Experiment 16
Factors affecting Reaction Rate
(1) Reaction between Marble Chips and Hydrochloric Acid
A. Following the progress of the reaction
Marble chips (calcium carbonate) react with dilute hydrochloric acid according to the equation:
CaCO3 (s) + 2 H+ (aq)  Ca2+ (aq) + CO2 (g) + H2O (1)
While the reaction is proceeding, the mass of the reaction flask (with its contents) decreases as gas escapes. By observing the
loss in mass (which is the mass of carbon dioxide gas evolved), we can investigate how the reaction rate changes during the
course of the reaction.
1. Use a measuring cylinder to measure out 40 cm3 of 2 M hydrochloric acid into a conical flask. Plug the neck of the flask
loosely with cotton wool. Weigh out 20 g of large marble chips on a filter paper on top-pan balance. Put also the flask on
the pan of the balance (see Figure). Note the total mass and record the reading in the table overleaf.
What is the use of the cotton wool plug?
2. Remove the cotton wool plug. Add the marble chips to the acid and quickly replace the plug. Start the stopwatch
immediately. (Leave the flask with its contents and the filter paper on the balance pan.) Record the total mass every 15
seconds for the first two minutes and every half-minute for the rest of the time, until two or three consecutive constant
readings have been obtained.
Put down your results and complete the following table.
Time (min)
Total mass (g)
Total loss in mass (g)
Time (min)
Total mass (g)
Total loss in mass (g)
0
0.25
0.5
0.75
1.0
1.25
1.5
1.75
2.0
2.5
3.0
3.5
4.0
4.5
5.0
5.5
6.0
6.5
7.0
7.5
8.0
8.5
Which reactant is in excess in this experiment - marble chips or hydrochloric acid?
3. Plot a graph of total loss in mass (g) against time (minute). Draw a smooth curve through as many points as possible.
Answer the following questions:
(a) What would be the shape of the rate curve if the rate was constant throughout the reaction?
(b) Is the rate of this reaction constant throughout the reaction?
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(c) At which stage of the reaction is the rate fastest? How do you know this from the graph?
(d) Does the graph have a horizontal portion? If so, what is its significance?
(e) What is the general trend in the change in rate during the reaction? Explain why it should be so.
B. Effect of surface area on reaction rate
4. Crush some of the large marble chips to about one-quarter of the original size in a mortar. Repeat part A of the experiment
using 20 g of the crushed marble chips. Take readings as before and record them in the table below.
Time (min)
Total mass (g)
Total loss in mass (g)
Time (min)
Total mass (g)
Total loss in mass (g)
0
0.25
0.5
0.75
1.0
1.25
1.5
1.75
2.0
2.5
3.0
3.5
4.0
4.5
5.0
5.5
6.0
6.5
7.0
7.5
8.0
8.5
Plot the results on the same graph paper for part A.
Since particle size of marble chips would decrease as the reaction proceeds, initial rates in the two cases should be
compared.
Now try to answer the following questions:
(a) Is the initial rate of reaction faster or slower in this case?
(b) Which has the greater surface area the crushed marble chips used in part B or the marble chips used in part A?
(c) What effect does an increase in surface area of reactants have on the reaction rate?
(d) Why do the two rate curves become horizontal at about the same value on the Y-axis?
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(2) Investigating Effects of Concentration on Reaction Rate
Sodium thiosulphate solution and hydrochloric acid react according to the equation:
S2 O32 (aq)  2H  (aq) 
 SO2 ( g)  S ( s)  H2 O(l)
At first the sulphur is not visible. But as reaction goes on, more and more sulphur forms, turning the mixture white and then light
yellow. The degree of cloudiness of the reaction mixture at any given moment therefore represents the extent of the reaction.
In this experiment, the reaction is carried out in a beaker placed upon a piece of white paper on which a cross has been marked.
The time for the total disappearance of the cross when the contents of the beaker are viewed vertically from above is noted. This
gives a measure of the time taken for a certain amount of sulphur to be formed.
1. Mark a thick, big cross in pencil on a piece of white paper and place it on the bench.
2. You are provided with 2 M hydrochloric acid and sodium thiosulphate solution with a concentration of 40 g dm3.
Measure out 50 cm3 of the sodium thiosulphate solution with a measuring cylinder and pour it into a dry clean 100 cm 3
beaker. Stand the beaker over the cross on the paper. Have a stopwatch ready. Now quickly add 5 cm3 of the hydrochloric
acid to the beaker from a 10 cm3 measuring cylinder and start the stopwatch at the same time (see the above figure). Stir the
mixture with a glass rod to mix the reactants thoroughly. Observe the cross by looking vertically down through the solution
from above the beaker.
Note the time taken when the cross just disappears completely (see the following figure). Record this in the table overleaf.
Is the reaction complete when enough sulphur is formed to 'blot out' the cross?
3. Clean and dry the beaker and repeat Step 2 using 40 cm3 of the sodium thiosulphate solution mixed with 10 cm3 of water in
the 100 cm3 measuring cylinder. Again add 5 cm3 of the acid and note the time for the cross to just disappear completely.
Repeat the experiment using the other mixtures shown in the table below. (Note that the same depth of liquid and the same
marked piece of paper are used in all cases.)
Volume of
Na2S2O3 (aq) (cm3)
Volume of water
(cm3)
Volume of HCl (aq)
(cm3)
50
40
30
20
10
0
10
20
30
40
5
5
5
5
5
Time for cross to
disappear (s)
1
(s1)
Time
In each case, work out the reciprocal of the time for the cross to disappear (i.e. 1/time). This is proportional to the average
reaction rate during that interval of time. (Note that in each case, the amount of sulphur needed to 'blot out’ the cross is the
same.)
The volume of the original thiosulphate solution used in each case can be taken to be its relative concentration.
Explain why this is so.
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Now plot a graph of 1/time against volume of the original thiosulphate solution used.
From your graph, explain how the reaction rate depends on the concentration of a reactant.
(3) Effect of Catalyst of Rate of Decomposition of Hydrogen Peroxide
1. Using a measuring cylinder, add 50 cm3 of '1-volume' hydrogen peroxide solution to a conical flask. Observe the solution
carefully.
Do you see any signs of decomposition?
2. Set up the apparatus as shown in the above figure. Make sure that the syringe reading is zero. Remove the bung. Add 0.5 g of
manganese(IV) oxide granules (10 - 30 mesh) to the solution and immediately replace the bung. Start the stopwatch as quickly
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as possible. Read the volume every 15 seconds for the first 2 minutes and every 30 seconds for the rest of the experiment,
until you have obtained 2 or 3 consecutive constant readings. Before taking each reading, rotate the plunger of the syringe to
ensure that it has not become jammed. Record your results in the table below.
3. Repeat the experiment with fresh '1-volume' hydrogen peroxide solution but this time add 1.5 g of the manganese (IV) oxide
granules. Record your results in the table below.
Time (min)
0.5 g MnO2 added
1.5 g MnO2 added
Time (min)
0.5 g MnO2 added
1.5 g MnO2 added
0.25
0.5
0.75
1.0
1.25
1.5
1.75
2.0
2.5
3.0
3.5
4.0
4.5
5.0
5.5
6.0
6.5
7.0
Plot graphs of volume of oxygen (cm3) against time (minute) below.
Draw a tangent at each of the rate curves at time t = 0. Now answer the following questions:
(a) What is the effect of manganese(IV) oxide catalyst on the rate of decomposition of hydrogen peroxide?
(b) What is the effect of increased mass of manganese(IV) oxide on the (i) initial rate (ii) total time of reaction?
(c) Do you think the manganese(IV) oxide is used up in the experiments?
4. Design an experiment to test your prediction to (c) above.
Describe your method.
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