Higher

Chemistry

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Higher

Exercise Book

1

Exercise 1 Standard grade revision - bonding

1. Use your data book to write down the number of electrons in the outer shell of an atom of the following

elements.

(a) hydrogen (b) oxygen (c) phosphorus (d) fluorine

2. The arrangement for the outer electrons of an oxygen atom is

shown left.

Draw similar diagrams to show the arrangement of the electrons

in the outer shell of an atom of the following elements.

(a) carbon (b) sulphur (c) nitrogen (d) bromine

3. An electron diagram to show the overlap of electron clouds in a

water molecule is shown left.

Draw a diagram to show electrons in the outer shells in the

following molecules.

(a) nitrogen hydride (NH3)

(b) hydrogen bromide (HBr)

(c) Silane (SiH4)

4. Copy and complete the table.

Particle No. of

protons

No. of

electrons

Charge

Na

Ca

Al

Na+

Ca2+

Al3+

5. Copy and complete the table

Particle No. of

protons

No. of

electrons

Charge

Cl

S

N

Cl−

S2−

N3−

6. Copy and complete the table

Particle No. of

protons

No. of

electrons

Charge

Cu2+

Pb4+

54 −

10 2-

18 2+

36 +

O

O

H

H

2

7. Predict the type of bonding (metallic, covalent, covalent network or ionic) for each chemical.

(i) water (ii) magnesium oxide

(iii) ammonia (iv) hydrogen chloride

(v) copper (vi) silicon carbide

(vii) iron (III) sulphate (viii) ammonium chloride

(ix) caesium fluoride (x) propane

(xi) cyclohexane (xii) manganese

(xiii) lead (IV) oxide (xiv) silicon

8. The three dimensional shape of a methane molecule is

shown left.

Draw similar diagrams to show the three-dimensional

shapes of the following molecules.

a) silane b) ammonia

c) hydrogen chloride d) water

9. Identify the following types of bonding from the following descriptions.

a) White solid that melts at 160°C; dissolves in water; does not conduct electricity.

b) White solid that melts at 801°C; dissolves in water; conducts electricity in solution (but not in the solid state).

c) Grey solid that melts at 1535°C; does not dissolve in water; conducts electricity in solid state.

d) White solid that melts at1610°C; does not dissolve in water; does not conduct electricity

e) A gas with a boiling point of –165°C; does not dissolve in water; does not conduct electricity.

11. Atoms can form ions with either a positive or negative charge.

A B C

LI

+

K

+

Mg2+

D E F

Cl

−

Br

−

O

2−

(a) Identify two ions with the same electron arrangement as an argon atom. ____________________

(b) Identify two ions, which combine to form an insoluble substance. ____________________

C

H

H H

H

3

Exercise 2 Standard grade revision - formulae and equations

1. Copy and complete the table.

Compound Formula Compound Formula hydrochloric acid nitric acid

sulphuric acid carbonic acid

ethanoic acid sodium hydroxide

calcium hydroxide barium hydroxide

magnesium oxide copper (II) oxide

aluminium oxide calcium carbonate

potassium carbonate lithium carbonate

2. Copy and complete the following to give balanced equations.

a) NaOH + HCl

b) H2SO4 + KOH

c) CuO + HCl

d) MgO + HNO3

e) CaCO3 + HCl

f) Na2CO3 + H2SO4

g) LiOH + HCl

h) H2SO4 + HNO3

i) PbO2 + HCl

j) Al2O3 + HNO3

Exercise 3 Standard grade revision - calculations

A. Titration calculations can be worked out in the following way:

Example: 40 cm3 of potassium hydroxide were neutralised by 10 cm

3 0f 2 mol/l sulphuric acid. What is the

concentration of the potassium hydroxide?

equation: 2KOH + H2SO4 K2SO4 + 2H2O

2 moles of KOH 1 mole of H2SO4

This means that there are twice as many moles of KOH as there are H2SO4

No. of moles H2SO4 = concentration x volume

= 2 x 0.01

= 0.02

∴ number of moles of KOH = 0.02 x 2

= 0.04

concentration of KOH = number of moles

volume

= 0.04

0.04

= 1 mol/l

1. Calculate the concentration of 20 cm3 sulphuric acid that would exactly neutralise

(a) 30 cm3 of 2 mol / l potassium hydroxide.

(b) 10 cm3 of 0.2 mol/l barium hydroxide.

10 cm3 = 0.01 l

40 cm3 = 0.04 l

4

2. Calculate the concentration of 10 cm3 hydrochloric acid that would exactly neutralise:

(a) 20 cm3 of 0.4 mol / l sodium hydroxide solution?

(b) 5.0 cm3 of 0.2 mol / l calcium hydroxide solution?

B. Calculations of mass from equations can be worked out in the following way:

Example: What is the mass of CO2 produced by burning 4.4g of C3H8?

Equation: C3H8 + 5O2 3CO2 + 4H2O

1 mole of C3H8 3 moles of CO2

Number of moles of C3H8 = Mass of C3H8 ÷R.F.M. of C3H8

= 4.4 ÷44

= 0.1

∴Number of moles of CO2 = 3 x 0.1

= 0.3

∴ Mass of CO2 = 0.3 x 44g

= 13.2g

3. What mass of carbon dioxide is produced by completely burning: -

(a) 36 g of carbon

(b) 1.5 g of ethane

4. What mass of mass of water is made by burning completely: -

(a) 6.4 g of methane

(b) 10 tonnes of ethane

(c) 30 g of hydrogen

Exercise 4 Standard grade revision - acids and alkalis

1. Ann added 20 cm3 0f 1mol/l sodium hydroxide solution to 20 cm

3 of 1mol/l sulphuric acid.

Identify the statement(s), which can be applied to this experiment.

A The number of H+ (aq) ions in the beaker decreased.

B The pH of the solution decreased.

C The number of SO42−

(aq) ions in the beaker decreased.

D Water molecules formed during the reaction.

E A precipitate formed during the reaction.

F The final solution contained equal numbers of H+(aq) and OH− (aq) ions.

2. Chromium (VI) oxide dissolves in water to form chromic acid.

CrO3 + H2O H2CrO4

(a) Suggest what is unusual about the reaction between chromium (VI) oxide and water.

(b) Chromic acid forms a salt with potassium hydroxide.

(i) How many moles of potassium hydroxide would be needed to neutralise completely 1 mole of chromic

acid?

(ii) Name the salt formed in the reaction. (You may wish to use page 7 of the data booklet to help you.)

Formula = C3H8

R.F.M. = 36 + 8

= 44

Formula = CO2

R.F.M. = 12 + 32

= 44

5

3. Many chemical compounds contain ions.

A B C

BaO

CaBr2

CuO

D E F

KCl

NaBr

SrCl2

Identify the two compounds, which are bases

4. Many chemical reactions involve water.

Identify the neutralisation reaction.

A C2H5OH C2H4 + H2O

B CH4 + 2O2 CO2 + 2H2O

C C2H2 + H2O C2H4O

D 2H2 + O2 2H2O

E C6H12O6 6C + 6H2O

F HNO3 + KOH KNO3 + H2O

5. The grid shows pairs of reactants.

A B C

Mg(s) + HCl(aq)

CuSO4 (aq) + Na2CO3(aq)

Fe(s) + CuSO4(aq)

D E F

NaOH(aq) + HCl(aq)

NH4NO3(s) + NaOH(aq) MgO(s)) + H2SO4(aq)

(a) Identify the pair(s) which is/are neutralisation reactions.

(b) Identify the reaction in which a precipitate is formed.

(c) Identify the reaction(s) in which a gas is formed.

6. The pH values of 1 mol/l solutions of some salts are shown in the table.

Salt pH

iron (II) sulphate

aluminium chloride

zinc (II) sulphate

copper (II) nitrate

sodium chloride

potassium sulphate

sodium carbonate

potassium carbonate

1

3

3

3

7

7

10

11

Identify the statement(s), which can be made about salts in the table.

A The salts are neutral.

B The salts of transition metals are acidic.

C The sulphates are acidic.

D The salts of group 1 metals are neutral.

E The salts of hydrochloric acid are neutral.

6

7. 2 mol/l hydrochloric acid is often used in the chemistry laboratory.

Identify the correct statement(s) about this solution.

A It produces hydrogen when electrolysed.

B It contains more H+(aq) ions than Cl−

(aq) ions.

C It does not react with magnesium.

D 20 cm3 is neutralised by 40 cm3 of 1 mol/l sodium hydroxide.

E 200 cm3 contains one mole of hydrogen chloride.

8. Chlorine reacts with water.

Cl2(g) + H2O(l) 2H+

(aq) + OCl−

(aq) + Cl−

(aq)

What happens to the pH of the water in a swimming pool when it is chlorinated?

9. Zinc reacts with sulphuric acid to produce hydrogen.

Adding copper sulphate solution is thought to speed up the

reaction.

You are given the following reactants:

1 mol/l sulphuric acid 1 mol/l copper sulphate zinc lumps

Describe how you would investigate the effect of adding copper sulphate solution on the speed of the

reaction between zinc and sulphuric acid.

10. A solution of sodium hydroxide used in photography has a concentration of 80g/l.

Calculate the concentration of sodium hydroxide in mol/l.

11. Magnesium sulphate can be converted to magnesium nitrate in 2 stages as follows.

Stage 1

MgSO4(aq) + Na2CO3(aq) Na2SO4(aq) + MgCO3(s)

Stage 2

MgCO3(s) + HNO3(aq) Mg(NO3)2(aq) + H2O(l) + CO2(g)

A B C

reduction

oxidation

neutralisation

D E F

precipitation

addition displacement

(a) Identify the type of reaction involved in Stage 1.

(b) Identify the type of reaction involved in Stage 2.

7

12. Solid sodium hydroxide often contains sodium carbonate as an impurity. This is caused by some of the

sodium hydroxide reacting with carbon dioxide from the air.

Instruction Card

Preparation of approximately 1 mol/l sodium hydroxide solution

CARE: Sodium hydroxide is highly corrosive.

1. Weigh out 41 g of sodium hydroxide pellets.

2. Dissolve in about 400 cm3 of water.

3. Add a few drops of barium chloride solution to remove

the carbonate ions.

4. Filter

5. Make the volume of solution up to exactly 1 litre.

The equation foe Step 3 is:

BaCl2(aq) + Na2CO3(aq) BaCO3(s) + 2NaCl(aq)

(i) Name an ion, which will be an impurity in the final sodium hydroxide solution.

(ii) Rewrite the equation as an ionic equation omitting spectator ions.

8

Exercise 5 Reaction rates (a)

1. Excess marble chips were added to 25 cm3 of 2 mol l

−1 hydrochloric acid.

Which measurement, taken at regular intervals and plotted against time, would give the graph shown below?

Measurement

taken

Time

A. Temperature B. Volume of gas produced C. pH of solution D. Mass of the beaker and contents

2. 1 g of magnesium ribbon was added to excess dilute hydrochloric acid. The mass of hydrogen gas produced was

measured and the results plotted in the graph below.

(a) From the graph estimate

(i) the total mass of gas produced (ii) time taken for the reaction to finish.

(b) Calculate the average rate of reaction for the first 20 seconds.

(c) Calculate the average rate of reaction for the period 20 seconds to 40 seconds.

(d) Explain the change in reaction rate as the reaction proceeds.

3. Some marble chips (excess) were added to 50 cm3 of 0.1 mol l−1 hydrochloric acid and the volume of carbon

dioxide produced was measured over a period of time. The results were as follows.

Time (seconds) 0 20 40 60 80 100

Volume of CO2 (cm3) 0 40 60 70 75 75

(a) Draw a diagram of the apparatus that could be used to make these measurements.

Mass of

hydrogen

gas

released

/ g

0.10

0.08

0.06

0.04

0.02

0

0 20 40 60 80 100

Time / s

9

Question 3 (continued)

(b) Draw a line graph of the measurements.

(c) Using your graph, calculate the average reaction rate for the first 30 seconds of the experiment.

(d) The experiment was repeated with excess calcium carbonate powder and 25 cm3 of 0.1 mol l

−1 hydrochloric acid.

Draw a dotted line on the above graph to show the graph you would expect to get with this second experiment.

4. Sodium thiosulphate was added to hydrochloric acid in a series of experiments. The time taken for the reaction

was recorded.

The results were recorded in a table:

Volume of

Na2S2O3 (aq)

used / cm3

Volume of

water added

/ cm3

Relative

concentration of

Na2S2O3 (aq)

Volume of

HCl (aq)

added / cm3

Time (t) for X

to vanish / s Rate / s−−−−1

A 50 0 5 10 13.5

B 40 10 4 10 16.9

C 30 20 3 10 22.5

D 20 30 2 10 33.8

E 10 40 1 10 67.5

(a) Describe how the time taken for the reaction was actually measured.

(b) Copy and complete the table by calculating the rate of reaction for experiments A to E.

(c) From the first two columns of the table, which factor is being kept constant throughout the series of

experiments?

(d) (i) What conclusion can be drawn concerning the concentration of the thiosulphate solution and reaction

rate?

(ii) Explain why changing the concentration affects the reaction rate.

(e) State one other way of increasing the reaction rate.

Exercise 6 Reaction rates (b)

1. The following potential energy diagram

represents the energy changes in a chemical

reaction.

The activation energy for the reaction in kJ

mol−1, is

A. 20 B. 30

B. 50 D. 70

reactant

products

20

40

70

Potential

energy /

kJ mol−1

Reaction

pathway

10

2. Which of the following is not a correct statement about the effect of a catalyst?

The catalyst

A. Provides an alternative route to the products B. Lowers the energy which molecules need for

successful collisions

C. Provides energy so that more molecules have

successful collisions

D. Forms bonds with reacting molecules.

3. When copper carbonate is reacted with excess acid,

carbon dioxide is produced. The curves shown opposite

were obtained under different conditions.

The change from P to Q could be brought about

by

A. increasing the concentration of the acid

B. decreasing the mass of copper carbonate

C. decreasing the particle size of the copper

carbonate

D. adding a catalyst.

4. For any chemical reaction, the temperature is a measure of

A. The average kinetic energy of the

particles which react

B. The average kinetic energy of all the

particles

C. The activation energy D. The minimum kinetic energy required

before a reaction occurs.

5. An experiment was carried out at four temperatures. The table shows the times taken for the reaction to occur.

Temperature / °C 20 30 40 50

Time / s 60 30 14 5

The results show that

A. The reaction endothermic B. The activation energy increases with increasing

temperature

C. The rate of reaction is directly proportional to the

temperature

D. A small rise in temperature results in a large

increase in reaction rate.

6. The diagram on the right could represent

A. the fermentation of glucose

B. neutralisation of an acid by an alkali

C. the combustion of sucrose

D. the reaction of a metal with acid.

7. The potential energy diagram opposite refers to

the reversible reaction involving reactants R and

products P.

What is the enthalpy change in kJ mol−1 for the

reverse reaction P R?

A. +30 B. +10

C. -10 D. -40

P

Q Volume

of CO2

Time

Rate of

reaction

Temperature

40

50

80

Potential

energy /kJ mol−1

Reaction pathway

11

8. The enthalpy change for the forward reaction in the

graph can be represented by

A. x B. y

C. x + y D. x – y

9.

A B C

D E F

Identify the graph, which shows how the rate varies with temperature in

(a) the reaction between sodium thiosulphate and dilute sulphuric acid

(b) the action of amylase on starch

10. The relative volumes of hydrogen produced in a given time for three reactions are plotted on the graph.

Curve I is for the reaction of excess zinc powder with 100 cm3 of 0.1 mol l

−1 hydrochloric acid.

A B C

Excess magnesium powder

50 cm3 of 0.1 mol l

−1

hydrochloric acid

50 cm3 of 0.2 mol l

−1

sulphuric acid

D E F

Excess iron powder

200 cm3 of 0.1 mol l

−1

hydrochloric acid

100 cm3 of 0.05 mol l

−1

sulphuric acid

(a) Identify the two chemicals that would react to give the results plotted in curve II.

(b) Identify the two chemicals that would react to give the results plotted in curve III.

rate

temperature

rate

temperature

rate

temperature

rate

temperature

rate

temperature

II I

III volume of

hydrogen

time

rate

temperature

Potential

energy

Reaction pathway

x

y

12

11. Many factors influence the rates of reactions.

A B C

Particle size of reactants temperature Surface area available for

reaction

D E F

Activation energy concentration Average kinetic energy of

reactant molecules

(a) Identify the factor which, if increased, causes an increase in the factor shown in box F.

(b) Identify the factor(s) which, if increased, would make the reaction slower.

12. Hydrogen peroxide can be used to clean contact lenses. In this process, the enzyme catalase is added to break

down hydrogen peroxide.

The equation for the reaction is:

2H2O2 2H2O + O2

The rate of oxygen production was measured in three laboratory experiments using the same volume of

hydrogen peroxide at the same temperature.

Experiment Concentration of H2O2 / mol l-1

Catalyst used

A 0.2 Yes

B 0.4 Yes

C 0.2 no

The curve obtained from experiment A is shown

20 40 60 80 100 120 140 160 180

time / s

(a) Calculate the average rate of the reaction over the first 40 s.

(b) Copy the graph (no graph paper needed) and add curves to the graph to show the results of experiments B

and C. Label each curve clearly.

13. Graph 1 shows the distribution of kinetic energies of

molecules in a gas at 30°C.

(a) Add a dotted line to Graph 1 to show the

distribution of kinetic energies at 20°C.

70

60

50

40

30

20

10

0

Volume of

oxygen

produced

/ cm3

A

Graph 2

Graph 1

Number of

molecules

Kinetic energy

13

Question 13 (continued)

(b) In Graph 2, the shaded area represents the

number of molecules with the required

activation energy EA, for the reaction to

occur.

Draw a line to show how a catalyst affects the energy of

activation.

14. Copy and complete the following table on industrial reactions and the catalysts used:

Catalyst Process Reaction Product(s)

Vanadium (V) oxide Contact sulphuric acid

Haber

4NH3 + 5O2 4NO + 6H2O

Breaking down long chain hydrocarbon

molecules

nickel Unsaturated oils + H2 saturated fats

Exercise 7 Reaction rates (c)

1. Magnesium ribbon reacts with dilute hydrochloric acid to form a salt, and hydrogen gas.

(a) Write a balanced equation for the above reaction.

(b) Indicate using labelled diagrams, two methods of following the course of the reaction. In each case, state

what measurements are necessary to determine the rate of the chemical reaction.

(c) Draw graphs (no graph paper required) of the expected results from each of the experiments described

in (b).

2. 5 g of marble chips were added to 50 cm3 of 1 mol l

−1 hydrochloric acid. State three ways in which the rate of

the chemical reaction can be increased.

3. Write a brief note explaining the difference between catalyst poisoning and catalyst regeneration.

4. Powdered silver metal behaves as a heterogeneous catalyst in the decomposition of ‘5 volume’ hydrogen

peroxide solution to give water and oxygen as the only products.

Note: Concentrations of peroxide solutions are quoted by ‘volume number’ with higher numbers relating to

solutions of greater concentrations.

(a) Write a balanced equation for the decomposition of hydrogen peroxide solution.

(b) What does a heterogeneous catalyst mean?

(c) Explain any observed differences in the rates of oxygen production, with time, in each of the following

pairs of silver metal/hydrogen peroxide solution reactions.

(i) 1 g silver powder in 250 cm3 ‘5 volume’ solution at 20 °C, and then at 40 °C.

(ii) 1 g of silver powder in 250 cm3 ‘5 volume’ solution at 20 °C, then with 250 cm

3 ’10 volume’

solution at the same temperature.

(iii) 250 cm3 ’10 volume’ solution at 30 °C with 1 g of silver powder, then with 1 g of silver foil at the

same temperature.

EA

Number

of

molecules

Kinetic energy

14

5. All new cars are fitted with a catalytic converter to deal with exhaust emissions.

(a) Name two metals which are used in the construction of catalytic converters in cars.

(b) Name three gases emitted by car exhausts, which are considered to be harmful.

(c) Explain how the catalyst changes these gases into harmless ones.

(d) Explain why cars fitted with a catalytic converter must use unleaded fuel.

6. The stomach produces an enzyme called pepsin. This helps digest food while it is in the stomach.

(a) Define the term ‘enzyme’.

(b) What are the two optimum conditions for the enzyme pepsin?

(c) Which type of food does pepsin digest?

7. Study the potential energy diagram for an

uncatalysed reaction.

(a) Using a dotted line, add a line to the

diagram to show energy change for a

catalysed reaction.

(b) Using the values shown on the diagram,

calculate the energy values for the

uncatalysed reaction.

(i) EA

(ii) ∆H

(c) What type of enthalpy reaction has accompanied this reaction? Explain.

(d) Explain why the catalyst speeds the rate of the reaction.

8. Copy and complete the following table to identify the functions of the enzymes.

Enzyme Reactions

zymase

papain

cellulase

sucrase

lipase

amylase

reactants

products

40

80

150

Potential

energy /

kJ mol −1

Reaction pathway

15

Exercise 8 Excess chemicals

1. 1.6 g of methane was mixed with 0.2 mol of oxygen gas and exploded. Which chemical was in excess?

2. 0.2 mol of magnesium was added to 100 cm3 of 0.5 mol l

−1 of hydrochloric acid. Which chemical was in

excess?

3. 0.06 g of sodium carbonate was added to 100 cm3 of 0.01 mol l

−1 of sulphuric acid. Which chemical was in

excess?

4. 1 g of propane was mixed with 10 g of oxygen and exploded. Calculate the mass of carbon dioxide produced.

5. 0.327 g of zinc metal was added to 100 cm3 of 0.02 mol l

−1 sulphuric acid. Calculate the mass of hydrogen

released.

6. 0.8 g of magnesium was added to 50 cm3 of 0.4 mol l−1 of hydrochloric acid. Which chemical was in excess and

hence calculate the mass of the remaining reactant.

7. A 1 g sample of marble (impure calcium carbonate) was added to excess nitric acid. The reaction produced 0.4 g

of carbon dioxide. Calculate the mass of the calcium carbonate that has reacted and hence, the percentage of

calcium carbonate in the marble.

Exercise 9 Enthalpy (a)

1. Copy and complete the following equation:

∆H = H (products)

2. (a) Complete the labelling of the energy diagram by identifying X and Y.

(b) Is this reaction exothermic or endothermic? Explain.

3. Redraw the energy diagram in question 2 and add a dotted line to show the reaction when a catalyst is used.

4. Write out the balanced equations to show the standard enthalpies (do not forget to show the state symbols):

(a) combustion of ethane

(b) solution of copper (II) sulphate

(c) addition of nitric acid to potassium hydroxide

Potential

energy

Reaction pathway

X

Y

16

5. The following is the equation for calculating the energy change in an experiment

Eh = c x m x ∆t

Explain what is meant by the symbols c, m and ∆∆∆∆T.

6. When 2 g of sodium hydroxide were added to water 1.68 kJ was liberated. Calculate the standard enthalpy of

solution of sodium hydroxide.

7. 6.93 kJ of heat was released when 0.24 g of methanol (CH3OH) completely burned. Calculate the standard

enthalpy of combustion of methanol.

8. When 100 cm3 of dilute nitric acid, concentration 0.1 mol l−1, is neutralised by 100 cm3 of 0.1 mol l−1 sodium

hydroxide solution, 570 J of heat is liberated. Calculate the enthalpy of neutralisation.

9. 9.24 kJ of heat was produced when 1 g of sulphur burned. Calculate the enthalpy of combustion of sulphur.

Exercise 10 Enthalpy (b)

1. When 0.1 g of propanol (C3H7OH) was burned, the heat produced raised the temperature of 100 cm3 water from

18°C to 23.2°C.

(a) Use the above data to calculate the enthalpy of combustion of propanol.

(b) A data booklet gives the value for the enthalpy combustion of propanol as ∆H = −2010 kJ mol−1

Explain the differences between this and your calculated value.

2. Discuss how the terms lattice-breaking and hydration are involved in the enthalpy of solution of an ionic solid.

3. Dissolving sodium hydroxide in 100 cm3 of water, releases 4.2 kJ of heat energy in the solution.

Use the data to calculate the rise in temperature.

4. The mixing of 50 cm3 of 0.5 mol l

−1 hydrochloric acid and 50 cm

3 of 0.5 mol l

−1 of sodium hydroxide produced a

temperature rise of 3.2°C. Calculate the enthalpy of neutralisation.

5. The following results are taken from a notebook of a pupil.

Experiment: Addition of 50 cm3 of 2.0 mol l−1 sodium hydroxide to 50 cm3 of 1.0 mol l−1 hydrochloric acid.

NaOH(aq) + HCl(aq) NaCl(aq) + H2O(l)

Mass of sodium hydroxide solution = 50 g

Mass of hydrochloric acid solution = 50 g

Initial temperature of NaOH(aq) = 21.9°C

Initial temperature of HCl(aq) = 21.5°C

Average temperature on mixing = X

Highest temperature during experiment = 28.5°C

(a) Work out which solution is in excess.

(b) Calculate the enthalpy of neutralisation of hydrochloric acid by sodium hydroxide.

17

6. When some propane was completely burned 222 kJ of heat was produced. If the enthalpy of

combustion of propane is –2220 kJ mol l−1

, what mass of propane was burned?

7. 0.01 moles of a hydrocarbon was completely burned to give 32.68 kJ of heat. Calculate the enthalpy

of combustion of the hydrocarbon and, using the data booklet, identify the hydrocarbon.

8. In an experiment 1.4 g of ethene were completely burned and the heat produced was 70.55 kJ. What

would be the temperature rise if the heat was absorbed by 500 cm3 of water.

Exercise 11 Periodic Table (a)

1. (a) What is meant by the first ionisation energy?

(b) (i) Is the process involved exothermic or endothermic?

(ii) Explain your answer.

(c) Write equations corresponding to the first ionisation energies of sodium and calcium and alongside each

equation write in the ∆H value with its appropriate sign.

2. Write equations corresponding to:

B. the second ionisation energy of magnesium

C. the third ionisation energy of aluminium

and alongside each equation write in the ∆H value with its appropriate sign.

3. (a) Explain why the second ionisation energy of each alkali metal is much greater than

that of the halogen in the same period.

(b) Explain why the third ionisation energy of magnesium is greater than that for aluminium.

(c) Why is there no value given in the data booklet for the fourth ionisation energy of lithium?

4. Calculate the energy required to bring about each of the following changes.

(a) Al(g) Al3+

(g) + 3e

(b) K+

(g) K3+

(g) + 2e

5. (a) Which three elements (in the first 20) have the lowest first ionisation energies?

(b) (i) In which group are these elements?

(ii) Explain why they have the lowest first ionisation energies?

6. (a) Which two elements (in the first 20) have the highest first ionisation energies?

(b) (i) In which group are these elements?

(ii) Explain why they have the highest first ionisation energies?

7. (a) Which elements show the greatest difference between their first ionisation energies

and their second ionisation energies?

(b) Explain your answer.

18

Exercise 12 Periodic Table (b)

1. Which element would require the most energy to convert 1 mole of gaseous atoms into gaseous ions carrying

two positive charges?

A. scandium B. titanium C. vanadium D. chromium

2. Which element has the greatest attraction for bonding electrons within a bond?

A. caesium B. oxygen C. fluorine D. iodine

3. Which equation represents the first ionisation enthalpy for chlorine?

A. Cl(g) + e Cl−

(g) B. Cl(g) Cl+

(g) + e

C. Cl+(g) + e Cl(g) D. Cl−

(g) Cl(g) + e

4. How does the pattern of density change going down group 7 of the Periodic Table? Explain your answer.

5. What is the trend in boiling points going down group 8 of the periodic table? Explain your answer.

6. Explain how the atomic size changes as you go across the periodic table from lithium to fluorine.

7. What is meant by a periodic property?

8. What is meant by the term electronegativity?

9. For each pair of elements identify the atom which has the greater electronegativity.

(a) oxygen and sulphur (b) strontium and calcium

(c) caesium and neon (d) lithium and iodine

Exercise 13 Bonding; structure and properties (a)

1 Decide whether each of the following molecules is polar or non-polar.

(a) oxygen

(b) Hydrogen chloride

(c) Carbon monoxide

(d) Hydrogen

2. By drawing the shape of and considering the symmetry, decide whether each of the following molecules is

polar.

(a) water (H2O)

(b) ammonia (NH3)

(c) methane (CH4)

(d) chloroform (CHCl3)

(e) methanol (CH3OH)

(f) carbon dioxide (CO2)

(g) carbon tetrachloride (CCl4)

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3. Hydrogen bonds will be set up between polar molecules, when these molecules contain a hydrogen atom

covalently bonded to an atom such as nitrogen, oxygen and fluorine.

Which of the following substances will exhibit hydrogen bonding?

(a) hydrogen bromide (HBr)

(b) propan–1–ol (CH3-CH2-CH2-OH)

(c) aminoethane (CH3-CH2-NH2)

(d) methoxymethane (CH3-O-CH3)

(e) tetrafluoromethane (CF4)

(f) propanone (CH3-CO-CH3)

(g) ethanoic acid (CH3-COOH)

4. Ammonia, NH3, exhibits a high degree of hydrogen bonding, yet there is no such bonding in hydrogen.

(a) Explain why this so.

(b) How do the strengths of the forces involved in hydrogen bonding compare with the strengths of van der

Waals’ forces?

5. Analysis of hydrogen fluoride shows the existence of molecules with relative molecular masses of 20, 40 and 60.

Explain the origin of these molecules.

Exercise 14 Bonding; structure and properties (b)

1. Taking water as an example, describe what happens when a covalent bond is formed between the atoms.

Illustrate your answer.

2. Draw shapes of the following molecules:

(a) methane (b) ammonia (c) water (d) hydrogen fluoride

3. Explain what is meant by the following terms:

(a) Ionic bond (b) Dipole (c) van der Waals forces (d) Hydrogen bond

4. State what metallic bonding means. Illustrate your answer.

5. (a) For each chemical, state whether the type of bonding is metallic, covalent, covalent network, polar or

ionic.

Chemicals

(i) water (ii) methane (iii) sodium

chloride

(iv) methanol (v) silicon (vi) carbon

tetrachloride

(vii) potassium

fluoride

(viii) hydrogen

gas

(ix) copper

(x) copper

oxide

(xi) aluminium

chloride

(xii) aluminium

fluoride

(xiii) carbon

dioxide

(xiv) ammonia (xv) ethene

(b) Which of the above chemicals contain non-polar molecules even though the bonding is polar covalent.

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6. When the following solids melt, what forces of attraction are overcome by the heating?

(a) sodium metal (b) hydrogen gas

(c) water (d) pentane

(e) propene (f) potassium chloride

(g) hydrogen fluoride (h) silicon dioxide

7. Explain why propanone has a higher boiling point than butane even though they have the same relative

molecular mass.

8. Explain why sodium chloride has a melting point of 800°C while that of aluminium chloride has a melting point

of 192°C.

Exercise 15 Bonding; structure and properties (c)

1 The diagram illustrates the trend in the boiling points of Group 7 hydrides HCl, HBr and HI.

250

200

Boiling

Point/K 150

100

50

0 HF HCl HBr HI

(a) Why is the boiling point of HI higher than that of HBr?

(b) Estimate the value for the boiling point of hydrogen fluoride (HF) if it were to follow the trend in the

graph.

(c) The actual boiling point of HF is 292 K

Account for this unexpectedly high value.

2. Liquid ammonia boils at − 33°C but liquid phosphine, PH3, boils at – 87.5°C.

Explain this difference in terms of bonding.

3. A pupil is comparing the boiling points of an alkane with ethanol, C2H5OH, to examine the influence of

hydrogen bonding.

(a) To examine the influence of hydrogen bonding, the boiling point of which alkane should be compared

with that of ethanol?

(b) Explain your answer.

(c) Which of the two compounds has hydrogen bonding between the molecules?

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4. Arrange the molecules below in order of increasing viscosity.

CH3CH2OH CH2 CH CH2 CH2 CH2

OH OH OH OH OH

ethanol propane – 1,2,3 – triol ethane – 1,2, - diol

5. Taking sodium chloride as your example, explain, in terms of bonding, what happens to an ionic crystal when it

dissolves in water.

6. Explain the following:

(a) Potassium chloride dissolves readily in water but not in cyclohexane.

(b) tetrachloromethane dissolves in hexane but not in water.

7. Lithium iodide is moderately soluble in non-polar solvents whereas caesium fluoride is not.

On the basis of this evidence, deduce the difference in bonding in the two compounds.

Exercise 16 The Mole (a)

1. Calculate the mass of each of the following:

(a) 0.5 mol of sulphur dioxide, SO2.

(b) 5 mol of calcium nitrate, Ca(NO3)2.

(c) 2 mol of sodium carbonate.

(d) 0.1 mol of copper (II) carbonate.

2. Calculate the number of moles which are contained in each of the following:

(a) 14g of nitrogen, N2.

(b) 202.2 g of potassium nitrate.

(c) 6.605 g of ammonium sulphate

(d) 10.2 g of aluminium oxide.

3. How many moles of sodium hydroxide must be dissolved to make each of the following solutions?

(a) 200 cm3 of 1 mol l

−1

(b) 500 cm3 of 0.5 mol l

−1

(c) 100 cm3 of 2 mol l−1

4. What is the concentration of each of the following solutions of hydrochloric acid?

(a) 2 mol of hydrogen chloride dissolved in 200 cm3 of solution.

(b) 0.5 mol of hydrogen chloride dissolved in 250 cm3 of solution.

(c) 0.1 mol of hydrogen chloride dissolved in 500 cm3 of solution.

5. What is the volume of each of the following solutions sodium hydroxide?

(a) 2 mol l−1 solution which contains 0.4 mol solute?

(b) 0.5 mol l−1 solution which contains 0.1 mol solute?

(c) 0.1 mol l−1

solution which contains 0.2 mol solute?

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6. How many grams of each of the following substances must be weighed out to make up each of the following

solutions?

(a) 500 cm3 of 0.1 mol l−1 magnesium sulphate, MgSO4.

(b) 100 cm3 of 2 mol l

−1 sodium nitrate, NaNO3.

(c) 2 litres of 0.1 mol l−1

sodium hydroxide, NaOH.

7. What is the concentration of each of the following solutions?

(a) 13.82 g of potassium carbonate, K2CO3, dissolved in 2 litres of solution.

(b) 15.96 g of copper (II) sulphate dissolved in 250 cm3 of solution.

(c) 10 g of ammonium nitrate dissolved in 100 cm3 of solution.

(d) 3.65 g of hydrogen chloride dissolved in 250 cm3 of solution.

Exercise 17 The Mole (b) 1. Decide which has more atoms:

(a) 16 g of sulphur or 27 g of aluminium

(b) 2.3 g of sodium or 2.3 g of potassium

(c) 10 g of argon or 10 g of neon

2. Decide which has more molecules:

(a) 28 g of carbon monoxide or 32 g of sulphur dioxide

(b) 32 g of oxygen or 1 g of hydrogen

(c) 6.4 g of sulphur dioxide or 6.4 g of carbon dioxide

3. Decide which has more atoms:

(a) 64 g of sulphur dioxide or 28 g of carbon monoxide

(b) 4.4 g of carbon dioxide or 8 g of argon

(c) 20 g of neon or 2 g of hydrogen

4. Decide which has more hydrogen atoms:

(a) 17 g of ammonia or 16 g of methane

(b) 3 g of ethane or 3 g of water

(c) 0.2 g of hydrogen or 2 g of hydrogen fluoride

5. Decide which has more ions:

(a) 16 g of copper (II) sulphate or 58 g of magnesium hydroxide

(b) 48 g of ammonium carbonate or 35 g of potassium hydroxide

(c) 45 g of sodium sulphate or 45 g of potassium nitrate

6. Calculate the number of atoms in each of the following elements.

(a) 10 g of calcium

(b) 24 g of carbon

(c) 2.43 g of magnesium

(d) 4 g of argon

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7. Calculate the number of molecules in each of the following compounds.

(a) 1.1 g of carbon dioxide

(b) 34 g of ammonia

(c) 360 g of water

(d) 3.2 g of methane

8. Calculate the number of formula units in each of the following compounds.

(a) 9.42 g of potassium oxide

(b) 14.8 g of calcium hydroxide

(c) 5.85 g of sodium chloride

(d) 3.038 g of iron (II) sulphate

Exercise 18 The Mole (c)

1. Decide which has the greater volume, measurements being made at the same temperature and pressure.

(a) 17 g of ammonia or 17 g of methane

(b) 20 g of neon or 20 g of nitrogen

(c) 4 g of hydrogen or 44 g of carbon dioxide

(d) 6.4 g of sulphur dioxide or 8 g of oxygen

2. The volume of 0.22 g of propene is 118 cm3.

Calculate the volume of 2 mol of propene at this temperature and pressure.

3. Calculate the volume of 2.4 g of ethene. (molar volume = 23.6 litres mol−1

)

4. The volume of 1 g of hydrogen is 11.6 litres.

Calculate the volume of 4 mol of hydrogen at this temperature and pressure.

5. Using the densities given in the data booklet, calculate the volume of 10 g of:

(a) hydrogen

(b) argon

6. 3 g of an alkane occupies a volume of 2.24 litres.

What is the molecular formula for the alkane? (molar volume = 22.4 litres mol−1

)

7. Using the densities given in the data booklet, calculate the number of moles in 10 litres of:

(a) helium

(b) nitrogen

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Exercise 19 The Mole (d)

1. From the following data calculate the approximate formula of the hydride gas X.

Mass of plastic bottle (empty) = 112.80 g

Mass of plastic bottle + gas X = 113.52 g

Capacity of plastic bottle = 1 litre

Molar volume of gas X = 23.6 litres mol-1

2. A flask, capacity 600 cm3, was used to calculate the molar volume of sulphur dioxide.

The following data was obtained

Mass of evacuated flask = 512.97 g

Mass of flask + sulphur dioxide = 514. 57 g

Calculate the volume of 1 mol of sulphur dioxide (under the experimental conditions).

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

What volume of hydrogen gas is produced when 5 g of magnesium is added to 100 cm3 of 2 mol l−1 hydrochloric

acid? (Molar volume = 23.6 litres mol−1.)

3. What volume of oxygen is needed to completely burn 4 g of sulphur to form sulphur dioxide? (Use the density

of oxygen given in the data booklet page 3.)

4. What mass of magnesium oxide is obtained when excess magnesium is ignited in 5 litres of oxygen? (Molar

volume = 22.8 litres mol−1

.)

5. What mass of water is produced when 2 litres of hydrogen is burned in excess oxygen? (Use the density of

hydrogen given in the data booklet page 3.)

6. 2 litres of hydrogen is ignited with 2 litres of oxygen. Calculate the mass of water produced. (Use the density of

hydrogen and oxygen given in the data booklet page 3.)

7. Calculate the volume of oxygen required for the complete combustion of 1 g of ethene. (molar volume = 23.2

litres mol−1

.)

8. Using the density of hydrogen given in the data booklet, what volume of hydrogen gas is produced when 13 g of

zinc is added to excess acid?

9. The first stage in the extraction of copper from sulphide ores such as chalcopyrite (CuFeS2) involves heating the

ore with sand and oxygen.

4 CuFeS2(s) + 2SiO2(s) + 5O2(g) 2Cu2S.FeS(l) + 2FeSiO3(l) + 4SO2(g)

Calculate the volume of oxygen required to react completely with 1472 kg chalcopyrite. (Molar volume = 24.8

litres mol−1.)

11. The flow diagram shows how benzene (C6H6) can be obtained from naphtha.

hydrogen

naphtha hexane benzene

C6H6 C6H14 catalytic reforming

Calculate the volume of hydrogen produced when 1 kg of hexane is reformed. (Use the density of hydrogen

given in the data booklet page 3.)

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Exercise 20 The Mole (e)

1. In each of the following reactions decide the ratio of the volume of products to the volume of reactants.

(a) H2(g) + Cl2(g) 2HCl(g)

(b) N2(g) + 3H2(g) 2NH3(g)

(c) 2C(s) + O2(g) 2CO(g)

(d) C2H4(g) + 3O2(g) 2CO2(g) + 2H2O(l)

(e) CuO(s) + CO(g) Cu(s) + CO2(g)

2. N2(g) + 2O2(g) NO2(g)

What volume of nitrogen dioxide is produced when 100 cm3 of nitrogen is sparked in excess oxygen?

3. When each of the following volumes of gases is burned completely, calculate

(a) the volume of oxygen required

(b) the volume of carbon dioxide produced.

(i) 100 cm3 of methane

(ii) 2 litres of carbon dioxide

(iii) 250 cm3 of ethene

(iv) 150 cm3 of propane

4. 10 cm3 of butane gas is mixed with 75 cm3 of oxygen and the mixture exploded.

C4H10(g) + 13/2 O2(g) 4CO2(g) + 5H2O(l)

Calculate the volume and composition of the resulting gas mixture.

5. A mixture of 60 cm3 of hydrogen and 40 cm3 of carbon monoxide is passed over hot copper (II) oxide until no

further reaction occurred. Calculate the volume of the resulting gas.

6. The reaction of sulphur dioxide with oxygen to produce sulphur trioxide does not go to completion. Assuming

only 50% of the available sulphur dioxide is converted to sulphur trioxide, calculate the volume and composition

of the resulting gas mixture when 100 cm3 of sulphur dioxide is mixed with 100 cm

3 of oxygen.

7. If 100 cm3 of propene is burned completely with 900 cm3 of oxygen, what will be the composition of the

resulting mixture?

8. 50 cm3 of ethyne, C2H2, is completely burned in 220 cm

3 of oxygen.

(a) What will be the volume and composition of the resulting gas mixture?

(b) What will be the volume if the reaction is repeated in a container heated to 200°C?

9. A mixture of 80 cm3 CO and 150 cm

3 of O2 is exploded. After cooling, the residual gas is shaken with sodium

hydroxide solution.

(a) Which gas is absorbed by the sodium hydroxide solution?

(b) What is the reduction in volume of the residual gas on shaking with the solution?

(c) What volume of gas remains?