Rate of Reactions


Rate of reaction/Speed of reaction: It is the speed for a reactant to be used up or product to be formed.

2 ways to measure speed of reaction

1. Measuring time for reaction to complete

    • Speed of reaction is inversely proportional to time taken; the shorter the time needed for reaction to complete, the faster the speed of reaction is.

Speed of reaction =1/time taken

Speed of reaction A = = 0.333/s

Speed of reaction B = = 0.667/s

Therefore reaction B is faster than reaction A as time taken for B is shorter

Number of times B faster than A = = 2 times

2. Measuring the amount of product produced in a period of time or measuring the amount of reactant remain in a period of time.

    • Can be measured by plotting change in volume of gas evolved, mass of reaction mixture as reaction proceeds and change of pressure of gas formed.

A. Measuring the amount of gas evolved.

    • Consider reaction of limestone with acid to produce carbon dioxide.

    • A syringe is used to help in measurement of gas produced in volume every time interval.

    • A graph of volume of gas against time is plotted.

      • Gradient largest at start indicating speed at its greatest.

      • Gradient decreases with time – speed decreases with time.

      • Gradient becomes zero, speed is zero. The reaction has finished.

B. Measuring change in mass of reaction mixture.

    • Marble is reacted with acid in a flask with cotton wool stucked at top to prevent splashing during reaction but it allows gas to be free.

    • The reading on balance is plotted on a graph on every time interval.

Factors Affecting Speed of Reaction

1. Particle Size of Reactant

    • When large marble is reacted with acid and compared to reaction of fine marble solids being reacted with acid and the graph of volume of gas against time is plotted, it is found that the reaction involving finer marble chips produces gas faster than the one with larger marble chunk as the graph of finer chips is steeper.

    • The volume of gas at the end is the same for both reactions.

    • Therefore, reactions of solids with liquid/gas is faster when the solids are of smaller pieces

    • Reactions occur when particles collide.

    • Small particles creates larger surface area for more collisions between reacting particles which increases speed of reaction.

    • Explosions: chemical reactions occuring extremely rapid rate producing heat+gas

    • Examples

      • Coal dust burn faster than large pieces as it has larger surface area. In coal mines, when air contains too much coal dust, explosion can occur from a single spark or match. Water is sprayed into the air to remove coal dust.

      • Flour in mills can ignite easily due to large surface area.

2. Concentration of Reactant

    • In the increase of concentration means there are more solute particles per unit volume of the solution which favours for more effective collision resulting in an increase in speed of reaction.

3. Pressure of Reactant

    • Only gaseous reactions are affected as gas is compressible.

    • At higher pressure, molecules are forced to move closely together, hence increasing the particles per unit volume of gas and effectively increases the collision between reacting molecules so the speed of reaction increases.

    • High pressure is used in industrial processes (e.g. Haber Process Plant) so that the reaction goes faster.

4. Temperature of Reaction

    • Speed of reaction increases when temperature increases.

    • Particles don’t always react upon collision but just bounce as they don’t have enough activation energy to react.

    • With increase in temperature, particles absorb the energy and having enough activation energy, they move faster and collide more effectively per second.

    • Therefore, speed of reaction is increased.

    • Usually, speed of reaction doubles for every 10oC rise in temperature.

5. Effect of Catalyst

    • Catalysts are chemical substances which alters speed of reaction without itself being used at the end of a reaction.

    • It can be reused and only small amount of catalyst is needed to affect a reaction.

    • transition metals (e.g. Titanium, Nickel, Iron, Copper) are good catalysts

    • most catalyst catalyse one kind of reaction (except titanium)

    • Catalysts lower the need of energy to break bonds so activation energy is lower.

      • Consequently, bond breaking occurs easily and more often when particles collide

Factors Affecting Speed of Catalysed Reactions:

Speed of catalysed reactions can be increased by:

    • increasing temperature

    • increasing concentration of solutions

    • increasing pressure of gas reactions

Catalyst provide “alternative path” which results in lower activation energy.


Enzymes are biological catalysts

Characteristics of enzymes:

    • They are very specific. One enzyme catalyse one type of reaction.

    • Sensitive to temperature. They work best at 40oC. Too high or too low temperatures destroy enzymes.

    • Sensitive to pH. They function within narrow range of pH.

Industrial uses of enzymes:

    • They are added to detergents from bacteria, and also to make tough meat tender. These enzymes can be found in papaya fruit.

    • Yeast convert sugars into alcohol and carbon dioxide by fermentation. Beer, wine and soy sauce are made this way.

    • Fungal enzymes can be used to make antibiotics such as penicillin.

Exothermic Reaction

    • Exothermic change is one which heat energy is given out.

    • Reaction is written as:

      • Reactants → Products + heat (or)

      • Reactants → Products [ΔH = – n kJ], where n is amount of heat energy released

Examples of exothermic changes

1. Changes of State

    • When gas condenses to water or water freezes to solid, heat is given out.

    • Eg. Condensation of steam to water

      • H2O (g) → H2O (l) + heat

2. Combustion reactions

    • All combustion (burning) reactions are exothermic.

    • Eg. Burning of hydrogen in air

      • 2H2 (g) + O2 (g) → 2H2O (l) + heat

3. Dissolving of anhydrous salts/acids in water

    • Dissolving solid salt to aqueous solution of the salt gives out heat

    • Eg. Dissolving of Na2CO3 in water (or CuSO4)

      • Na2CO3 (s) → Na2CO3 (l) + heat

    • Eg. Dissolving of concentrated acid in water

      • HCl (aq) + H2O (l) → less concentrated HCl (aq) + heat

4. Neutralization

    • When acid and alkali react it gives out heat due to combining of H+ ions from acid and OH- ions from alkali to form water

      • H+ (aq) + OH- (aq) → H2O (l) + heat

5. Metal Displacement

    • Magnesium reacting with copper(II) sulphate

    • Mg (s) + Cu2+ (aq) → Mg2+ (s) + Cu (s) + heat

Endothermic Reaction

    • Endothermic change is one which heat energy is absorbed.

    • This is to break bonds between the reactants which needs more energy in them.

    • Reaction is written as:

      • Reactants + heat → Products (or)

      • Reactants → Products [ΔH = + n kJ], where n is amount of heat energy absorbed

Examples of endothermic changes:

1. Changes of states

    • When solid melts to water & boils to steam, heat is absorbed to break the bond.

    • Eg. Condensation of steam to water

      • H2O (s) + heat → H2O (l)

2. Photolysis

    • Reaction of light sensitive silver chloride in camera reel in light

      • 2AgBr (s) + heat → 2Ag (s) + Br2 (g)

3. Dissolving of Ionic Compounds

    • Ionic compounds such as NH4Cl, KNO3, CaCO3 absorb heat from surroundings.

      • Eg. NH4Cl (s) + heat → NH4Cl (aq)

      • Eg. CuSO4 (s) + heat → CuSO4 (aq)

4. Photosynthesis

    • Light energy is absorbed by plants to produce starch.

5. Decomposition by heat

    • Many compounds require heat for decomposition, e.g. CaCO3 to CO2 + CaO

      • CaCO3 (s) + heat → CO2 (g) + CaO (s)

6. Acid + Bicarbonates (HCO3)

    • NaHCO3 (s) + H2SO4 (aq) + heat → NaSO4 (aq) + CO2 (g) + H2O (l)

Endothermic vs Exothermic Reactions

Heat of Reaction

    • The amount of energy given out or absorbed during a chemical reaction is enthalpy change.

    • The symbol is ΔH measured in kilojoules (kJ).

Exothermic reaction:

    • Mg (s) + CuSO4 (aq) → MgSO4 (aq) + Cu (s) [ΔH = –378 kJ]

    • 378 kJ of heat energy is given out when 1 mol of Mg react with 1 mol CuSO4 to produce 1 mol of MgSO4 and 1 mol of Cu.

Endothermic reaction:

    • CaCO3 (s) → CO2 (g) + CaO (s) [ΔH = +222 kJ]

    • 222 kJ of heat energy is absorbed when 1 mol of CaCO3 decompose to 1 mol of CO2 and 1 mol of CaO.

Heat Energy and Enthalpy Change in Reaction

    • When bonds are formed, heat energy is given out, it’s exothermic and ΔH is negative

    • When bonds broken, heat energy is absorbed, it’s endothermic and ΔH is positive

Worked Example

Hydrogen bromide is made by reacting H2 gas with Br2 gas. Calculate the heat change of the reaction given the equation and bond energy table below.

H2 (g) + Br2 (g) → 2HBr (g)

Bonds of H2 and Br2 molecules must be broken first to make HBr.

Heat energy is absorbed to break these bonds by endothermic reaction.

H – H + Br – Br → H H Br Br

Broken bonds are used to make H – Br bonds of HBr. Heat energy is released.

H H Br Br → 2H – Br

Heat change can be calculated by:

ΔH = heat released in making bonds + heat absorbed in breaking bonds

Exothermic ΔH

= the bond energy of 2 H – Br bonds

= 2(366)

= – 732 kJ

Endothermic ΔH

= the bond energy of 1 H – H bond + 1 Br – Br bond

= 436 + 224

= + 660 kJ

ΔH = – 732 + 660 = – 72 kJ

Therefore more heat is given out in making bond than absorbed in breaking bond.

The overall change is to give out heat and it’s exothermic with ΔH negative.

Activation energy

    • Activation energy is the minimum energy needed to start a reaction.

    • It is the energy needed to break the reactant bonds before new bonds are formed.

    • Reactions occur because of collision of particles and sufficient kinetic energy is needed to provide activation energy to break the bonds and start the reaction by providing extra energy from a heat source.

Exothermic and Endothermic Reaction Graph

    • In exothermic reaction, enough energy given out in the reaction of particles to provide activation energy therefore less energy is needed to form products.

    • In endothermic reaction, insufficient energy is given out when bonds are made to provide activation energy for reaction to continue.

    • More energy is needed to form products and heat must be continually added to fulfill energy requirement.


    • The combustion of fuels gives out large amount of energy in industries, transport & homes.

    • These fuel mainly methane from coal, wood, oil, natural gas & hydrogen.

    • Combustion in air provides energy and gives out heat. Hence, exothermic reaction.

Hydrogen as a Fuel

    • Hydrogen provides twice as much as heat energy per gram than any other fuel and burns cleanly in air to form steam.

    • They are mainly used as rocket fuel.

Production of Hydrogen

    • Hydrogen is produced either by electrolysis of water or by cracking of hydrocarbon

By cracking of hydrocarbon:

    • First, methane (hydrocarbon) and steam are passed over a nickel catalyst to form hydrogen and carbon monoxide.

      • CH4 (g) + H2O (g) --> CO (g) + 3H2 (g)

    • The by-product carbon monoxide is not wasted. It is reacted with more steam to form carbon dioxide and hydrogen.

      • CO (g) + H2O (g) --> CO2 (g) + H2 (g)

    • Now you get more hydrogen.

By electrolysis:

    • Water is electrolysed according to equation:

    • 2H2O (l) --> 2H2 (g) + O2 (g)

    • However, electrolysis is costly.

Creation of the Fuel

In Engines:

    • The hydrogen created is reacted with oxygen to form steam and heat energy

    • 2H2 (g) + O2 (g) --> H2O (g) + heat

This heat is needed to thrust the vehicle forward. However, we don’t use heat energy for our daily appliances.

    • Instead we use electrical energy and to make electrical energy from hydrogen, we use fuel cell

Fuel Cells

    • A fuel cell converts chemical energy directly into electrical energy.

    • Hydrogen reacts with hydroxide ions into electrolyte on the platinum catalyst on electrode to make the electrode negatively-charged.

      • H2 + 2OH- --> 2H2O + 2e-

    • Electrons flows past the load and to the other electrode. That negatively-charged electrode is now anode. Hydroxide ions constantly deposit electrons here to make water. While then, the other electrode is now cathode.

    • Oxygen reacts with water created on from hydrogen on the cathode to gain electrons from it:

      • O2 + 2H2O + 4e- --> 4OH-

    • If we combine the ionic equations, we still get water as product of hydrogen and oxygen, but the energy produced is now electrical energy:

      • 2H2 (g) + O2 (g) --> H2O (g) + electrical energy


    • Petroleum is a mixture of hydrocarbons, which are compounds made up of carbon and hydrogen only.

    • Crude oil, freshly extracted from underground, undergo refining – a process where oil undergoes fractional distillation to be separated into its fractions.

    • First, crude oil is heated up to 350oC and the vapours rise up a tower, divided with trays on some certain heights for the fractions to be collected.

    • The fractionating column is cooler on top, hence upper trays collects fractions of low boiling points while the lower ones, being hotter, collect those with higher boiling points.


    • Plants take in carbon dioxide and water in presence of chlorophyll and synthesize them in the presence of sunlight to produce glucose and release oxygen:

      • 6CO2 + 6H2O --> C6H12O6 + 6O2

    • Plants get their energy by using the glucose formed. Scientists believe that we can use the stored energy in glucose as combustible fuels.

    • First, glucose fermented to make ethanol by microorganisms such as yeast. This is fermentation. The glucose is usually derived from corn plant or sugar cane.

      • C6H12O6 → 2C2H6O + 2CO2

    • Then, water is removed from ethanol by fractional distillation by heating it up until 78oC (boiling point of ethanol).

    • Some water might still be present as the boiling point is close to ethanol. The ethanol produced is then mixed with fuel to be combusted to produce energy. This is biofuel, and it’s a renewable energy source.

MCQ Questions

1. If a strip of magnesium is dropped into excess hydrochloric acid an exothermic reaction occurs. The rate of reaction increases during the first few seconds because

a. the amount of magnesium is decreasing

b. the magnesium is acting as a catalyst

c. the solution is becoming hotter

d. the surface area of the magnesium is increasing

2. The energy profile diagram show how adding a substance to a reaction mixture changes the reaction pathway.

Which change is likely to be observed when X is added to the reaction mixture?

a. the reaction becomes less exothermic

b. the reaction becomes more exothermic

c. the speed of the reaction decreases

d. the speed of the reaction increases

3. Dilute sulphuric acid reacts with copper(II) oxide to form copper(II) sulphate and water. What would not alter the rate of this reaction?

a. the concentration of the sulphuric acid

b. the pressure at which the reaction takes place

c. the size of the particles of copper(II) oxide

d. the temperature of the reacting mixture

4. Zinc reacts with an excess of dilute sulphuric acid. The graph shows how the volume of hydrogen gas given off changed with time.

Why does the graph become horizontal at X?

a. all the sulphuric acid has reacted

b. all the zinc has reacted

c. hydrogen is being produced at a constant rate

d. the reaction is beginning to slow down

5. Why is the reaction H2 + Cl2 --> 2HCl exothermic?

a. energy involved in the bonds breaking is greater than that of the bonds forming

b. energy involved in the bonds forming is greater than that of the bonds breaking

c. more bonds are broken than are formed

d. more bonds are formed than are broken

6. Which of the following is an endothermic process?

a. the addition of water to anhydrous copper(II) sulphate

b. the combustion of ethanol in air

c. the formation of a carbohydrate and oxygen from carbon dioxide and water

d. the oxidation of carbon to carbon dioxide

7. Which process is exothermic?

a. burning petrol in a car engine

b. cracking of oil fractions

c. fractional distillation of oil

d. melting bitumen for roads

8. Curve I is obtained by observing the decomposition of 100cm3 of 1mol/dm3 aqueous hydrogen peroxide, catalysed by manganese(IV) oxide.

2H2O2 (aq) --> 2H2O (l) + O2 (g)

Which alteration to the conditions will produce curve II?

a. adding some 0.1 mol/dm3 aqueous hydrogen peroxide

b. lowering the temperature

c. using a better catalyst

d. using less manganese(IV) oxide

9. Why is Vanadium(IV) oxide used in the oxidation of sulphur dioxide to sulphur trioxide?

a. it acts as a reducing agent

b. it prevents the decomposition of sulphur trioxide

c. it removes impurities

d. it speeds up the reaction

10. Nitrogen and hydrogen react in a closed vessel.

N2 (g) + 3H2 (g) <---> 2NH3 (g)

How do the speeds of the forward and reverse reactions change, if the pressure in the vessel is increased but the temperature is kept constant?

11. In the graph, curve X represents the results of the reaction between 1.0 g of granulated zinc and an excess of acid at 30oC.

Which changes will produce curve Y?

a. using 1.0 g of powdered zinc at 20oC

b. using 1.0 g of granulated zinc at 20oC

c. using 0.5 g of granulated zinc at 40oC

d. using 0.5 g of granulated zinc at 20oC

12. Ethanol is produced by the fermentation of sugar. During the reaction, carbon dioxide is given off. The graph shows how the volume of carbon dioxide produced per minute varies with temperature.

Using the graph, decide which statement is correct?

a. the rate of reaction always increases with temperature

b. the rate of reaction reaches a maximum at about 40

c. the reaction is slowest at 0

d. the reaction takes a long time to begin

13. Which change will increase the speed of the reaction between 1 mol of each of the two gases?

a. a decrease in surface area of the catalyst

b. a decrease in temperature

c. a decrease in the volume of the reaction flask

d. an increase in the volume of the reaction flask

14. Magnesium reacts with hydrochloric acid. Which solution would give the fastest initial rate of reaction?

a. 40g of HCl in 1000cm3 of water

b. 20g of HCl in 1000cm3 of water

c. 10g of HCl in 100cm3 of water

d. 4g of HCl in 50cm3 of water

15. In which reaction is the pressure least likely to affect the rate of reaction

a. C (s) + CO2 (g) ---> 2CO (g)

b. 2SO2 (g) + O2 (g) ---> 2SO3 (g)

c. N2 (g) + 3H2 (g) ---> 2NH3 (g)

d. NaOH (aq) + HCl (aq) ---> NaCl (aq) + H2O (l)

16. Which of these changes is exothermic?

a. evaporation

b. thermal decomposition

c. respiration

d. melting

17. Which of these changes is endothermic?

a. freezing

b. neutralisation

c. photosynthesis

d. combustion

18. Which element is always present in fuels like coal, oil, and natural gas?

a. hydrogen

b. carbon

c. oxygen

d. nitrogen

19. Natural gas burns more easily than other fuels because

a. it is more exothermic

b. it is a gas

c. it is colourless

d. it has a low density

20. Most fuels contain small amounts of sulphur. When they burn, a gas which pollutes the atmosphere and causes acid rain is formed. it is called

a. carbon monoxide

b. carbon dioxide

c. sulphur dioxide

d. sulphur trioxide

21. Methane (CH4) is the main constituent of natural gas. When it burns, it produces 890 kJ of heat per mole. How much heat would be produced if 64 g of methane were burnt?

[Ar of C = 12, Ar of H = 1]

a. 890 kJ

b. 1780 kJ

c. 2670 kJ

d. 3560 kJ

22. Which of these reactions would you expect to be endothermic?

a. 4K (s) + O2 (s) --> 2K2O (s)

b. H2 (g) --> 2H (g)

c. 2NaOH (aq) + H2SO4 (aq) --> Na2SO4 (aq) + 2H2O (l)

d. 2H2 (g) + O2 (g) --> 2H2O (l)

23. An endothermic reaction is one in which

a. the reaction vessel gets hot

b. light is given out

c. the products contain more energy than the reactants

d. chemical bonds are made

24. Which pair of elements, described by their proton number, will react together most exothermically?

a. 2 and 8

b. 12 and 16

c. 3 and 10

d. 19 and 9

25. A fuel is a device for

a. converting electrical energy efficiently into chemical energy

b. converting chemical energy efficiently into electrical energy

c. recharging accumulators

d. burning a fuel efficiently

26. The formation of hydrogen iodide from hydrogen and iodine is an endothermic reaction.

H-H + I-I ---> H-I + H-I

What may be deduced from this information?

a. The number of bonds broken is greater than the number of bonds formed.

b. The formation of H-I bonds absorbs energy.

c. The products possess less energy than the reactants.

d. The total energy change in bond formation is less than that in bond breaking.

27. Methane gas reacts extremely slowly with air at room temperature. If a piece of warm platinum is held in a methane-air mixture, the methane ignites. What differences are there between the reaction with the platinum and the reaction without the platinum?

For the reaction with the platinum:

I. The activation energy is lower.

II. The energy change is greater.

III. The energy of the reactants is higher.

IV. The rate of reaction is greater.

a. I and II only

b. I and III only

c. I and IV only

d. II and IV only

28. The graph shows how the total volume of hydrogen produced changes when iron fillings reacted with excess dilute sulphuric acid.

Which statement best describes the section PQ of the curve?

a. The acid is slowly used up which results in the reaction slowing down.

b. The decreasing mass of the iron filings results in the reaction slowing down.

c. Water is produced in the reaction that dilutes the acid which slows down the reaction.

d. Hydrogen gas produced slows down the reaction.

29. Two experiments were carried out in which hydrochloric acid was added to limestone.

Experiment A: 500 cm3 of 1.0 mol/dm3 hydrochloric acid added to an excess of limestone.

Experiment B: 100 cm3 of 5.0 mol/dm3 hydrochloric acid added to an excess of limestone.

The initial rate of evolution of carbon dioxide and the total volume of carbon dioxide evolved were measured in each experiment. How do the results in experiment A compare with those in experiment B when all other conditions are identical?

Rate of evolution of carbon dioxide Total volume of carbon dioxide

a. It is slower in A than in B. It is the same in A and B.

b. It is faster in B than in A. It is greater in B than in A.

c. It is slower in B than in A. It is greater in B than in A.

d. It is the same in A and B. It is greater in A than in B.

30. Why is the reaction H2 + Cl2 --> 2HCl exothermic?

a. Energy involved in the bonds breaking is greater than that of the bonds forming.

b. Energy involved in the bonds forming is greater than that of the bonds breaking.

c. More bonds are broken than are formed.

d. More bonds are formed than are broken.

MCQ Answers

1. c

2. d

3. b

4. b

5. b

6. c

7. a

8. a

9. d

10. a

11. c (since the total vol of H2 collected is about half of X, the mass of Zn used should be halved. The steeper gradient of Y suggests a faster rate of reaction eg higher temperature)

12. b

13. c

14. c

15. d (there are no gaseous reactants in the reaction)

16. c

17. c

18. b

19. a

20. c

21. d

22. b

23. c

24. d

25. b

26. d

27. c

28. b

29. a

30. b

Worked Solutions

1. A metallic element forms compounds in which its oxidation states are II and III. The element is displaced from solutions of its salts by copper.

a. Using the symbol El for the element, write the formulae for the chlorides and oxides of this element.

chlorides ______ and ________

oxides _______ and _________

b. Why is it necessary for the symbols of the majority of the elements to consist of two letters rather than one only?

c. Would you expect this element El to react with dilute sulphuric acid? Explain your answer.

d. Write an equation to show the reduction of an ion of this element El from oxidation state III to oxidate state II


a. chlorides: ElCl2 and ElCl3

oxides: ElO and El2O3

b. There are elements with names that start with the same letter eg. copper (Cu) and carbon (C)

c. No. Because copper does not react with dilute sulphuric acid and El is less reactive than copper.

d. 2El3+ (aq) + Cu (s) --> 2El2+ (aq) + Cu2+ (aq)

2. Give one physical property and one chemical property possessed by all metals.


a. good conductor of electricity/heat

b. form oxides which show basic properties

3. At 150oC and a pressure of one atmosphere, the reversible reaction between gas A and gas B reaches a dynamic equilibrium.

A (g) + 2B (g) <---> AB2 (g) : heat change = -220kJ/mol

a. Is the formation of AB2 exothermic or endothermic? Explain your answer.

bi. What is meant by the phrase dynamic equilibrium?

bii. Predict how the proportion of AB2 at equilibrium changes if the pressure is increased.

c. What effect will an increase in temperature have on the rate of formation of AB2? Explain your answer in terms of the movement of the molecules.


a. exothermic. This is because the forward reaction has a negative heat change, indicating that heat is released.

bi. The amounts of A, B and AB2 are constant (equilibrium) but the forward and backward reactions are still taking place (dynamic).

bii. amount of AB2 increases

c. Rate is faster. At higher temperature, the molecules have more kinetic energy. More molecules possess energy greater than the required minimum energy for a reaction to take place. Furthermore, they collide more frequently with one another. Hence, the rate of reaction increases.

4. The element oxygen exists in two forms, O2 and O3. Both forms are gases.

ai. Ozone molecules O3 decompose when heated into O2 molecules. Construct the equation for this decomposition.

aii. What volume of oxygen O2 is formed when 40cm3 of ozone O3 is decomposed, both volumes being measured at the same temperature and pressure?

b. Ozone is an oxidizing agent. Describe the color change you expect to see when ozone is bubbled into aqueous potassium iodide.

color before:

color after:


ai. O3 (g) --> 3/2 O2 (g)

aii. 60 cm3

bi. colorless

bii. brown

5a. Ammonia is made in the Haber process by the reversible reaction between nitrogen and hydrogen.

i. How is nitrogen obtained from liquid air for use in this process?

ii. State the name of the catalyst and the conditions used in the Haber process.

b. Ethanol can be manufactured by reacting ethene and steam in the presence of phosphoric acid as a catalyst.

The reaction is reversible and forms an equilibrium mixture.

C2H4 (g) + H2O <--> C2H5OH (g)

i. Predict how increasing the pressure will change the percentage of ethanol present at equilibrium. Explain your answer.

ii. The table shows how the percentage of ethanol present at equilibrium changes with temperature at a pressure of 60 atmospheres.

Is the formation of ethanol exothermic or endothermic? Explain your answer.

c. One use of ethanol is the manufacture of an acid. Name the acid and draw its structural formula.


ai. fractional distillation

aii. iron(III) oxide catalyst at 450oC and 200 atm

bi. Increasing pressure will shift the equilibrium to the right where there is a reduction in total amount of gases since 2 moles of gases combine to give only 1 mole. Hence, the percentage of ethanol increases.

bii. As temperature increases, amount of ethanol produced decreases, showing that the reverse reaction is favored. Hence, the backward reaction is endothermic where heat is absorbed. Therefore, the formation of ethanol is exothermic.

c. ethanoic acid

6. The table below shows some bond energies, measured in kilojoules per mole. Bond energy is the energy required to break the bonds between pairs of atoms

a. Which of the bonds listed above is the strongest?

b. Is the double bond between two carbon atoms twice as strong as a single bond? Use the info given above to explain your answer

c. Use the info given to calculate the total energy required to break one mole of methane into atoms.

CH4 --> C + 4H; ∆H = ? kJ

d. Bond-making is an exothermic process (∆H, -ve)

Complete the statements below to calculate the energy change expected from the reaction of one mole of hydrogen with one mole of chlorine to form two moles of hydrogen chloride

H2 + Cl2 --> 2HCl; ∆H = ? kJ

Energy change in breaking the bonds in one mole of H2 = ______kJ

Energy change in breaking the bonds in one mole of Cl2 = ______kJ

Total energy change = ______kJ

Energy change in making the bonds in two moles of HCl = ______kJ

Hence ∆H for this reaction = ______kJ


6a. C = C bond

6b. No. If C = C bond is twice as strong as a C - C bond, then the C = C bond energy should be 2 x 348 = 696 kJ/mol, but the bond energy of C = C bond is actually less than that.

6c. 4 x 412 = 1648 kJ

6di. 436 kJ

6dii. 242 kJ

1diii. 436 + 242 = 678 kJ

1div. -(2 x 431) = - 862 kJ

1dv. 678 + (- 862) = - 184 kJ

7. In the Contact process for the manufacture of sulphuric acid, sulphur dioxide is converted into sulphur dioxide.

2SO2 (g) + O2 (g) <---> 2SO3 (g) ∆H = - 98 kJ/mol

a. The reaction reaches a dynamic equilibrium. Explain the term.

b. A mixture of sulphur dioxide, oxygen, and sulphur trioxide was allowed to reach equilibrium, then the temperature was raised. Predict the effect of raising the temperature on

i. the composition of the equilibrium mixture

ii. the rate of reaction

Explain your answers.

c. Describe how you would prepare a pure dry sample of sodium sulphate starting with dilute sulphuric acid.


7a. At dynamic equilibrium for a reversible reaction, the forward rate of reaction is equal to the backward rate of reaction, and both rates are not equal to zero.

7bi. Amount of SO3 decreases. When temperature is raised, the equilibrium shifts to the left where heat is absorbed (backward reaction is endothermic) so as to counteract the temperature change.

7bii. Faster rate of reaction. Since temperature is raised, both forward and backward rates of reaction are increased.

7c. First, perform a titration on 25.0 cm3 of aqueous NaOH, with dilute sulphuric acid using a suitable indicator such as phenolphthalein. The volume of acid used is noted at the end point when the colour of the indicator changes. Repeat the experiment but without the indicator. The final solution will be sodium sulphate solution. Crystalise sodium sulphate by evaporation. The crystals are filtered, washed with water, and dried with filter paper and dessicator.

8. In the future, fuel cells may be used to power cars. In a fuel cell, the overall reaction is represented by the equatin

2H2 (g) + O2 (g) --> 2H2O (l)

a. This is the energy profile diagram for the above reaction

8aiii. The amount of energy released when covalent bonds in the water molecules are formed is greater than the energy required to break covalent bonds in the hydrogen and oxygen molecules. Hence, there is a net amount of energy released.

8b. Titanium, because transition metals are good catalysts.

9. Methane, CH4, is used as a fuel. The complete combustion of methane can be represented by the equation below

CH4 + 2O2 -------> CO2 + 2H2O

a. Explain why this reaction is exothermic in terms of the energy changes that take place during bond breaking and bond forming.

b. Calculate the energy released when 4.0g of methane is completely combusted.

c. Draw the energy profile diagram for the complete combustion of methane.


9a. The reaction is exothermic because the heat of reaction, ∆H, has a negative value. The energy released during formation of covalent bonds in products is greater than the energy absorbed to break covalent bonds in reactants. Hence, the overall reaction is exothermic. A total of 890 kJ of energy is released.

9b. Mr of CH4 = 16

No. of mol of CH4 = 4/16 = 0.25

Energy released = 0.25 x 890 = 225 kJ


i. Label on the diagram the activation energy of the reaction.

ii. The fuel cell contains a catalyst. Draw a second curve on the diagram to show the energy profile for the catalysed reaction.

iii. Explain why this reaction is exothermic in terms of bond breaking and bond forming.

b. Choose from the following list the metal that is most likely to act as a catalyst. Give a reason for your answer.







10. Methane is a fuel. It completely burns to form carbon dioxide and water. When 1 mole of methane is burned, 890 kJ of energy is released.

a. Calculate the energy released when 0.32g of methane is burned.

b. Use ideas of bond breaking and bond forming to explain why the reaction exothermic.


10a. 0.32g of methane contains 0.32/16 = 0.020 moles of methane

energy released = 0.02 x 890 = 17.8 kJ

10b. The energy released by formation of bonds in carbon dioxide and water is far greater than the energy absorbed to break bonds in methane and the oxygen molecules.

11. Hydrogen is used as a fuel in some space rockets.

a. Describe how hydrogen is manufactured from a named hydrocarbon source.

b. A highly exothermic reaction occurs between hydrogen as oxygen to form water.

i. Suggest why liquid hydrogen, rather than hydrogen gas is carried by space rockets.

ii. Explain what is meant by an exothermic reaction.

iii. Which bonds are broken in this reaction.

iv. What type of energy change occurs when bonds are formed?

v. When 1 mole of hydrogen molecules reacts with oxygen, the energy ÔH = -285kJ

Calculate the energy change which occurs when 100g of hydrogen reacts with oxygen.


11a. Hydrogen is formed from the reaction between steam and methane.

H2O (g) + CH4 (g) --> CO (g) + 3H2 (g)

bi. Liquid hydrogen occupies a smaller volume and so is easier to transport.

bii. Exothermic reactions are those where heat energy is released in the course of the reaction.

biii. H-H and O=O bonds

biv. Heat is released. Chemical energy is converted into heat energy, sometimes with the emission of light.

bv. 100g of H contains 100/2 = 50 moles of H

Energy changes = 50(-285) = -14250kJ