IGCSE Chemistry

Cambridge IGCSE™

Chemistry STUDENT’S BOOK

Chris Sunley, Sam Goodman

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In this section you will explore the speed or rate of chemical reactions as well as how the rate can be controlled. You will learn about reactions that go from reactants to products and go back again, known as reversible or equilibrium reactions. These are a particular challenge for chemists who are trying to make a product when it keeps reverting to the reactants! You will also learn about redox reactions which, as the name suggests, involve both reduction and oxidation. You will also have opportunities to practise writing and balancing chemical equations, as introduced in Section 3, Stoichiometry.

Starting points 1. Give an example of a very rapid, almost instantaneous, chemical reaction. Now give an example of a very slow one.

2. What is a catalyst? Can you give any examples of catalysts used in everyday life?

3. Can you name a process involving a physical change that can be easily reversed. Now can you name a chemical process that is reversible?

SYLLABUS SECTIONS COVERED 6.1 Physical and chemical changes 6.2 Rate of reaction 6.3 Reversible reactions and equilibrium 6.4 Redox

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6

Chemical reactions

∆ Rusting is a very slow reaction.

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Rate of reaction INTRODUCTION

Some chemical reactions take place extremely quickly. For example, when petrol is ignited it combines with oxygen almost instantaneously. Reactions like these have a high rate. Other reactions are much slower, for example when an iron bar rusts in the air; reactions like these have a low rate.

Chemical reactions can be controlled and made to be quicker or slower. This can be very important in situations like food production, either by slowing down or ∆ Fig. 6.1 Petrol igniting. increasing the rate at which food ripens, or in the chemical industry, where the rate of a reaction can be adjusted to an optimum level.

LEARNING OBJECTIVES ✓ Identify physical and chemical changes, and describe the differences between them. ✓ Describe the effect on the rate of reaction of changing: the concentration of solutions, the pressure of gases, the surface area of solids, the temperature, and adding or removing a catalyst including enzymes. ✓ State that a catalyst increases the rate of a reaction and is unchanged at the end of a reaction. ✓ Describe practical methods for investigating the rate of a reaction including change in mass of a reactant or a product and the formation of a gas. ✓ Interpret data, including graphs, from rate of reaction experiments. ✓ SUPPLEMENT Describe collision theory in terms of: number of particles per unit volume; frequency of collisions between particles; kinetic energy of particles; activation energy, Ea. ✓ SUPPLEMENT Describe and explain the effect on the rate of reaction of changing: the concentration of solutions, the pressure of gases, the surface area of solids, the temperature, and adding or removing a catalyst including enzymes, using collision theory. ✓ SUPPLEMENT State that a catalyst decreases the activation energy, Ea, of a reaction. ✓ SUPPLEMENT Evaluate practical methods for investigating the rate of a reaction including change in mass of a reactant or a product and the formation of a gas.

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CHEMICAL REACTIONS

KNOWLEDGE CHECK ✓ The particles in the three states of matter (solid, liquid and gas) have different arrangements, movement and energy. ✓ The course of a reaction can be shown in a reaction pathway diagram. ✓ Balanced chemical equations are used to describe reactions.

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PHYSICAL AND CHEMICAL CHANGES A chemical change, or chemical reaction, is quite different from physical changes that occur, for example, when sugar dissolves in water.

In a chemical change, one or more new substances are produced. In many cases an observable change is apparent, for example, the colour changes or a gas is produced.

An apparent change in mass can occur. This change is often quite small and difficult to detect unless accurate balances are used. Mass is conserved in all chemical reactions – the apparent change in mass usually occurs because one of the reactants or products is a gas (whose mass may not have been measured).

An energy change is almost always involved. In most cases energy is released and the surroundings become warmer. In some cases energy is absorbed from the surroundings, and so the surroundings become colder. Note: Some physical changes, such as evaporation, also have energy changes.

RATE OF REACTION A quick reaction takes place in a short time. It has a high rate of reaction. As the time taken for a reaction to be completed increases, the rate of the reaction decreases. In other words: Speed Quick or fast Slow

Rate High Low

Completion time Short Long

∆ Table 6.1 Speed, rate and time.

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Collision theory For a chemical reaction to occur, the reacting particles (atoms, molecules or ions) must collide. The energy involved in the collision must be enough to break the chemical bonds in the reacting particles – or the particles will just bounce off one another. The more particles there are in a particular volume, the greater the frequency of collisions.

Similarly the greater the kinetic energy of the particles, the greater the frequency of collisions.

A collision that has enough energy to result in a chemical reaction is an effective collision.

RATE OF REACTION

SUPPLEMENT

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unsuccessful

successful

∆ Fig. 6.2 Particles must collide with sufficient energy to make an effective collision.

Some chemical reactions occur extremely quickly (for example, the explosive reaction between petrol and oxygen in a car engine) and some more slowly (for example, iron rusts over days or weeks). This is because they have different activation energies. Activation energy acts as a barrier to a reaction. It is the minimum amount of energy required in a collision for a reaction to occur. As a general rule, the bigger the activation energy, the slower the reaction will be at a particular temperature.

Energy

activation energy (Ea)

reactants ∆H = – (energy transferred) products

Reaction pathway

REMEMBER

The ‘barrier’ preventing a reaction from occurring is called the activation energy. If the activation energy of a reaction is low, more of the collisions will be effective and the reaction will proceed quickly. If the activation energy is high, a smaller proportion of collisions will be effective and the reaction will be slower.

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CHEMICAL REACTIONS

∆ Fig. 6.3 Reaction pathway diagram.

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QUESTIONS 1. State the main difference between a physical change and a chemical change.

2. In a chemical change there is often an apparent change in mass even though mass cannot be created or destroyed in a chemical reaction. Suggest a possible cause of this apparent change in mass.

3. SUPPLEMENT In the collision theory, state what two things must happen for two particles to react.

4. SUPPLEMENT Define an effective collision.

5. SUPPLEMENT Describe, using a diagram, what is meant by the term activation energy.

Monitoring the rate of a reaction The rate of a reaction changes as the reaction proceeds. There are some easy ways of monitoring this change.

When marble (calcium carbonate) reacts with hydrochloric acid, the following reaction starts straight away: calcium carbonate + hydrochloric acid → calcium chloride + carbon dioxide + water CaCO3(s) + 2HCl(aq) → CaCl2(aq) + CO2(g) + H2O(l) The reaction can be monitored as it proceeds either by measuring the volume of gas being formed or by measuring the change in mass of the reaction flask.

The volume of gas produced in this reaction can be measured using the apparatus shown in Fig. 6.4. The hydrochloric acid is put into the conical flask, the marble chips are added, the stopper is quickly fixed into the neck of the flask and the stopclock is started. gas syringe

calcium carbonate (marble chips)

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∆ Fig. 6.4 Monitoring the rate of a reaction.

RATE OF REACTION

dilute hydrochloric acid

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The reaction will start immediately, effervescence (bubbling) will occur in the flask as the carbon dioxide gas is produced and the plunger on the syringe will start to move. Measuring the volume of gas in the syringe every 10 seconds will indicate how the total amount of gas produced changes as the reaction proceeds. The change in the rate of the reaction with time can be shown on a graph of the results (see Fig. 6.6).

To measure the change in mass in the same reaction, the apparatus shown in Fig. 6.5 can be used. The hydrochloric acid is put into the conical flask, the marble chips are added, the cotton wool plug is put in the neck of the flask and the stopclock is started. The mass of the flask and contents is measured as soon as the plug is inserted and then every 10 seconds as the reaction occurs. The mass will decrease as carbon dioxide gas escapes from the flask. loose cotton wool plug to stop spray escaping

dilute hydrochloric acid

stopclock

marble chips

balance

∆ Fig. 6.5 Measuring the change in mass.

Drawing a graph of the results shows the change in the rate of the reaction over time.

Graphs of the results from both experiments have almost identical shapes. The rate of the reaction decreases as the reaction proceeds.

100

3

90

0.18

Loss in mass /g

70

2

60 50 40 30 20 10

3

0.16

0.10 0.08 0.06 0.02

10 20 30 40 50 60 70 80 90 100

2

0.12

0.04

1

0

0.14

0

1 10 20 30 40 50 60 70 80 90 100

Time / seconds

Time / seconds

∆ Fig. 6.6 Volume of carbon dioxide produced or loss in mass.

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CHEMICAL REACTIONS

Gas volume / cm3

80

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Loss in mass during the reaction The rate of the reaction at any point can be calculated from the gradient of the curve. The shapes of the graphs can be divided into three regions.

1. At this point, the curve is the steepest (has the highest gradient) and the reaction has its highest rate. The maximum number of reacting particles are present and the number of effective collisions per second is at its greatest.

2. The curve is not as steep (has a lower gradient) at this point and the rate of the reaction is lower. Fewer reacting particles are present and so the number of effective collisions per second is less.

3. The curve is horizontal (gradient is zero) and the reaction is complete. At least one of the reactants has been completely used up and so no further collisions can occur between the two reactants.

REMEMBER

In experiments like these it is helpful to have a good understanding of the types of variables involved. The factor you are investigating is called the independent variable – when investigating how the reaction between marble and hydrochloric acid changes over time, time is the independent variable. A dependent variable is changed by the independent variable – in the marble and hydrochloric acid reaction, the volume of carbon dioxide produced is the dependent variable. Other variables involved are control variables and are not allowed to change to ensure a ‘fair test’. So temperature could be a control variable in the reaction between marble and hydrochloric acid.

In chemical reactions it is very rare that exact (as predicted by the equation) quantities of reactants are used. In the marble and hydrochloric acid reaction all the marble may be used up (it is called the limiting reactant) but not all the hydrochloric acid; some is left when the reaction has stopped (it is in excess).

QUESTIONS 1. State what piece of apparatus can accurately measure the volume of gas produced in a reaction.

2. On a volume versus time graph, explain what a horizontal line shows. reaction have a steeper volume/time gradient at the beginning?

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What can change the rate of a reaction?

There are five key factors that can change the rate of a reaction: • concentration (of a solution) • pressure (of a gas)

RATE OF REACTION

3. When comparing two reactions, will the slower or quicker

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• temperature • surface area/particle size (of a solid) • a catalyst (including enzymes) Concentration Increasing the concentration of a reactant will increase the rate of reaction. When a piece of magnesium ribbon is added to a solution of hydrochloric acid, the following reaction occurs: magnesium + hydrochloric acid Mg(s) + 2HCl(aq)

→ magnesium chloride + hydrogen + H2(g) → MgCl2(aq)

As the magnesium and acid come into contact, there is effervescence (‘bubbling’) and hydrogen gas is given off. Two experiments were performed using the same length of magnesium ribbon, but different concentrations of acid. In experiment 1 the hydrochloric acid used was 2.0 mol/dm3, in experiment 2 the acid was 0.5 mol/dm3. The graph in Fig. 6.7 shows the results of the two experiments.

Experiment 1 finishes here

Experiment 2 finishes here

t1

80

rim

en

60

t2

en

pe

40

Ex

Volume of hydrogen gas / cm3

100

m eri

p

Ex

20

0

10

20

30 Time / seconds

40

50

60

∆ Fig. 6.7 Volume of hydrogen produced in the reaction between magnesium and hydrochloric acid.

SUPPLEMENT

In the 2.0 mol/dm3 hydrochloric acid solution there are more hydrogen ions in a given volume, a higher concentration of hydrogen ions, and so there will be a lot more effective collisions per second with the surface of the magnesium ribbon than in the 0.5 mol/dm3 hydrochloric acid.

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CHEMICAL REACTIONS

In experiment 1 the curve is steeper (has a higher gradient) than in experiment 2. In experiment 1 the reaction is complete after 20 seconds, whereas in experiment 2 it takes 60 seconds. The initial rate of the reaction is higher with 2.0 mol/dm3 hydrochloric acid than with 0.5 mol/dm3 hydrochloric acid.

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