© MHS 2010 1

Chapter 1: Rate of Reaction 1. Rate of Reaction

- Rate of reaction is the speed at which reactants are converted into products in a

chemical reaction.

- Different chemical reactions occur at different rates.

Fast reaction

- The time taken for a fast reaction is short.

- The reactants are quickly converted to the

products.

- The rate of reaction is said to be high

- Examples:

- Fading of dyes on a shirt under hot sun

- Cooking a chicken using a microwave oven

- Burning of petrol in a car engine

- Striking a match

- Ripening of tomatoes

Slow reaction

- The time taken for a slow reaction is

long.

- The rate of reaction for a slow reaction is

low.

- Examples:

- A piece of newspaper turning yellow

- The weathering of limestone by acid rain

- Rusting of a water pipe

2. Observable changes for measuring rate of reaction

Precipitation formation

Colourchange

Pressurechanges

Temperature changes

Change inmass during the reaction

Volume of a gas liberated

Observable changes

- During the reaction of magnesium ribbon

with dilute hydrochloric acid, HCl, two

visible changes are

- Decrease in the mass of marble chips

(reactant)

- Increase in the volume of carbon dioxide

gas, CO2 (product)

Mg (s) + 2HCl (aq) MgCl2 (aq) + H2 (g)

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3. Rate of reaction is measured by:

time taken=change in selested quantity

Rate of reaction

Example: Magnesium ribbon and dilute hydrochloric acid

Rate of reaction

decrease in mass of magnesium ribbon (reactant)

= time taken

time taken= increase in volume of hydrogen gas (product)

Rate of reaction

4 Average rate of reaction

- is the average value of the rate of reaction within a specified period of time.

5. Rate of on reaction at a give time

- is the actual rate of reaction at that instant

- It is also known as the instantaneous rate of reaction

- Gradient of the curve at that instant

© MHS 2010 3

Activity

Aim: To determine the average rate of reaction and the instantaneous rate of reaction

Apparatus:

150 cm3 conical flask, 50 cm3 measuring cylinder, stopper with delivery tube, burrette, retort

stand and clamp, stopwatch, basin and electronic balance

Materials:

5 g of granulated zinc, 0.1 mol dm–3 hydrochloric acid and water.

Procedure:

1. 25 cm3 of 0.1 mol dm–3 hydrochloric acid is measured and pored into a conical flask.

2. About 5 g of granulated zinc is weighed using an electronic balance.

3. A burette is filled full with water. It is then inverted into a basin containing water and

clamped vertically using a retort stand.

4. The water level in the burette is adjusted to 50 cm3 mark.

5. The set up of the apparatus is shown in Figure 1.

6. The granulated zinc is added into the conical flask containing hydrochloric acid.

7. The conical flask is closed immediately with a stopper which is joined to delivery tube. At

the same time, the stopwatch is started.

8. The conical flask is shaken steadily throughout the whole activity.

9. The volume of gas collected in the burette by downward displacement of water is recorded

at 30-second intervals for a period of 5 minutes.

10. The results are recorded in a table.

Results: Time (s) 0 30 60 90 120 150 180 210 240 270 300

Burette reading (cm3)

50.00 38.00 30.00 26.00 23.00 20.00 18.50 17.50 17.00 17.00 17.00

Volume of gas (cm3)

0.00 12.00 20.00 24.00 27.00 30.00 31.50 32.50 33.00 33.00 33.00

Calculation:

From your graph,

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1. To calculate the average rate of reaction:

a) Average rate of reaction for the overall reaction

b) Average rate of reaction in the first minute

c) Average rate of reaction in the second minute

2. To calculate the instantaneous rate of reaction:

a) Instantaneous rate of reaction at 30 seconds

b) Instantaneous rate of reaction at 90 seconds

Observation Inference

1. Gas bubbles are produced. 1. Hydrogen gas is produced.

2. Liberation of gas is the fastest at the

beginning. Then, liberation of gas

slows down and stops at 240 s.

2. The reaction is the fastest at the

beginning. The reaction slows down and

then is complete at 240 s.

3. From the graph plotted, it can be

seen that the gradient of the curve

decreases with time. The gradient

becomes zero at 240 s.

3. The rate of reaction decreases with time.

The rate of reaction becomes zero at 240s.

Conclusion:

The rate of reaction decreases with time and then becomes zero.

Explanation:

1. Reaction between hydrochloric acid and zinc produces hydrogen gas and zinc chloride

solution.

2. The instantaneous rate of reaction at 30 seconds is higher than that of 90 seconds. This is

because

a) the concentration of hydrochloric acid is higher at 30 seconds

b) the total surface area of granulated zinc is larger at 30 seconds

3. The rate of reaction decreases with time because the concentration of hydrochloric acid and

the total surface area of granulated zinc decreases with time. When all the concentration of

hydrochloric acid is completely reacted, the rate of reaction becomes zero.

© MHS 2010 5

CatalystsTemperature and Pressure

Concentrationof Reactants

Size ofReactants

Factors Affecting the Rate of Reaction

© MHS 2010 6

B Factors Affecting the Rate of Reaction

I Effect of surface area on the rate of reaction

1. For a fixed mass of solid reactant, the smaller the size of the reactant, the larger will be the

total exposed surface area, thus the higher will be the rate of reaction.

Examples:

i. CaCO3 (s) + 2HCl (aq) CaCl2 (aq) + CO2 (g) + H2O (l)

ii. Mg (s) + 2HCl (aq) Mg Cl2 (aq) + H2 (g)

iii. Zn (s) + H2SO4 (aq) ZnSO4 (aq) + H2 (g)

Experiment

Aim:

To investigate the effect of size of a reactant on the rate of reaction

Problem statement:

Does small size of marble chips increase the rate of reaction?

Hypothesis:

The smaller the size of marble chips, the higher the rate of reaction.

Variables:

Manipulated variable: size of marble chips

Responding variable: Rate of reaction

Controlled variables : Mass of marble chips, volume and concentration of HCl acid,

Temperature

Apparatus:

50 cm3 measuring cylinder, 150 cm3 conical flask, stopper with delivery tube, basin, burette,

electronic balance and stopwatch

Materials:

0.1 mol dm–3 hydrochloric acid, 2 g of large marble chips, 2 g of small marble chips and water

Procedure:

1. 40 cm3 of 0.1 mol dm–3 hydrochloric acid is measured and pored into a conical flask.

2. About 2 g of large marble chips are weighed using an electronic balance.

3. A burette is filled full with water. It is then inverted into a basin containing water and clamped

vertically using a retort stand.

4. The water level in the burette is adjusted to 50 cm3 mark.

5. The set up of the apparatus is shown in Figure 2.

© MHS 2010 7

6. The large marble chips are added into the conical flask containing hydrochloric acid.

7. The conical flask is closed immediately with a stopper which is joined to delivery tube. At the

same time, the stopwatch is started.

8. The conical flask is shaken steadily throughout the whole experiment.

9. The volume of gas collected in the burette by downward displacement of water is recorded at

30-second intervals for a period of 5 minutes.

10. The results are recorded in a table.

11. The experiment is repeated using 2 g of small marble chips to replace 2 g of large marble

chips. All the conditions remain unchanged.

Results:

Experiment 1: Large marble chips Time (s) 0 30 60 90 120 150 180 210 240 270 300

Burette reading (cm3)

50.00 43.00 38.00 34.00 30.00 27.50 25.00 23.00 21.00 20.00 19.00

Volume of gas (cm3)

0.00 7.00 12.00 16.00 20.00 22.50 25.00 27.00 29.00 30.00 31.00

Experiment 2: small marble chips Time (s) 0 30 60 90 120 150 180 210 240 270 300

Burette reading (cm3)

50.00 39.00 32.00 27.00 22.50 19.00 16.00 13.50 11.50 10.00 9.00

Volume of gas (cm3)

0.00 11.00 18.00 23.00 27.50 31.00 34.00 36.50 38.50 40.00 41.00

Calculation:

From your graph,

1. To calculate the average rate of reaction at 200 seconds:

a) Experiment 1

b) Experiment 2

c) By comparison, the average rate of reaction for experiment 2 is higher than experiment 1 at

300 seconds.

2. To calculate the instantaneous rate of reaction at 120 seconds:

© MHS 2010 8

a) Experiment 1

b) Experiment 2

c) By comparison, the instantaneous rate of reaction for experiment 2 is higher than

experiment 1 at 120 seconds.

Observation Inference

1. Gas bubbles are produced. 1. Carbon dioxide gas is produced.

2. Liberation of gas is the fastest at the

beginning and it slows down later.

2. The reaction is the fastest at the

beginning. The reaction slows down and

then is complete at 240 s.

3. From the graph plotted, it can be

seen that the initial gradient of the

curve for the experiment using small

marble chips is higher than the initial

gradient of the curve for the

experiment using large marble chips.

3. The rate of reaction for the experiment

using small marble chips is higher that the

rate of reaction for the experiment using

large marble chips

Conclusion:

1. small marble chips have a higher rate of reaction.

2. The hypothesis is accepted.

Explanation:

1. The reaction between hydrochloric acid and marble chips produces carbon dioxide gas, calcium

chloride solution and water.

2. Based on the graphs plotted, the curve for experiment 2 is steeper than experiment 1

3. This means the rate of reaction for experiment 2 (using small marble chips) is higher than

experiment 1 (using large marble chips). This is because small marble chips have a larger total

exposed surface area.

4. The curves becomes less steep with time due to the decrease in concentration of hydrochloric

acid and mass of marble chips.

5. If both experiments are continued until the reaction is completed, the following graph will be

obtained.

© MHS 2010 9

6. Both experiments produce the same maximum volume of carbon dioxide gas. This is because

both experiments use the same

a) concentration and volume of hydrochloric acid

b) mass of marble chips (calcium carbonate)

7. The maximum volume of carbon dioxide gas collected is less than the theoretical volume

because a small volume of carbon dioxide gas has dissolved in the water when it is collected in

the burette.

8. To overcome the problem, one should pass the carbon dioxide gas through the water for a few

minutes before starting the experiment. This is to saturate the water with carbon dioxide before

collecting the gas in the burette.

Use of collision theory to explain the effect of the size of reactant /

total surface area on the rate of reaction

1. When the size of a fixed mass of solid reactant is smaller, the total surface area exposed to

collision with the particles of the other reactants is bigger. Thus, the frequency of collision among the reacting particles at the surface of the solid

reactant increases.

This leads to an increase in the frequency of effective collision and hence, a higher rate of reaction.

© MHS 2010 10

II Effect of Concentration on the Rate of Reaction

1. The higher the concentration of liquid reactant, the higher the rate of reaction.

Examples:

i. Na2S2O3 (aq) + H2SO4 (aq) Na2SO4 (aq) + S (s) + SO2 (g) + H2O (l)

ii. CaCO3 (s) + 2HCl (aq) CaCl2 (aq) + CO2 (g) + H2O (l)

iii. CaCO3 (s) + 2HNO3 (aq) Ca(NO3) (aq) + CO2 (g) + H2O (l)

Experiment

Aim:

To investigate the effect of concentration on the rate of reaction

Problem statement:

Does high concentration of sodium thiosulphate solution, the shorter the time taken for the

mark ‘X ‘ to disappear from sight?

Hypothesis:

The higher the concentration of sodium thiosulphate solution, the shorter the time taken for the

mark ‘X ‘ to disappear from sight.

Variables:

Manipulated variable: Concentration of sodium thiosulphate solution

Responding variable: Time taken for the mark ‘X ‘ to disappear from sight

Controlled variables : Temperature, total volume of the reacting mixture, concentration and

volume of sulphuric acid, size of conical flask

Apparatus:

10 cm3 measuring cylinder , 50 cm3 measuring cylinder, 150 cm3 conical flask and stopwatch

Materials:

1.0 mol dm–3 hydrochloric acid, 0.2 mol dm–3 sodium thiosulphate solution, distilled water and

white paper with a mark ‘X ‘ at the centre

Procedure:

1. 45 cm3 of 0.2 mol dm–3 sodium thiosulphate solution is measured using a 50 cm3

measuring cylinder and pored into a conical flask.

© MHS 2010 11

2. The conical flask is placed on top of a piece of white paper with a mark ‘X ‘ at the centre.

3. 5 cm3 of 1.0 mol dm–3 sulphuric acid is measured using a 10 cm3 measuring cylinder.

4. The sulphuric acid is poured quickly and carefully into the conical flask. At the same time,

stopwatch is started immediately.

5. The mixture in the conical flask is swirled for a few time. It is then placed back on the white

paper.

6. The mark ‘X‘ is observed vertically from the top through the solution as shown in Figure 3.

7. The stopwatch is stopped immediately once the mark ‘X‘ disappear from sight.

8. The time t required for the mark ‘X‘ to disappear from sight is recorded.

9. The experiment is repeated four more times using different volumes of 0.2 mol dm–3

sodium thiosulphate solution which is diluted with different volumes of distilled water as

shown in table below. All other conditions remain unchanged.

10. The results are recorded.

Results:

Set I II III IV V

Volume of 0.2 mol dm–3 sodium thiosulphate

solution, V1 (cm3) 45 40 30 20 10

Volume of distilled water (cm3) 0 5 15 25 35

Volume of 1.0 mol dm–3 sulphuric acid (cm3) 5 5 5 5 5

Total volume of the mixture, V2 (cm3)

Concentration of sodium thiosulphate solution

that reacts, M2 (mol dm–3)

2

112 V

VMM

Time taken (s) 18 20 27 41 82

)(1 1sTime

Observation Inference

1. A yellow precipitate formed 1. A yellow precipitate formed is sulphur 2. A pungent smell is produced 2. Sulphur dioxide gas is produced 3. When the concentration of sodium

thiosulphate decreases, the time taken for the mark ‘X‘ to disappear from sight becomes longer.

3. When the concentration of sodium thiosulphate decreases, the production of the yellow precipitate of sulphur becomes slower. Thus, the rate of reaction decreases

Conclusion:

1. An increase in concentration of sodium thiosulphate increases the rate of reaction 2. The hypothesis is accepted.

© MHS 2010 12

Explanation: 1. The reaction between sulphuric acid and sodium thiosulphate solution produces sodium

sulphate solution, sulphur, sulphur dioxide gas and water. Chemical equation:

Ionic equation:

2. Based on the graph I, the curve implies that as the concentration of sodium thiosulphate

solution becomes lower, the time taken for the mark ‘X‘ to disappear from sight becomes

longer.

3. Based on graph II, the straight line implies that the concentration of sodium thiosulphate

solution is directly proportional to Time

1 .

4. In other words, the concentration of sodium thiosulphate solutions is direct proportional to

the rate of reaction.

5. Conical flask of the same size and shape are used in this experiment. If a bigger conical

flask is used, the time taken for the mark ‘X‘ to disappear from sight becomes longer.

6. This is because bigger conical flask has a larger base area. The mixture of 50 cm3 solution

becomes shallower. A bigger amount of yellow precipitate is needed to turn the mark ‘X‘

invisible from sight.

7. If the experiments is repeated using 1.0 mol dm–3 ydrochloric acid to replace 1.0 mol dm–3

sulphuric acid, the rate of reaction will decreases.

8. This because hydrochloric acid is a strong monoprotic acid whereas sulphuric acid is a

strong diprotic acid. Although the concentration of acids is the same, the concentration of

hydrogen ions in sulphuric acid is twice the concentration of hydrogen ions in hydrochloric

acid.

Use of collision theory to explain the effect of concentration on the rate of reaction

1. When the concentration of the solution of a reactant increase, the number of particles per unit volume of the solution of this reactant also increase.

2. with more particles per unit volume of the solution, the number of collisions per unit time between the reacting particles increases.

3. Thus, the frequency of effective collision increases and hence, the rate of reaction increases

© MHS 2010 13

III Effect of temperature on the rate of reaction

1. As the temperature of the reactant increases, the rate of reaction increases.

2. At higher temperature, the reactant particles have greater kinetic energy resulting in a higher

speed of movement and thus more frequent effective collisions.

3. Cooling a mixture will slow down the particles and fewer collisions take place.

4 Temperature is directly proportional to the rate of reaction. That is, the higher the temperature,

the higher the rate of reaction.

Experiment

Aim:

To investigate the effect of temperature on the rate of reaction

Problem statement:

Does high temperature of sodium thiosulphate solution decrease the time taken for the mark

‘X ‘ to disappear from sight?

Hypothesis:

The higher the temperature of sodium thiosulphate solution, the shorter the time taken for the

mark ‘X ‘ to disappear from sight.

Variables:

Manipulated variable: Temperature of sodium thiosulphate solution

Responding variable: Time taken for the mark ‘X ‘ to disappear from sight

Controlled variables : Concentration and volume of sodium thiosulphate solution,

concentration and volume of sulphuric acid, size of conical flask

Apparatus:

10 cm3 measuring cylinder , 50 cm3 measuring cylinder, 150 cm3 conical flask, stopwatch,

thermometer, Bunsen burner and wire gauze.

Materials:

1.0 mol dm–3 hydrochloric acid, 0.2 mol dm–3 sodium thiosulphate solution, distilled water and

white paper with a mark ‘X ‘ at the centre

Procedure:

1. 50 cm3 of 0.2 mol dm–3 sodium thiosulphate solution is measured using a 50 cm3

measuring cylinder and pored into a conical flask.

2. The temperature of the solution is measured using a thermometer.

3. The conical flask is placed on top of a piece of white paper with a mark ‘X ‘ at the centre.

4. 5 cm3 of 1.0 mol dm–3 sulphuric acid is measured using a 10 cm3 measuring cylinder.

5. The sulphuric acid is poured quickly and carefully into the conical flask. At the same time,

stopwatch is started immediately.

6. The mixture in the conical flask is swirled for a few time. It is then placed back on the white paper.

© MHS 2010 14

7. The mark ‘X‘ is observed vertically from the top through the solution as shown in Figure 4.

8. The stopwatch is stopped immediately once the mark ‘X‘ disappear from sight.

9. The time t required for the mark ‘X‘ to disappear from sight is recorded.

10. The experiment is repeated using 50 cm3 0.2 mol dm–3 sodium thiosulphate solution at

350C, 400C, 450C and 500C respectively. All other conditions remain unchanged.

11. The results are recorded.

Results:

Set Temperature (0C) Time, t (s) Time

1 (s–1)

I 30.0 50

II 35.0 28

III 40.0 20

IV 45.0 15

V 50.0 12

Observation Inference

1. A yellow precipitate formed 1. A yellow precipitate formed is sulphur

2. A pungent smell is produced 2. Sulphur dioxide gas is produced

3. When the temperature of sodium

thiosulphate decreases, the time

taken for the mark ‘X‘ to disappear

from sight becomes shorter.

3. When the temperature of sodium

thiosulphate decreases, the

production of the yellow precipitate of

sulphur becomes faster. Thus, the

rate of reaction increases

Conclusion:

1. An increase in temperature of sodium thiosulphate increases the rate of reaction

2. The hypothesis is accepted.

Explanation:

1. The reaction between sulphuric acid and sodium thiosulphate solution produces sodium

sulphate solution, sulphur, sulphur dioxide gas and water.

Chemical equation:

Ionic equation:

2. Based on the graph I, the curve implies that as the temperature of sodium thiosulphate

solution becomes higher, the time taken for the mark ‘X‘ to disappear from sight becomes

© MHS 2010 15

shorter.

3. Based on graph II, the straight line implies that the temperature of sodium thiosulphate

solution is directly proportional to Time

1 .

4. In other words, the temperature of sodium thiosulphate solutions is directly proportional to

the rate of reaction.

Use of collision theory to explain the effect of Temperature on the rate of reaction

1. When the temperature increase, the average kinetic energy of the reacting particles

increases.

2. The particles move faster and collide more often with one another. The frequency of

collision among the reacting particles increases.

3. Particles also have more energy to overcome the activation energy.

4. These changes increase the frequency of effective collision.

5. Hence, the rate of reaction increase.

Use of collision theory to explain the effect of Temperature on the rate of reaction

1. when the pressure of gaseous reactant increases, the particles are compressed to

occupy a smaller volume. The number of gas particles per unit volume increases.

2. The frequency of collision among the reacting particles increases.

3. This increases the frequency of effective collision.

4. Hence, the rate of reaction increases.

Lower pressure results in less collision among

particles as the particles are far from one

another.

Higher pressure results in more collision

among particles as the particles are pushed

near to one another.

© MHS 2010 16

IV Effect of catalyst on the rate of reaction

1. A catalyst is a substance which alters the rate of a chemical reaction while it remains

chemically unchanged at the end of the reaction.

2. A positive catalyst increases the rate of reaction whereas a negative catalyst (inhibitors)

decreases the rate of reaction.

3. Most catalysts are transition elements or compounds of transition elements such as iron,

nickel and copper (II) sulphate, CuSO4.

It is specific in its action. Itcan only catalyse aparticular reaction

During a reaction, catalystremains chemicallyunchanged but mayundergo physicalchanges. For example,catalyst may turn intopowder during the reaction

Only a small amount ofcatalyst is needed toincreases the rate ofreaction. An increase inthe quantity of catalystwill increase the rate ofreaction but only a veryslight increase

Does not change thequantity of products formed

Alters the rate of reaction

Catalyst

Experiment

Aim: To investigate the effect of catalyst on the rate of reaction

Problem statement:

Does the presence of a catalyst increase the rate of reaction?

Hypothesis:

The presence of a catalyst increase the rate of reaction.

Variables:

Manipulated variable: the presence of the catalyst

Responding variable: Rate of reaction

Controlled variables : Mass of zinc granules, volume and concentration of HCl acid,

Temperature

© MHS 2010 17

Apparatus:

50 cm3 measuring cylinder, 150 cm3 conical flask, stopper with delivery tube, basin, burette,

electronic balance and stopwatch

Materials:

0.1 mol dm–3 hydrochloric acid, 5 g of zinc granules, 0.5 mol dm–3 copper (II) sulphate solution

and water

Procedure:

1. 50 cm3 of 0.1 mol dm–3 hydrochloric acid is measured and pored into a conical flask.

2. About 5 g of granulated zinc is weighed using an electronic balance.

3. A burette is filled full with water. It is then inverted into a basin containing water and

clamped vertically using a retort stand.

4. The water level in the burette is adjusted to 50 cm3 mark.

5. The set up of the apparatus is shown in Figure below.

6. The granulated zinc is added into the conical flask containing hydrochloric acid.

7. 5 cm3 of 0.5 mol dm–3 copper (II) culphate solutions is added to the conical flask.

8. The conical flask is closed immediately with a stopper which is joined to delivery tube. At

the same time, the stopwatch is started.

9. The conical flask is shaken steadily throughout the whole experiment.

10. The volume of gas collected in the burette by downward displacement of water is recorded

at 30-second intervals for a period of 5 minutes.

11. The experiment is repeated without adding 5 cm3 of 0.5 mol dm–3 copper (II) culphate

solutions to the conical flask.

12. The results are recorded in a table.

Results:

Experiment 1: With 5 cm3 of 0.5 mol dm–3 copper (II) culphate solutions (with catalyst) Time (s) 0 30 60 90 120 150 180 210 240 270 300

Burette reading (cm3)

50.00 37.00 29.00 23.00 19.00 16.00 14.00 12.00 10.50 9.50 9.00

Volume of gas (cm3)

0.00 13.00 21.00 27.00 31.00 34.00 36.00 38.00 39.50 40.60 41.00

© MHS 2010 18

Experiment 2: Without 5 cm3 of 0.5 mol dm–3 copper (II) culphate solutions (without catalyst) Time (s) 0 30 60 90 120 150 180 210 240 270 300

Burette reading (cm3)

50.00 43.00 38.00 34.00 30.00 27.00 25.00 22.00 21.00 19.00 18.00

Volume of gas (cm3)

0.00 7.00 12.00 16.00 20.00 23.00 25.00 28.00 29.00 31.00 32.00

Calculation:

From your graph,

1. To calculate the average rate of reaction at 200 seconds:

a) Experiment 1

b) Experiment 2

c) By comparison, the average rate of reaction for experiment 2 is higher than experiment

1 at 300 seconds.

2. To calculate the instantaneous rate of reaction at 120 seconds:

a) Experiment 1

b) Experiment 2

c) By comparison, the instantaneous rate of reaction for experiment 2 is higher than

experiment 1 at 120 seconds.

Observation Inference

1. Gas bubbles are produced. 1. Hydrogen gas is produced.

2. Liberation of gas is the fastest at the

beginning and it slows down later.

2. The reaction is the fastest at the beginning

and then it slows down.

3. From the graph plotted, it can be

seen that the gradient of the curve for

the experiment with a catalyst (copper

(II) sulphate solution) is higher than

the gradient of the curve for the

experiment without a catalyst.

3. The rate of reaction for the experiment

using a catalyst is higher that the rate of

reaction for the experiment without a

catalyst.

Conclusion:

1. The presence of catalyst increase the rate of reaction.

2. The hypothesis is accepted.

Explanation:

1. The reaction between hydrochloric acid and zinc produces hydrogen gas and zinc chloride

solution.

2. Based on the graphs plotted, the curve for experiment 1 is steeper than experiment 2.

3. This means the rate of reaction for experiment 1 (using catalyst) is higher than experiment

© MHS 2010 19

2 (without catalyst).

4. The curves becomes less steep with time due to the decrease in concentration of

hydrochloric acid and mass of zinc granules.

5. If both experiments are continued until the reaction is completed, the following graph will

be obtained.

6. Both experiments produce the same maximum volume of hydrogen gas. This is because

both experiments use the same

a) concentration and volume of hydrochloric acid

b) mass of zinc granules

Use of collision theory to explain the effect of catalyst on the rate of reaction

1. When a positive catalyst is used in a chemical reaction, it provides as alternative path

with a lower activation energy.

2. More colliding particles are able to overcome the lower activation energy.

3. The increases the frequency of effective collision.

4. Hence the rate of reaction increases.

© MHS 2010 20

Activation energy

1. According to the kinetic theory of matter, particles of matter are in continuous motion and

constantly in collision with each other.

2. During a reaction, the particles of the reactants, whether atoms, molecules or ions, must

collide with each other for bond breaking and then bond formation to occurs.

3. Only those collisions which achieved a minimum amount of energy, called activation

energy, and with the correct orientation, will result in a reactions.

4. These collisions are known as called effective collisions.

5. If the particles collide with energy less than the activation energy needed for reaction or

with the wrong orientation, they simply bounce apart without reacting.

6. These collisions are known as ineffective collisions.

7. In the energy profile diagram, the activation energy is the difference is energy between the

energy of the reactants and the energy shown by the peak of the curve,

8. It is the energy barrier that must be overcome by the colliding particles of the reactants in

order for reaction to occur.

9. The number of effective collisions occurring in one second is called the frequency of

effective collision.

10. when the frequency of effective collision is high, the rate of reaction is high.