) Run 1
10
20
30
40
50
60
70
80
90
100
110
120
Temperature
(K)
pH
The graph of % absorbency against time for this concentration of hydrogen peroxide is
shown overleaf
Implementing 20
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
% Absorbency
Run 2
Run 3
0.13
0.21
0.29
0.37
0.45
0.52
0.60
0.66
0.72
0.78
0.84
0.90
294
0.12
0.21
0.29
0.39
0.47
0.55
0.63
0.69
0.76
0.83
0.89
0.95
294
0.11
0.19
0.27
0.33
0.43
0.50
0.58
0.64
0.70
0.76
0.82
0.88
294
Average
absorbency
0.12
0.20
0.28
0.36
0.45
0.52
0.60
0.66
0.73
0.79
0.85
0.91
294
1.0
1.0
1.0
1.0
Time (seconds) Run 1
10
20
30
40
50
60
70
80
90
100
110
120
Temperature
(K)
pH
The graph of % absorbency against time for this concentration of hydrogen peroxide is
shown overleaf
Implementing 22
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
% Absorbency
Run 2
Run 3
0.08
0.17
0.26
0.36
0.43
0.51
0.61
0.68
0.77
0.82
0.85
0.88
294
0.1
0.18
0.25
0.35
0.46
0.52
0.6
0.68
0.73
0.81
0.85
0.9
294
0.09
0.17
0.25
0.35
0.41
0.5
0.59
0.67
0.75
0.81
0.84
0.89
294
Average
absorbency
0.09
0.17
0.25
0.35
0.43
0.51
0.6
0.68
0.75
0.81
0.85
0.89
294
1.0
1.0
1.0
1.0
Time (seconds) Run 1
10
20
30
40
50
60
70
80
90
100
110
120
Temperature
(K)
pH
The graph of % absorbency against time for this concentration of hydrogen peroxide is
shown overleaf
Implementing 24
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
% Absorbency
Run 2
Run 3
0.06
0.13
0.18
0.24
0.29
0.35
0.40
0.47
0.52
0.54
0.59
0.63
294
0.06
0.11
0.17
0.23
0.30
0.35
0.42
0.48
0.50
0.55
0.60
0.62
294
0.07
0.12
0.19
0.22
0.30
0.36
0.41
0.46
0.51
0.56
0.58
0.64
294
Average
absorbency
0.06
0.12
0.18
0.23
0.30
0.35
0.41
0.47
0.51
0.54
0.59
0.63
294
1.0
1.0
1.0
1.0
Time (seconds) Run 1
10
20
30
40
50
60
70
80
90
100
110
120
Temperature
(K)
pH
The graph of % absorbency against time for this concentration of hydrogen peroxide is
shown overleaf.
Implementing 26
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
% Absorbency
Run 2
Run 3
0.16
0.28
0.43
0.54
0.65
0.75
0.85
0.93
1.01
1.09
1.12
1.17
294
0.14
0.29
0.42
0.55
0.66
0.76
0.88
1.00
0.98
1.08
1.13
1.18
294
0.15
0.27
0.38
0.48
0.68
0.76
0.79
0.90
0.99
1.07
1.12
1.16
294
Average
absorbency
0.15
0.28
0.41
0.52
0.66
0.76
0.84
0.94
0.99
1.08
1.12
1.17
294
1.0
1.0
1.0
1.0
Time (seconds) Run 1
10
20
30
40
50
60
70
80
90
100
110
120
Temperature
(K)
PH
The graph of % absorbency against time for this concentration of potassium iodide is shown
overleaf
Implementing 28
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
% Absorbency
Run 2
Run 3
0.13
0.24
0.34
0.44
0.52
0.61
0.71
0.81
0.90
0.90
0.96
1.00
294
0.16
0.29
0.34
0.48
0.56
0.68
0.74
0.81
0.90
0.96
0.98
1.01
294
0.12
0.20
0.29
0.45
0.48
0.63
0.69
0.82
0.89
0.99
0.97
1.00
294
Average
absorbency
0.14
0.24
0.32
0.46
0.52
0.64
0.71
0.81
0.90
0.95
0.97
1.00
294
1.0
1.0
1.0
1.0
Time (seconds) Run 1
10
20
30
40
50
60
70
80
90
100
110
120
Temperature
(K)
pH
Implementing 30
When I initially conducted this experiment, I got the following anomalous result:
Time (seconds)
10
20
30
40
50
60
70
80
90
100
110
120
Temperature
(K)
pH
Result
0.17
0.35
0.57
0.84
1.01
1.19
1.34
1.45
1.54
1.60
1.65
1.69
294
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
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
% Absorbency
Run 2
Run 3
0.11
0.19
0.26
0.33
0.41
0.48
0.54
0.61
0.67
0.73
0.72
0.78
294
0.10
0.20
0.29
0.34
0.38
0.47
0.55
0.62
0.69
0.76
0.82
0.79
294
0.08
0.15
0.21
0.28
0.44
0.40
0.54
0.61
0.68
0.71
0.76
0.78
294
Average
absorbency
0.10
0.18
0.25
0.32
0.41
0.45
0.54
0.61
0.68
0.73
0.77
0.78
294
1.0
1.0
1.0
1.0
Time (seconds) Run 1
10
20
30
40
50
60
70
80
90
100
110
120
Temperature
(K)
pH
The graph of % absorbency against time for this concentration of potassium iodide is shown
overleaf
Implementing 33
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
% Absorbency
Run 2
Run 3
0.06
0.11
0.17
0.22
0.27
0.32
0.37
0.41
0.46
0.50
0.54
0.58
294
0.07
0.12
0.18
0.24
0.29
0.34
0.39
0.44
0.49
0.53
0.57
0.62
294
0.04
0.09
0.18
0.18
0.22
0.36
0.38
0.44
0.48
0.52
0.54
0.59
294
Average
absorbency
0.06
0.11
0.18
0.21
0.26
0.34
0.38
0.43
0.48
0.52
0.55
0.60
294
1.0
1.0
1.0
1.0
Time (seconds) Run 1
10
20
30
40
50
60
70
80
90
100
110
120
Temperature
(K)
pH
The graph of % absorbency against time for this concentration of potassium iodide is shown
overleaf
Implementing 35
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
% Absorbency
Run 2
Run 3
0.02
0.03
0.05
0.07
0.09
0.13
0.16
0.17
0.19
0.20
0.22
0.22
294
0.02
0.05
0.07
0.10
0.09
0.14
0.15
0.18
0.20
0.20
0.20
0.23
294
0.03
0.05
0.08
0.11
0.14
0.12
0.15
0.17
0.18
0.21
0.20
0.21
294
Average
absorbency
0.02
0.04
0.07
0.09
0.11
0.13
0.15
0.17
0.19
0.20
0.21
0.22
294
1.0
1.0
1.0
1.0
Time (seconds) Run 1
10
20
30
40
50
60
70
80
90
100
110
120
Temperature
(K)
pH
The graph of % absorbency against time for this concentration of potassium iodide is shown
overleaf.
Implementing 37
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
% Absorbency
Run 2
Run 3
0.10
0.18
0.29
0.37
0.48
0.55
0.63
0.70
0.76
0.82
0.87
0.90
294
0.10
0.17
0.28
0.38
0.48
0.55
0.62
0.69
0.76
0.82
0.86
0.89
294
0.11
0.18
0.28
0.36
0.48
0.56
0.63
0.70
0.77
0.83
0.85
0.89
294
Average
absorbency
0.10
0.18
0.28
0.37
0.48
0.55
0.63
0.70
0.76
0.82
0.86
0.89
294
1.0
1.0
1.0
1.0
Time (seconds) Run 1
10
20
30
40
50
60
70
80
90
100
110
120
Temperature
(K)
PH
The graph of % absorbency against time for this concentration of H+ ions is shown overleaf
Implementing 39
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
% Absorbency
Run 2
Run 3
0.08
0.16
0.25
0.34
0.41
0.50
0.58
0.67
0.73
0.78
0.81
0.84
294
0.07
0.18
0.26
0.33
0.41
0.50
0.58
0.67
0.73
0.78
0.81
0.84
294
0.08
0.17
0.25
0.33
0.42
0.49
0.57
0.66
0.72
0.79
0.82
0.83
294
Average
absorbency
0.08
0.17
0.25
0.33
0.41
0.50
0.58
0.67
0.73
0.78
0.81
0.84
294
1.0
1.0
1.0
1.0
Time (seconds) Run 1
10
20
30
40
50
60
70
80
90
100
110
120
Temperature
(K)
pH
Implementing 41
When I initially conducted this experiment, I got the following anomalous result:
Time (seconds)
10
20
30
40
50
60
70
80
90
100
110
120
Temperature
(K)
pH
Result
0.11
0.19
0.26
0.34
0.41
0.48
0.55
0.61
0.67
0.73
0.78
0.83
294
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
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
% Absorbency
Run 2
Run 3
0.07
0.15
0.21
0.28
0.35
0.42
0.50
0.56
0.60
0.66
0.70
0.72
294
0.07
0.14
0.21
0.29
0.36
0.41
0.50
0.57
0.60
0.65
0.70
0.73
294
0.06
0.14
0.22
0.28
0.35
0.42
0.49
0.56
0.61
0.63
0.71
0.73
294
Average
absorbency
0.07
0.14
0.21
0.28
0.35
0.42
0.50
0.56
0.60
0.65
0.70
0.73
294
1.0
1.0
1.0
1.0
Time (seconds) Run 1
10
20
30
40
50
60
70
80
90
100
110
120
Temperature
(K)
pH
The graph of % absorbency against time for this concentration of H+ ions is shown overleaf
Implementing 44
When I initially conducted this experiment, I got the following anomalous result:
Time (seconds)
10
20
30
40
50
60
70
80
90
100
110
120
Temperature
(K)
pH
Result
0.10
0.21
0.29
0.36
0.43
0.50
0.53
0.58
0.64
0.71
0.77
0.82
294
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
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
% Absorbency
Run 2
Run 3
0.06
0.12
0.17
0.24
0.29
0.35
0.39
0.44
0.49
0.53
0.57
0.62
294
0.06
0.13
0.16
0.23
0.30
0.35
0.41
0.46
0.50
0.55
0.59
0.63
294
0.07
0.12
0.18
0.23
0.28
0.34
0.39
0.43
0.48
0.53
0.57
0.61
294
Average
absorbency
0.06
0.12
0.17
0.23
0.29
0.35
0.40
0.44
0.49
0.54
0.58
0.62
294
1.0
1.0
1.0
1.0
Time (seconds) Run 1
10
20
30
40
50
60
70
80
90
100
110
120
Temperature
(K)
pH
The graph of % absorbency against time for this concentration of H+ ions is shown overleaf
Implementing 47
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
% Absorbency
Run 2
Run 3
0.04
0.09
0.13
0.16
0.21
0.25
0.29
0.32
0.36
0.39
0.44
0.45
294
0.05
0.10
0.14
0.18
0.22
0.26
0.30
0.34
0.38
0.41
0.46
0.48
294
0.04
0.10
0.14
0.19
0.23
0.27
0.31
0.35
0.39
0.43
0.42
0.46
294
Average
absorbency
0.04
0.10
0.14
0.18
0.22
0.26
0.30
0.34
0.38
0.41
0.44
0.46
294
1.0
1.0
1.0
1.0
Time (seconds) Run 1
10
20
30
40
50
60
70
80
90
100
110
120
Temperature
(K)
pH
The graph of % absorbency against time for this concentration of H+ ions is shown overleaf.
Implementing 49
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:
Concentration
(mol dm-3)
0.00010
0.00020
0.00030
0.00040
0.00050
0.00060
0.00070
0.00080
0.00090
0.0010
0.0011
0.0012
0.0013
0.0014
0.0015
Run 1
Percentage absorbency (%)
Run 2
Run 3
Average
0.04
0.10
0.15
0.21
0.27
0.30
0.38
0.44
0.49
0.52
0.56
0.63
0.65
0.70
0.76
0.06
0.11
0.15
0.20
0.26
0.29
0.34
0.41
0.48
0.50
0.57
0.60
0.64
0.70
0.81
0.05
0.10
0.15
0.20
0.26
0.30
0.35
0.41
0.48
0.50
0.56
0.61
0.64
0.70
0.77
0.06
0.10
0.15
0.20
0.24
0.32
0.34
0.39
0.47
0.49
0.55
0.60
0.63
0.69
0.75
The graph for this is shown overleaf.
Implementing 49
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)
11.0
22.0
31.0
40.5
51.5
62.0
69.5
79.5
Temperature (K)
284.0
295.0
304.0
313.5
324.5
335.0
342.5
352.5
Time (seconds)
645.2
414.6
237.3
156.6
77.9
39.4
22.1
9.6
Implementing 51
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
Time
(seconds)
10
20
30
40
50
60
70
80
90
100
110
120
Temperature
(K)
pH
Hydrogen
peroxide =
5 cm3,
distilled
water = 0
cm3
0.08
0.16
0.24
0.31
0.38
0.44
0.50
0.56
0.61
0.66
0.71
0.76
294
Percentage absorbency (%)
Hydrogen
Hydrogen
Hydrogen
peroxide =
peroxide =
peroxide =
4 cm3,
3 cm3,
2 cm3,
distilled
distilled
distilled
water = 1
water = 2
water = 3
3
3
cm
cm
cm3
0.05
0.06
0.05
0.13
0.13
0.09
0.19
0.18
0.14
0.24
0.24
0.18
0.30
0.29
0.22
0.35
0.34
0.25
0.40
0.30
0.29
0.44
0.43
0.32
0.49
0.47
0.36
0.53
0.52
0.39
0.58
0.56
0.42
0.62
0.60
0.45
294
294
294
Hydrogen
peroxide =
1 cm3,
distilled
water = 4
cm3
0.02
0.03
0.05
0.07
0.08
0.09
0.11
0.12
0.14
0.15
0.16
0.18
294
1.0
1.0
1.0
1.0
1.0
The graphs for each of these concentrations are shown overleaf.
Implementing 52
Analysis
Implementing 51