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Higher
Chemistry
A
Higher
Exercise Book
Exercise 1
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
2.
Standard grade revision - bonding
(b) oxygen
(c) phosphorus
(d) fluorine
The arrangement for the outer electrons of an oxygen atom is
shown left.
O
Draw similar diagrams to show the arrangement of the electrons
in the outer shell of an atom of the following elements.
(a) carbon
3.
(b) sulphur
(c) nitrogen
(d) bromine
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.
4.
(a)
nitrogen hydride (NH3)
(b)
hydrogen bromide (HBr)
(c)
Silane (SiH4)
O
H
H
Copy and complete the table.
Particle
No. of
protons
No. of
electrons
Charge
No. of
protons
No. of
electrons
Charge
No. of
protons
No. of
electrons
Charge
54
10
18
36
−
22+
+
Na
Ca
Al
Na+
Ca2+
Al3+
5.
Copy and complete the table
Particle
Cl
S
N
Cl−
S2−
N3−
6.
Copy and complete the table
Particle
Cu2+
Pb4+
1
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
8.
(xiv) silicon
H
The three dimensional shape of a methane molecule is
shown left.
C
Draw similar diagrams to show the three-dimensional
shapes of the following molecules.
H
H
H
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
LI+
B
K+
C
Mg2+
D
Cl−
E
Br−
F
O2−
(a)
Identify two ions with the same electron arrangement as an argon atom.
____________________
(b)
Identify two ions, which combine to form an insoluble substance.
____________________
2
Exercise 2
1.
Standard grade revision - formulae and equations
Copy and complete the table.
Compound
Formula
Compound
hydrochloric acid
sulphuric acid
ethanoic acid
calcium hydroxide
magnesium oxide
aluminium oxide
potassium 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
A.
Formula
nitric acid
carbonic acid
sodium hydroxide
barium hydroxide
copper (II) oxide
calcium carbonate
lithium carbonate
Standard grade revision - calculations
Titration calculations can be worked out in the following way:
Example:
40 cm3 of potassium hydroxide were neutralised by 10 cm3 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
∴ number of moles of KOH
concentration of KOH
=
concentration x volume
=
2 x 0.01
=
0.02
=
0.02 x 2
=
0.04
=
number of moles
volume
=
=
1.
10 cm3 = 0.01 l
40 cm3 = 0.04 l
0.04
0.04
1 mol/l
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.
3
Calculate the concentration of 10 cm3 hydrochloric acid that would exactly neutralise:
2.
B.
(a)
20 cm3 of 0.4 mol / l sodium hydroxide solution?
(b)
5.0 cm3 of 0.2 mol / l calcium hydroxide solution?
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
Formula = C3H8
R.F.M. = 36 + 8
= 44
∴Number of moles of CO2
Formula = CO2
R.F.M. = 12 + 32
= 44
3.
4.
(a)
36 g of carbon
(b)
1.5 g of ethane
=
4.4 ÷44
=
0.1
=
3 x 0.1
=
0.3
=
0.3 x 44g
=
13.2g
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
Standard grade revision - acids and alkalis
Ann added 20 cm3 0f 1mol/l sodium hydroxide solution to 20 cm3 of 1mol/l sulphuric acid.
Identify the statement(s), which can be applied to this experiment.
A
B
C
D
E
F
2.
Mass of C3H8 ÷R.F.M. of C3H8
What mass of carbon dioxide is produced by completely burning: -
Exercise 4
1.
∴ Mass of CO2
=
The number of H+ (aq) ions in the beaker decreased.
The pH of the solution decreased.
The number of SO42−(aq) ions in the beaker decreased.
Water molecules formed during the reaction.
A precipitate formed during the reaction.
The final solution contained equal numbers of H+(aq) and OH− (aq) ions.
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.)
4
3.
Many chemical compounds contain ions.
A
B
C
BaO
CaBr2
D
CuO
E
F
KCl
NaBr
SrCl2
Identify the two compounds, which are bases
4.
Many chemical reactions involve water.
Identify the neutralisation reaction.
A
B
C
D
E
F
5.
C2H5OH
CH4 + 2O2
C2H2 + H2O
2H2 + O2
C6H12O6
HNO3 + KOH
C2H4 + H2O
CO2 + 2H2O
C2H4O
2H2O
6C + 6H2O
KNO3 + H2O
The grid shows pairs of reactants.
A
B
Mg(s) + HCl(aq)
D
C
CuSO4 (aq) + Na2CO3(aq)
E
NaOH(aq) + HCl(aq)
Fe(s) + CuSO4(aq)
F
NH4NO3(s) + NaOH(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
iron (II) sulphate
aluminium chloride
zinc (II) sulphate
copper (II) nitrate
sodium chloride
potassium sulphate
sodium carbonate
potassium carbonate
MgO(s)) + H2SO4(aq)
pH
1
3
3
3
7
7
10
11
Identify the statement(s), which can be made about salts in the table.
A
B
C
D
E
The salts are neutral.
The salts of transition metals are acidic.
The sulphates are acidic.
The salts of group 1 metals are neutral.
The salts of hydrochloric acid are neutral.
5
7.
8.
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.
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)
A
Mg(NO3)2(aq) + H2O(l) + CO2(g)
B
reduction
D
C
oxidation
E
precipitation
(a)
Identify the type of reaction involved in Stage 1.
(b)
Identify the type of reaction involved in Stage 2.
neutralisation
F
addition
displacement
6
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.
7
Exercise 5
Reaction rates (a)
Excess marble chips were added to 25 cm3 of 2 mol l−1 hydrochloric acid.
1.
Which measurement, taken at regular intervals and plotted against time, would give the graph shown below?
Measurement
taken
Time
A. Temperature
2.
B. Volume of gas produced
C. pH of solution
D. Mass of the beaker and contents
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.
0.10
0.08
Mass of
hydrogen
gas
released
/g
0.06
0.04
0.02
0
0
20
40
60
80
100
Time / s
(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)
Volume of CO2 (cm3)
(a)
0
0
20
40
40
60
60
70
80
75
100
75
Draw a diagram of the apparatus that could be used to make these measurements.
8
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
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
Rate / s−1
(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)
(e)
(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.
State one other way of increasing the reaction rate.
Exercise 6
1.
Reaction rates (b)
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
70
Potential
energy /
kJ mol−1
40
products
20
reactant
Reaction
pathway
9
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.
P
The change from P to Q could be brought about
by
4.
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.
Time
For any chemical reaction, the temperature is a measure of
A.
C.
5.
Q
Volume
of CO2
The average kinetic energy of the
particles which react
The activation energy
B.
The average kinetic energy of all the
particles
The minimum kinetic energy required
before a reaction occurs.
D.
An experiment was carried out at four temperatures. The table shows the times taken for the reaction to occur.
20
60
Temperature / °C
Time / s
30
30
40
14
50
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
7.
Rate of
reaction
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.
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
Temperature
80
Potential
energy
/kJ mol−1
50
40
Reaction pathway
10
8.
The enthalpy change for the forward reaction in the
graph can be represented by
x
Potential
energy
A.
x
B.
y
C.
x+y
D.
x–y
y
Reaction pathway
9.
A
B
C
rate
rate
rate
temperature
temperature
D
temperature
E
F
rate
rate
rate
temperature
temperature
temperature
Identify the graph, which shows how the rate varies with temperature in
10.
(a)
the reaction between sodium thiosulphate and dilute sulphuric acid
(b)
the action of amylase on starch
The relative volumes of hydrogen produced in a given time for three reactions are plotted on the graph.
II
I
III
volume of
hydrogen
time
Curve I is for the reaction of excess zinc powder with 100 cm3 of 0.1 mol l−1 hydrochloric acid.
A
B
50 cm3 of 0.1 mol l−1
hydrochloric acid
Excess magnesium powder
D
C
E
Excess iron powder
50 cm3 of 0.2 mol l−1
sulphuric acid
F
3
200 cm of 0.1 mol l
hydrochloric acid
−1
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.
11
11.
Many factors influence the rates of reactions.
A
B
C
Particle size of reactants
temperature
D
Surface area available for
reaction
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.
Concentration of H2O2 / mol l-1
0.2
0.4
0.2
Experiment
A
B
C
Catalyst used
Yes
Yes
no
The curve obtained from experiment A is shown
70
60
Volume of
oxygen
produced
/ cm3
50
A
40
30
20
10
0
20
40
60
80
100
120
140
160
180
time / s
13.
(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.
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.
Graph 1
Number of
molecules
Graph 2
Kinetic energy
12
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.
Number
of
molecules
Draw a line to show how a catalyst affects the energy of
activation.
Kinetic energy
14.
EA
Copy and complete the following table on industrial reactions and the catalysts used:
Catalyst
Process
Vanadium (V) oxide
Contact
Reaction
Product(s)
sulphuric acid
Haber
4NH3 + 5O2
4NO + 6H2O
Breaking down long chain hydrocarbon
molecules
nickel
Unsaturated oils + H2
Exercise 7
1.
saturated fats
Reaction rates (c)
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 cm3 ’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.
13
5.
6.
7.
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.
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?
Study the potential energy diagram for an
uncatalysed reaction.
(a)
(b)
Using a dotted line, add a line to the
diagram to show energy change for a
catalysed reaction.
Using the values shown on the diagram,
calculate the energy values for the
uncatalysed reaction.
150
Potential
energy /
80
kJ mol −1
40
8.
(i)
EA
(ii)
∆H
products
reactants
Reaction
pathway
(c)
What type of enthalpy reaction has accompanied this reaction? Explain.
(d)
Explain why the catalyst speeds the rate of the reaction.
Copy and complete the following table to identify the functions of the enzymes.
Enzyme
Reactions
zymase
papain
cellulase
sucrase
lipase
amylase
14
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
1.
Enthalpy (a)
Copy and complete the following equation:
∆H
2.
(a)
=
H (products)
Complete the labelling of the energy diagram by identifying X and Y.
X
Potential
energy
Y
Reaction pathway
(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
15
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
1.
Enthalpy (b)
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 cm3 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)
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.
H2O(l)
16
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
1.
(a)
What is meant by the first ionisation energy?
(b)
(i)
Is the process involved exothermic or endothermic?
(ii)
Explain your answer.
(c)
2.
Periodic Table (a)
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.
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.
4.
(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?
Calculate the energy required to bring about each of the following changes.
(a)
Al(g)
Al3+(g)
+
3e
+
3+
+
2e
(b) K
5.
6.
7.
K
(g)
(g)
(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?
(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?
(a)
Which elements show the greatest difference between their first ionisation energies
and their second ionisation energies?
(b)
Explain your answer.
17
Exercise 12
1.
Which element would require the most energy to convert 1 mole of gaseous atoms into gaseous ions carrying
two positive charges?
A.
2.
scandium
B.
titanium
C.
vanadium
D.
chromium
D.
iodine
Which element has the greatest attraction for bonding electrons within a bond?
A.
3.
Periodic Table (b)
caesium
B.
oxygen
C.
fluorine
Which equation represents the first ionisation enthalpy for chlorine?
A.
C.
Cl(g)
Cl
+
(g)
+
+
Cl−(g)
e
e
B.
Cl(g)
D.
Cl(g)
−
Cl
(g)
Cl+(g) +
e
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
1
2.
Bonding; structure and properties (a)
Decide whether each of the following molecules is polar or non-polar.
(a)
oxygen
(b)
Hydrogen chloride
(c)
Carbon monoxide
(d)
Hydrogen
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)
18
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?
4.
5.
(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)
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?
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)
3.
methane
(b)
ammonia
(c)
water
(d)
hydrogen fluoride
Explain what is meant by the following terms:
(a)
Ionic bond
(b)
Dipole
(c)
van der Waals forces
4.
State what metallic bonding means. Illustrate your answer.
5.
(a)
(d) Hydrogen bond
For each chemical, state whether the type of bonding is metallic, covalent, covalent network, polar or
ionic.
Chemicals
(i)
water
(ii)
methane
(iv)
methanol
(v)
silicon
(vii) potassium
fluoride
(x)
copper
oxide
(xiii) carbon
dioxide
(b)
(viii) hydrogen
gas
(xi) aluminium
chloride
(xiv) ammonia
(iii)
sodium
chloride
(vi) carbon
tetrachloride
(ix) copper
(xii) aluminium
fluoride
(xv) ethene
Which of the above chemicals contain non-polar molecules even though the bonding is polar covalent.
19
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
1
Bonding; structure and properties (c)
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?
20
4.
Arrange the molecules below in order of increasing viscosity.
CH3CH2OH
ethanol
CH2
CH
CH2
CH2
CH2
OH
OH
OH
OH
OH
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)
(b)
7.
Potassium chloride dissolves readily in water but not in cyclohexane.
tetrachloromethane dissolves in hexane but not in water.
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
1.
2.
3.
4.
5.
The Mole (a)
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.
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.
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
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.
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?
21
6.
7.
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.
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
1.
2.
3.
4.
5.
6.
The Mole (b)
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
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
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
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
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
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
22
7.
8.
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
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
1.
2.
The Mole (c)
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
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.
4.
Calculate the volume of 2.4 g of ethene. (molar volume = 23.6 litres mol−1)
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.
6.
Using the densities given in the data booklet, calculate the volume of 10 g of:
(a)
hydrogen
(b)
argon
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
23
Exercise 19
1.
2.
The Mole (d)
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
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
C6H6
C6H14
benzene
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.)
24
Exercise 20
1.
2.
The Mole (e)
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)
N2(g)
+
2O2(g)
NO2(g)
What volume of nitrogen dioxide is produced when 100 cm3 of nitrogen is sparked in excess oxygen?
3.
4.
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
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 cm3 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 cm3 of oxygen.
9.
(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?
A mixture of 80 cm3 CO and 150 cm3 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?
25