Chemistry Experiments (Module Kinetic Experiment)
IJSO Training: Chemistry Experiment
Experiment Title: Kinetic Experiments
Experiment 1: Investigating the Effects of Concentration and Temperature on the Rate of a
Chemical Reaction
Objective:
To investigate the concentration and temperature effects on the rate of a chemical
reaction
Introduction
You have learnt in the classroom session that the rate of a chemical reaction depends on the
concentration of the reactants as well as the reaction temperature.
We are going to investigate
this chemical phenomenon. To facilitate our investigation, we need to choose a chemical reaction
that exhibits some easily observable physical changes so that we can monitor the progress of the
reaction conveniently. In this regard we will study the reaction between sodium thiosulphate and
diluted hydrochloric acid.
S2O32-(aq) + 2 H+(aq)

SO2(g) + S(s)
+ H2O(l)
Sodium thiosulphate reacts with diluted acid to give sulphur dioxide, sulphur and water. Both
sodium thiosulphate and diluted hydrochloric acid are colorless solution. Sulphur dioxide is a
very soluble gas and dissolves completely in the aqueous solution. The sulphur formed, however,
is not soluble and will exist in the mixture as white or yellow precipitate (or colloidal). It makes
the reaction mixture becomes opaque as the reaction occurs. Therefore, we can study the reaction
rate by monitoring the opaqueness of the reaction. This can be easily done by measuring the time
taken for forming a certain amount of precipitate.
In the experiment, you will carry out the reaction by mixing the two reactants into a small
beaker, and put the beaker on top of a piece of white paper which has a mark on it. Before the
reaction starts you can see the mark clearly from the top of the beaker through the solution.
However, as the reaction proceeds, sulphur precipitate forms and makes the solution become more
opaque, and so eventually the mark is completely masked. The time taken for the mark to become
totally disappeared indicates how fast the reaction has occurred.
Chemicals:
0.15 M sodium thiosulphate solution
1 M hydrochloric acid
Apparatus:
100 mL beakers
100 mL measuring cylinder
Thermometer (alcohol, up to 110oC)
Stopwatch
Part A: To study the concentration effect on the rate of a chemical reaction.
In this part of the experiment you will study the concentrate effect on the rate of the above
reaction by varying the concentration of the sodium thiosulphate solution.
Chemistry Experiments (Module Kinetic Experiment)
Procedure:
1. Take a piece of white paper and make a clear mark on it.
2. Put a dry clean beaker on top of the mark.
3. Measure 60 mL of 0.15 M sodium thiosulphate solution and put it into the beaker.
4. Quickly add 10 mL of 1 M HCl into the beaker in a single portion and start the stopwatch
immediately.
5. Stir the mixture gently with a glass rod for a few seconds.
6. Look at the mark from vertically above the beaker through its content, and record the time
taken for the mark to become totally disappeared.
7. Repeat steps 1-6 by substituting the 60 mL 0.15 M sodium thiosulphate solution with:
Solution Vol. of 0.15M Na2S2O3
Vol. of water
Total volume
solution
A
45 mL
15 mL
60 mL
B
30 mL
30 mL
60 mL
C
15 mL
45 mL
60 mL
Data
1. Time for the disappearance of the cross
Volume of
Na2S2O3(aq)
(mL)
Volume of
water
(mL)
Volume of
HCl(aq)
(mL)
60
0
10
45
15
30
30
15
45
Initial concentration
of Na2S2O3(aq) in the
reaction mixture/M
Time for the
disappearance of the
cross/s
Trial
Trial
Average
1
2
10
10
10
Data Analysis
1. Calculate the initial concentration of sodium thiosulphate in the reaction mixtures.
Run 1
Run 2
Chemistry Experiments (Module Kinetic Experiment)
Run 3
Run 4
2.
Plot the time required for the disappearance of the cross against the initial concentration of
sodium thiosulphate on a graph paper. Label your graph properly.
3.
How is the rate of reaction affected by the concentration of the reactant?
Chemistry Experiments (Module Kinetic Experiment)
Part B: To study the temperature effect on the rate of a chemical reaction.
In this part of the experiment you will study the temperature effect on the rate of the about
reaction. You will carry out the reaction at different temperature and compare the results.
Temperature: 0oC (ice-water bath)
Room temperature (refer to the data taken in Part A)
50oC (by warming the Na2S2O3 solution with a 50oC water bath)
Procedure
Reaction at 0oC
1. Measure 60 mL of 0.15 M sodium thiosulphate solution and put it into a clean, dry 100 mL
beaker.
2. Cool the solution in an ice-water bath for 15 minutes.
3. Measure the temperature of the solution.
4. Carry out the reaction of the cooled Na2S2O3 solution with 1 M HCl by following the steps 1-6
described in part A.
5. Measure 60 mL of 0.15 M sodium thiosulphate solution and put it into another dry, clean 100
mL beaker.
6. Warm the solution in a 50oC water bath for 15 minutes.
7. Measure the temperature of the warmed solution.
8. Carry out the reaction of the warmed Na2S2O3 solution with 1 M HCl by following the steps
1-6 described in part A.
Data
Time for the disappearance of the cross
Initial concentration of
Na2S2O3(aq) in the reaction
moisture/M
Temperature of the
reaction mixture/oC
Time for the disappearance of the
cross/s
Trial 1
Trial 2
Average
Data Analysis
1. From the data in the above table, how is the rate of reaction affected by temperature?
Chemistry Experiments (Module Kinetic Experiment)
Experiment 2: Investigating the effect of surface area (particle size of a solid reactant) on the
rate of a chemical reaction.
Objective:
To investigate the effect of surface area on the rate of a chemical reaction
Introduction
If a chemical reaction involves one or more solid reactants, the particle sizes of the solid
reactants will affect the rate of the reaction. It is because surface area increases as the particles
become smaller. Increasing the reactants’ surface area allows the reactants approach to each other
more frequently, and so the reaction rates are often enhanced. A daily example is the burning of
charcoal. Powdered charcoal burns more fiercely than large lumps of charcoal. It is because
powdered charcoal has a much larger surface for reacting with oxygen.
Zinc metal reacts with diluted hydrochloric acid to give zinc chloride and hydrogen gas. The
rate of hydrogen gas formation allows us to observe the reaction rate conveniently. In this
experiment we will compare the reaction rates of diluted hydrochloric acid with zinc powder and
zinc granules.
Chemicals:
1M hydrochloric acid
Zinc powder
Zinc granules
Apparatus:
test tubes, spatula
Procedure:
1. Measure about 5 mL of 1 M hydrochloric acid into a test tube and add some zinc granules.
Record the observation.
2. Measure about 5 mL of 1 M hydrochloric acid into another clean test tube and add zinc powder.
The amount of zinc powder added should be approximately equal to the amount of zinc
granules added in step 1. Record the observation.
3. Compare the rate of these two reactions.
Data Analysis
1. Which of the two reactions occurs faster?
reaction rate?
2.
Did zinc granules or zinc powder give a faster
How does the particle size of solid reactant affect the reaction rate?
Chemistry Experiments (Module Kinetic Experiment)
Experiment 3: Determination of the Rate Equation for a Chemical Reaction
Objective:
To determine the rate equation for the reaction of iodine and propanone in acidic
medium
Introduction
From experiment 1 you should have realized that the rate of a chemical equation depends on
its reactants’ concentration. The relationship between the rate of a chemical reaction and its
reactants’ concentration can be expressed into a rate equation. For example, propanone reacts with
iodine in acidic medium as shown below:
CH3COOCH3(aq) + I2(aq)  CH3COCH2I(aq) + H+(aq) + I-(aq)
The rate of the reaction may depend on the concentration of propanone (CH3COCH3), iodine
(I2) and acid (H+). The rate equation of the reaction can be expressed in the form of:
Rate = k[CH3COCH3]a[I2]b[H+]c
where [CH3COCH3], [I2] and [H+] are the concentration of propanone, iodine and acid,
respectively. k is the rate constant which is a quantity that depends on reaction temperature. a, b
and c are the orders of reaction with respect to propanone, iodine and acid, respectively. Reaction
orders (a, b and c) are quantities that must be determined experimentally and cannot be deduced
from the chemical equation. The reaction order with respect to a reactant is not necessarily equal
to the reactant’s coefficient present in the balanced chemical equation.
In this experiment you are going to determine the values of k, a, b, and c for the reaction of
propanone with iodine. You will do so by running the reaction several times and varying the
concentration of the reactants one at a time.
Chemicals:
1 M hydrochloric acid
1 M propanone
0.002 M iodine solution (prepared by dissolving 0.5 g of I2 and 3.3 g of KI in 1 L of
deionized water)
Apparatus:
measuring cylinders, conical flasks
Procedure:
1. Use measuring cylinders to measure the appropriate amounts of hydrochloric acid, propanone
solution and deionized water into dry conical flasks, according to the table shown below.
2. Add appropriate amounts of iodine solution into the conical flasks. Start the stop watch.
3. Swirl the flasks gently.
4. Measure the time taken for the colour of iodine disappears completely.
Data
Volume of 1 M HCl(aq) (mL)
Volume of 1 M CH3COCH3(aq) (mL)
Volume of deionised water (mL)
Volume of 0.002M I2(aq) (mL)
 [I 2] in the time taken (mol dm-3)
Run 1
20
8
0
4
Run 2
10
8
10
4
Run 3
20
4
4
4
Run 4
20
8
2
2
Chemistry Experiments (Module Kinetic Experiment)
Time for the disappearance of the
colour (s)
Trial 1
Trial 2
Average
Initial rate of reaction (mol dm-3 s-1)
Data Analysis
1.
Calculate the initial concentration of I2 in the reaction mixture and hence the change in iodine
concentration in the time taken.
Run 1
Run 2
Run 3
Run 4
Chemistry Experiments (Module Kinetic Experiment)
2.
Calculate the rate of consumption of I2, this is the initial rate of reaction in terms of I2.
Run 1
Run 2
Run 3
Run 4
3.
Determine the order of reaction with respect to each reactant.
Chemistry Experiments (Module Kinetic Experiment)
4.
Estimate the rate constant (with unit) for the reaction using the data in Run 1.
5.
Write the rate equation for the reaction.