Download CfE Higher Chemistry Unit 1: Chemical Changes and Structure

SCHOLAR Study Guide
CfE Higher Chemistry
Unit 1: Chemical Changes and
Structure
Authored by:
Emma Maclean (George-Heriot’s School)
Reviewed by:
Diane Oldershaw (Menzieshill High School)
Previously authored by:
Peter Johnson
Brian T McKerchar
Arthur A Sandison
Heriot-Watt University
Edinburgh EH14 4AS, United Kingdom.
First published 2014 by Heriot-Watt University.
This edition published in 2016 by Heriot-Watt University SCHOLAR.
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SCHOLAR Study Guide Unit 1: CfE Higher Chemistry
1. CfE Higher Chemistry Course Code: C713 76
ISBN 978-1-909633-20-9
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i
Contents
1 Reaction rates - collision theory
1.1 Prior knowledge . . . . . . .
1.2 Introduction . . . . . . . . .
1.3 Rate of reaction . . . . . . .
1.4 Collision theory . . . . . . .
1.5 Summary . . . . . . . . . . .
1.6 Resources . . . . . . . . . .
1.7 End of topic test . . . . . . .
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1
2
3
4
6
10
10
10
2 Reaction rates - reaction profiles
2.1 Prior knowledge . . . . . . .
2.2 Interpreting graphs . . . . .
2.3 Activation energy . . . . . .
2.4 Summary . . . . . . . . . . .
2.5 Resources . . . . . . . . . .
2.6 End of topic test . . . . . . .
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17
19
19
25
30
30
31
3 Catalysis
3.1 Prior knowledge . . . . . .
3.2 Introduction to catalysis . .
3.3 Catalyst mechanism . . .
3.4 Catalysts in industry . . . .
3.5 Potential energy diagrams
3.6 Catalysts and energy . . .
3.7 Summary . . . . . . . . . .
3.8 Resources . . . . . . . . .
3.9 End of topic test . . . . . .
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35
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67
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5 Bonding and structure
5.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
5.2 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
83
85
86
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4 The Periodic Table
4.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . .
4.2 Arrangement of elements in the Periodic Table: Introduction
4.3 History of the Periodic Table . . . . . . . . . . . . . . . . . .
4.4 Trends and patterns (periodicity) . . . . . . . . . . . . . . . .
4.5 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
4.6 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . .
4.7 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . .
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ii
CONTENTS
5.3
5.4
5.5
5.6
5.7
5.8
Metallic bonds .
Covalent bonds
Ionic bonds . .
Summary . . . .
Resources . . .
End of topic test
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6 Periodic Table trends
6.1 Prior knowledge . . . . . . . . . . . . . .
6.2 Covalent radius . . . . . . . . . . . . . .
6.3 Ionisation energies . . . . . . . . . . . .
6.4 Electronegativity . . . . . . . . . . . . . .
6.5 Summary of trends in the Periodic Table
6.6 Summary . . . . . . . . . . . . . . . . . .
6.7 Resources . . . . . . . . . . . . . . . . .
6.8 End of topic test . . . . . . . . . . . . . .
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7 Bonding continuum and polar covalent bonding
7.1 Prior knowledge . . . . . . . . . . . . . . . . .
7.2 Polar covalent bonds . . . . . . . . . . . . . .
7.3 Predicting bonding type using electronegativity
7.4 Polar molecules . . . . . . . . . . . . . . . . .
7.5 The bonding continuum . . . . . . . . . . . . .
7.6 Summary . . . . . . . . . . . . . . . . . . . . .
7.7 Resources . . . . . . . . . . . . . . . . . . . .
7.8 End of topic test . . . . . . . . . . . . . . . . .
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92
102
107
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108
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113
115
116
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128
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130
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137
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153
8 Intermolecular forces
8.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . .
8.2 Introduction . . . . . . . . . . . . . . . . . . . . . . . .
8.3 London dispersion forces . . . . . . . . . . . . . . . . .
8.4 Hydrogen bonding . . . . . . . . . . . . . . . . . . . . .
8.5 Relating properties to bonding . . . . . . . . . . . . . .
8.6 Viscosity . . . . . . . . . . . . . . . . . . . . . . . . . .
8.7 Predicting solubilities from solute and solvent polarities
8.8 Summary . . . . . . . . . . . . . . . . . . . . . . . . . .
8.9 Resources . . . . . . . . . . . . . . . . . . . . . . . . .
8.10 End of topic test . . . . . . . . . . . . . . . . . . . . . .
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157
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178
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180
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9 End of unit test
183
Glossary
193
Answers to questions and activities
1
Reaction rates - collision theory . . . . . . . . .
2
Reaction rates - reaction profiles . . . . . . . .
3
Catalysis . . . . . . . . . . . . . . . . . . . . .
4
The Periodic Table . . . . . . . . . . . . . . . .
5
Bonding and structure . . . . . . . . . . . . . .
6
Periodic Table trends . . . . . . . . . . . . . . .
7
Bonding continuum and polar covalent bonding
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196
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201
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213
218
© H ERIOT-WATT U NIVERSITY
CONTENTS
8
9
Intermolecular forces . . . . . . . . . . . . . . . . . . . . . . . . . . . .
End of unit test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
© H ERIOT-WATT U NIVERSITY
iii
221
224
1
Topic 1
Reaction rates - collision theory
Contents
1.1
1.2
1.3
1.4
1.5
1.6
1.7
Prior knowledge
Introduction . .
Rate of reaction
Collision theory
Summary . . .
Resources . . .
End of topic test
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2
3
4
6
10
10
10
Prerequisite knowledge
Before you begin this topic, you should already know that:
•
reactions can be monitored by measuring changes in concentration, mass and
volume of reactants and products;
•
reactions can be monitored and graphs drawn and interpreted in relation to rate;
•
the rates of reactions are affected by changes in concentration, particle size and
temperature;
•
calculations can be carried out to determine the average rate of a chemical
reaction from a graph of the change in mass or volume against time.
Learning objectives
By the end of this topic, you should be able to:
•
state that reaction rates can be controlled by chemists;
•
explain that if reaction rates are too low a manufacturing process will not be
economically viable;
•
explain that if reaction rates are too high there is a risk of thermal explosion;
•
describe, using collision theory, the effects of concentration, pressure, surface
area (particle size), temperature and collision geometry on reaction rates;
•
calculate the relative rate of a reaction using the formula Rate = 1t .
2
TOPIC 1. REACTION RATES - COLLISION THEORY
1.1
Prior knowledge
Test your prior knowledge
Q1:
Go online
What three variables can be altered to change the rate of a reaction?
..........................................
Q2:
Name two variables you could monitor to follow the course of a reaction.
..........................................
Q3: A student obtains the following set of results shown in the graph when carrying
out a reaction with marble chips and dilute hydrochloric acid.
What is the average rate, in g s-1 , of reaction between 60 and 120 seconds?
..........................................
..........................................
© H ERIOT-WATT U NIVERSITY
TOPIC 1. REACTION RATES - COLLISION THEORY
1.2
Introduction
Chemical reactions occur at different rates depending on the nature of the reaction and
the conditions under which it occurs.
Fireworks are fast reactions
Rusting is a slow reaction
Rate is very important to chemists. In industry, as in the lab, it is important to
have knowledge of how to control rates of reaction. If reaction rates are too low, a
manufacturing process will not be economically viable. However, if reaction rates are
too high, there is a risk of thermal explosion.
How can rate be measured and what is meant by rate anyway?
Think about this:
Imagine yourself having to travel 500 kilometres. There are three methods of travel
available:
a) walking;
b) fast car;
c) rocket.
In each case, think about it and decide on a rough time it would take to complete the
journey, then work out the rate of travel.
The rate of travel in a), b) or c) will be measured in units of distance per time interval,
possibly in these units:
i. km day-1
ii. km hour-1
iii. km second-1
In each case, the rate is 'per time interval'. Rate is measured in units involving the
reciprocal of time. If the time taken is low, the rate is high. If the time taken is high, the
rate is low.
© H ERIOT-WATT U NIVERSITY
3
4
TOPIC 1. REACTION RATES - COLLISION THEORY
Rate is proportional
1
t
The study of rates of reaction and the factors which influence the rate is known as
'kinetics'.
Knowledge of how to control the rate of reaction is very important in the chemical
industry.
1.3
Rate of reaction
During the course of a chemical reaction, reactants are being converted into products.
Measurement of the rate of reaction involves measurement of the change in the amount
of a reactant or product in a certain time. The rate of reaction changes as it progresses,
being relatively fast at the start and slowing towards the end. What is being measured
is the average rate over the time interval chosen.
Reactions can be followed by measuring changes in concentration, mass and volume,
or by using properties of reactants or products which change in step with concentration
such as pressure, conductivity, pH or colour intensity. If the change in concentration is
measured then:
Average rate =
Change in concentration
Change in time
The units of rate would be moles per second (written as mol s -1 ). The units will, of
course, depend on which variable is being measured. In general:
Average rate =
Change in variable
Change in time
It is worth mentioning at this stage that the unit of concentration of a solution is
measured as moles per litre. In previous courses, this may have been shortened to
mol/ or even as mol/cm3 . At this level, the same unit will always be written as mol -1 .
This means that a change in concentration will be measured as mol -1 s-1 .
© H ERIOT-WATT U NIVERSITY
TOPIC 1. REACTION RATES - COLLISION THEORY
5
Measuring rate
Measuring the speed of a car, or even a rocket, can be performed directly using a
speedometer. Measuring the rate of a reaction directly is more difficult since there is
no such instrument as a reaction rate meter. A change in the amount of reactant or
product over time can be used.
When zinc reacts with hydrochloric acid, hydrogen gas is released. The volume of gas
produced can be recorded at regular time intervals. The illustration shows the
apparatus used and the results.
Q4: What volume of hydrogen is released in the first 5 seconds?
..........................................
Q5: Calculate the rate of hydrogen production in the first 5 seconds (remember units).
..........................................
Q6: Calculate the rate of hydrogen production from 5 to 10 seconds in the reaction.
..........................................
Q7: Explain the pattern of these two results.
..........................................
Q8: Calculate the average rate of the reaction from the start to the end. (Hint: When
has it ended?)
..........................................
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TOPIC 1. REACTION RATES - COLLISION THEORY
1.4
Collision theory
All substances are made up of particles called atoms, ions or molecules, and these
particles are constantly moving. The degree of movement depends upon the state of
the substance. This is known as the 'kinetic model' of matter. In any sample of solution,
liquid or gas there is a range of kinetic energies known as an energy distribution.
The collision theory of reactions suggests that, for a chemical reaction to occur,
particles must collide. Simple collision is not enough however, as many collisions do
not have sufficient kinetic energy to successfully rearrange the reactants to form new
products. In many cases, the way the particles line up, sometimes called the collision
geometry or orientation, is wrong and successful collision cannot occur.
Collision theory, based on the kinetic model of matter, provides an explanation for the
effect that various factors have on the rate of chemical reactions in terms of the number
of successful collisions which occur. Collision theory can be stated thus:
•
particles must collide to react;
•
not all collisions are successful;
•
sufficient energy is needed;
•
orientation must be correct.
In the illustration the single step formation of products in the reaction is a simplification.
The reaction itself actually proceeds in a number of stages.
Key point
The rates of reactions are affected by changes in concentration, particle size and
temperature, and the collision theory can be used to explain these effects.
Earlier work in chemistry showed that the rate of a reaction can be influenced by
concentration of reactants, particle size and temperature. How can collision theory
explain these effects?
A suitable reaction which will be used to study this over the next three activities is the
reaction of hydrochloric acid with a sample of calcium carbonate (chalk).
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TOPIC 1. REACTION RATES - COLLISION THEORY
7
Collisions and concentration
Look at the illustrations showing the result of collisions between two different
concentrations of hydrochloric acid and calcium carbonate, both after 10 seconds of
reaction. The hydrochloric acid is represented as a large sphere and the calcium
carbonate as a small sphere. Products of the reaction are shown as a combination
of a small sphere and a large sphere.
Q9: How many successful collisions have occurred using the 1mol -1 acid after 10
seconds?
..........................................
Q10: What is the rate of reaction with calcium carbonate and 1 mol -1 acid expressed
as successful collisions per second?
..........................................
Q11: What is the rate of reaction with calcium carbonate and 2 mol -1 acid expressed
in the same units?
..........................................
Q12: Predict what the rate might be if 4 mol -1 acid were used in the same experiment.
..........................................
Key point
In many reactions the rate of reaction is directly proportional to the concentration
of reactant, but there is no simple way to predict the relationship in advance.
..........................................
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TOPIC 1. REACTION RATES - COLLISION THEORY
Collisions and particle size
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Look at the illustrations showing the result of collisions between hydrochloric acid and
two particle sizes of the same mass of calcium carbonate, both after 10 seconds of
reaction.
Q13: Which sample of calcium carbonate has the greater surface area?
(Hint: Compare the number of particles on the outside edges of the solids which would
have been available before reaction.)
..........................................
Q14: How many successful collisions have occurred with the large sample after 10
seconds?
..........................................
Q15: What is the rate of reaction with the large sample expressed as successful
collisions per second?
..........................................
Q16: What is the rate of reaction with the small sample expressed in the same units?
..........................................
Key point
In many reactions involving solid reactant the rate of reaction is raised if the
surface area of the reactant is increased. This is the same as saying the particle
size is decreased.
..........................................
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TOPIC 1. REACTION RATES - COLLISION THEORY
9
Collisions and temperature
Look at the illustrations showing the result of collisions between hydrochloric acid and
calcium carbonate at temperature T ◦ C and then at T + 10 ◦ C, both after 10 seconds of
reaction.
Q17: How many successful collisions have occurred at T◦ C after 10 seconds?
..........................................
Q18: What is the rate of reaction at T ◦ C expressed as successful collisions per second?
..........................................
Q19: What is the rate of reaction at T +10 ◦ C expressed in the same units?
..........................................
Q20: What happens to the rate when the temperature is increased by 10 ◦ C?
a) It stays the same.
b) It halves.
c) It doubles.
d) It quadruples.
..........................................
Key point
In many reactions a rise in temperature of 10 ◦ C causes the rate of reaction to
approximately double.
There are more successful collisions in the earlier parts of the reactions, when there are
plenty of reactant particles available to react. As some react, there are fewer available
for collision and therefore the rate of reaction falls. This is as expected as a consequence
of collision theory.
..........................................
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TOPIC 1. REACTION RATES - COLLISION THEORY
1.5
Summary
Summary
1.6
•
1.7
•
Rates of chemical reactions can be controlled by chemists.
•
If reaction rates are too low, a manufacturing process will not be
economically viable.
•
If reaction rates are too high, there is a risk of thermal explosion.
•
The rates of reactions are affected by changes in concentration, particle size
and temperature, and collision theory can be used to explain these effects.
•
The relative rate of a reaction can be calculated using the formula Rate = 1t .
Resources
Higher Chemistry for CfE: J Anderson, E Allan and J Harris, Hodder Gibson,
ISBN 978-1444167528
End of topic test
End of Topic 1 test
Q21:
Go online
For any chemical, the temperature is a measure of the:
a) average kinetic energy of all the particles.
b) minimum kinetic energy required before reaction occurs.
c) average kinetic energy of the particles which react.
d) activation energy.
..........................................
Q22:
Which of the following is not a factor which affects the rate of a reaction?
a) Time taken for reaction to complete.
b) Concentration of reactants.
c) Collision geometry.
d) Kinetic energy of reactants.
..........................................
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TOPIC 1. REACTION RATES - COLLISION THEORY
Q23: When marble pieces and hydrochloric acid are reacted, carbon dioxide is evolved.
The curves below, showing the mass of the reaction vessel, were obtained under
different conditions.
The change in form of the rate of the curve from P to Q would be obtained by:
a) increasing the concentration of the acid.
b) decreasing the volume of the acid.
c) increasing the particle size of the marble.
d) decreasing the temperature of the reactants.
..........................................
Q24: Calculate the average rate of reaction in g s -1 for curve Q over the first 50 seconds
of reaction.
..........................................
Q25: Calculate the average rate of reaction in g s -1 for curve Q over the 250 seconds
shown on the graph.
..........................................
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TOPIC 1. REACTION RATES - COLLISION THEORY
Q26: Magnesium was added to 1.0 mol -1-1 sulfuric acid. The volume of hydrogen gas
released was plotted as curve C.
Which curve shows the results of a repeat experiment using the same volume of 0.5 mol
-1-1 sulfuric acid?
a) Curve A
b) Curve B
c) Curve D
d) Curve E
..........................................
Q27: Which curve shows the results of a repeat experiment using double the volume of
0.5 mol -1-1 sulfuric acid?
a) Curve A
b) Curve B
c) Curve D
d) Curve E
..........................................
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TOPIC 1. REACTION RATES - COLLISION THEORY
13
Q28: The same reaction was carried out at four different temperatures. The table shows
the times taken.
Temperature (◦ C)
20
30
40
50
Time (s)
60
30
14
5
These results show that:
a) a small rise in temperature results in a large increase in reaction rate.
b) the activation energy increases with increasing temperature.
c) the rate of reaction is directly proportional to the temperature.
d) the reaction is endothermic.
..........................................
Q29: A reaction was carried out under different conditions. The curves shown were
obtained when copper carbonate was reacted with acid, carbon dioxide being produced.
The change from P to Q could be brought about by:
a) decreasing the mass of copper carbonate.
b) increasing the particle size of the copper carbonate.
c) decreasing the particle size of the copper carbonate.
d) increasing the concentration of the acid.
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TOPIC 1. REACTION RATES - COLLISION THEORY
..........................................
Q30: Using the apparatus shown, a 6.075 g block of magnesium was added to 100 cm 3
of 1.0 mol -1 hydrochloric acid and the change in mass noted at regular time intervals.
Which of the following graphs would be drawn from the results?
a)
b)
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TOPIC 1. REACTION RATES - COLLISION THEORY
c)
d)
..........................................
Q31: Calculate the number of moles of magnesium added to the acid.
..........................................
Q32: Calculate the number of moles of hydrochloric acid present at the start of the
reaction.
..........................................
Q33: A vitamin C solution, which is acidic, slowly breaks down any sugar present.
The sugar concentration measured over a number of hours was found to be
decomposing at an average rate of 0.0002 mol -1-1 min-1 .
If the starting concentration in a fresh solution was found to be 0.5 mol -1-1 , what would
be its sugar concentration after 5 hours?
..........................................
..........................................
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TOPIC 1. REACTION RATES - COLLISION THEORY
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Topic 2
Reaction rates - reaction profiles
Contents
2.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
19
2.2 Interpreting graphs . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
2.3 Activation energy . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
19
25
2.4 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
2.5 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
30
30
2.6 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
31
Prerequisite knowledge
Before you begin this topic, you should already know that:
•
reactions can be monitored and graphs plotted and interpreted (National 4, Unit
1);
•
when chemical reactions occur, they are often accompanied by a change in energy
(National 4, Unit 1);
•
reactions can be exothermic or endothermic - this is dependent on the overall
energy change taking place (National 4, Unit 1);
•
there are four factors that affect the rate of reaction: temperature, concentration,
catalyst and surface area (National 3, Unit 1);
•
the effects of concentration, pressure, surface area (particle size), temperature
and collision geometry on reaction rates can be described using collision theory
(Higher Unit 1, Topic 1).
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TOPIC 2. REACTION RATES - REACTION PROFILES
Learning objectives
By the end of this topic, you should be able to:
•
interpret rate graphs;
•
explain that a potential energy diagram can be used to show the energy pathway
for a reaction;
•
state that enthalpy change is the energy difference between products and
reactants;
•
explain that enthalpy change can be calculated from a potential energy diagram;
•
state that enthalpy change has a negative value for exothermic reactions and a
positive value for endothermic reactions;
•
explain that an activated complex is an unstable arrangement of atoms formed
during a reaction at the maximum of the potential energy barrier;
•
state that activation energy is the energy required by colliding particles to form an
activated complex;
•
explain that activation energy can be calculated from a potential energy diagram;
•
state that temperature is a measure of the average kinetic energy of the particles
of a substance (revision: Higher Unit 1, Topic 1);
•
state that activation energy is the minimum kinetic energy required by colliding
particles before reaction may occur;
•
describe that energy distribution diagrams can be used to explain the effect of
changing temperature on the kinetic energy of particles;
•
describe that the effect of temperature on reaction rate can be explained in terms
of an increase in the number of particles with energy greater than the activation
energy.
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2.1
19
Prior knowledge
Test your prior knowledge
Q1: During the course of a combustion reaction, energy in the form of heat is given
out. What type of reaction is this?
..........................................
Q2: During the course of a decomposition reaction, energy in the form of heat is taken
in. What type of reaction is this?
..........................................
Q3: Collision theory provides an explanation for the effect that various factors have on
the rate of chemical reactions. For successful collisions to occur, what must happen?
(Name two things.)
..........................................
2.2
Interpreting graphs
Investigations which measure the change in a property, such as concentration, during
the progress of a chemical reaction often have results displayed as a graph or a table.
In the reaction between zinc and hydrochloric acid, the total volume of hydrogen given
off was measured against time and the graph that resulted showed a steep slope which
tailed off to a level gradient as the reaction finished and the volume of hydrogen released
stayed constant. Revisit the activity in the previous topic if you are unsure about this
point.
It would also be possible to follow the progress of the same chemical reaction between
zinc and hydrochloric acid by measuring the mass of the beaker and its contents
against time using the apparatus shown in Figure 2.1. The graph which results is
shown in Figure 2.2.
Figure 2.1: Zinc and hydrochloric acid mass experiment
..........................................
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Graphs which use the same axes to plot results for change in concentration of both
reactant and product on the same graph may also be met. Figure 2.2 shows such a
graph.
Figure 2.2: Reactant and product on shared axes
..........................................
Graphs which use the same axes and place the results for different experiments in
which the concentration or temperature are varied from one experiment to the next are
common and show up how that variable affects the reaction progress.
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21
Measuring rate at different concentrations
A chemical reaction between zinc and hydrochloric acid is repeated at three different
concentrations of acid. The data collected, involving changes of volume, is graphed
against time, giving multiple graphs on the same axes.
The three features of each line which give information about the reaction, and are
commonly met in examination questions are:
•
the steepness of the slope;
•
the point at which the line becomes horizontal;
•
the final level of the horizontal line.
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Q4:
Which graph shows the steepest slope?
a) 0.5 mol -1
b) 1.0 mol -1
c) 2.0 mol -1
..........................................
Q5:
Which graph shows the fastest rate of reaction?
a) 0.5 mol -1
b) 1.0 mol -1
c) 2.0 mol -1
..........................................
Q6:
How long did it take for the 2.0 mol -1 reaction to finish?
..........................................
Q7:
What is the volume of hydrogen in the 2.0 mol -1 reaction after 30 seconds?
..........................................
Q8:
Compare the 1.0 mol -1 graph with the 2.0 mol -1 graph in three areas:
•
the steepness of the slope;
•
the point at which the line becomes horizontal;
•
the final level of the horizontal line.
Explain the differences.
..........................................
Q9: Predict the final level (in cm3 to 1 decimal place) of the horizontal line in the 0.5
mol -1 graph.
..........................................
In each of the graphs showing how a variable affects the reaction progress, it is the
slope of the graph which gives an indication of the average rate of reaction. The rate of
reaction changes as it progresses, being relatively fast at the start and slowing towards
the end.
Average rate is measured as:
Average rate =
Change in variable
Change in time
In the graphs developed previously, a comparison of rate of reaction of 0.5 mol -1 , 1.0
mol -1 and 2.0 mol -1 acid can be determined by looking at how much hydrogen has
been given off in the same time span.
If the volume of hydrogen given off after 5 seconds is measured, the question "How
does the rate of reaction depend on the concentration of acid?" can be answered.
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23
The table of results (Table 2.1) taken from the graphs is shown below. The resulting
graph of rate of reaction is plotted against concentration of hydrochloric acid (Figure
2.3). It is clearly a straight line showing that, in this case, the rate is directly
proportional to the concentration of hydrochloric acid. (Remember, however, that there
is no simple way to predict the relationship in advance and each concentration/rate
relationship must be investigated experimentally to determine any relationship.)
Table 2.1: Rate versus concentration of hydrochloric acid
Hydrochloric acid
concentration/mol -1
Volume of
hydrogen/cm3
Rate/(cm3 of H2 s-1 )
0.5
6.5
1.3
1.0
13.0
2.6
2.0
26.0
5.2
..........................................
Figure 2.3
..........................................
Q10: Predict the volume of hydrogen in cm 3 given off in 5 seconds in a similar
experiment using 1.5 mol -1 hydrochloric acid.
..........................................
Q11: Predict the average rate in cm 3 s-1 in a similar experiment using 4.0 mol -1
hydrochloric acid.
..........................................
Many investigations into how concentration of solutions or temperature changes affect
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TOPIC 2. REACTION RATES - REACTION PROFILES
the rate of a reaction involve 'clock reactions'. The time taken for a colour change to
appear may be an experiment you have seen. Sometimes a cross drawn on paper is
placed below the reaction vessel. The time taken for it be obscured is called the end
point. In each case, the time taken to reach the same stage in the reaction is being
measured.
The time is converted into rate by calculating the reciprocal of time:
Rate =
1
t
Conversions between time, rate and concentration or temperature are common.
Example
Problem
The graph shows a rate of reaction plotted against concentration of a reactant.
Figure 2.4: Rate versus concentration graph
..........................................
a) How long would it take to complete the reaction if the reactant concentration is 0.4
mol -1 ?
b) At what concentration was the reaction completed in 10 seconds?
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Solution
a) From the graph, at 0.4 mol -1 the rate = 0.2 s-1 .
1
t
1
0.2 =
t
1
t=
0.2
t = 5 seconds
Rate =
b) Since the rate is given by:
1
t
1
Rate =
10
Rate = 0.1
Rate =
From the graph, concentration = 0.2 mol -1
..........................................
2.3
Activation energy
Collision theory states that colliding particles must have 'sufficient energy' to react.
Some reactions start as soon as the reactants are mixed. A neutralisation reaction
between hydrochloric acid and sodium hydroxide solutions would react immediately on
mixing. The colliding particles, in this case, must have sufficient energy to react.
Some reactions require an input of energy to start. For example, a firework burns well
in air only after lighting. The particles in the firework had insufficient energy to react at
room temperature.
The activation energy is the minimum kinetic energy required by colliding particles
before reaction will occur. This is often thought of as an 'energy barrier' which has to be
overcome. The kinetic model of matter shows the energy distribution of particles in a
gas, liquid or solution. The average kinetic energy of the particles is measured as
'temperature'. Only those molecules with a kinetic energy greater than the activation
energy (shown as EA or as Ea ) are capable of successful collision.
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TOPIC 2. REACTION RATES - REACTION PROFILES
Figure 2.5: Kinetic energy distribution graph
..........................................
These questions refer to the four points on the distribution graph.
Q12: Which point represents particles with the lowest kinetic energy?
a)
b)
c)
d)
A
B
C
D
..........................................
Q13: Which point represents the least number of particles?
a)
b)
c)
d)
A
B
C
D
..........................................
Q14: The average kinetic energy of particles determines the 'temperature'. Which point
represents particles nearest to the average kinetic energy?
a)
b)
c)
d)
A
B
C
D
..........................................
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27
Q15: Which point represents particles with sufficient kinetic energy to react
successfully?
a)
b)
c)
d)
A
B
C
D
..........................................
The effect of temperature on kinetic energy
Temperature is a measure of the average kinetic energy of the particles of a substance.
An increase in temperature changes the energy distribution and increases the average
kinetic energy.
Controlling the rate - temperature and kinetic energy
The following graphs show the kinetic energy distribution of a group of particles involved
in a chemical reaction at T◦ C and T + 10 ◦ C.
Compare the graphs and consider the effect of increasing the temperature by
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10 ◦ C.
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TOPIC 2. REACTION RATES - REACTION PROFILES
Q16: How has the increased temperature affected the value of the activation energy?
a) Increased
b) Decreased
c) Not changed
..........................................
Q17: How has the increased temperature affected the average kinetic energy of the
particles?
a) Increased
b) Decreased
c) Not changed
..........................................
Q18: At the increased temperature, which point or points represent particles with
sufficient kinetic energy to react successfully?
a)
b)
c)
d)
D
C and D
B,C and D
A,B,C and D
..........................................
Q19: What does the shaded area to the right of E A represent?
a)
b)
c)
d)
Particles with sufficient energy to react.
Particles with insufficient energy to react.
Particles with the average kinetic energy.
Particles with the same kinetic energy.
..........................................
Q20: How has the 10 ◦ C increase in temperature affected the number of particles with
energy greater than the activation energy?
a) Roughly the same.
b) Roughly halved.
c) Roughly doubled.
..........................................
Remember, the Chemical Industry is not in existence to manufacture chemicals: like any
other industry it exists to create wealth and wealth can only be created if it can make
profits.
For this reason, an understanding of rate is very important to chemists. In industry, as
in the lab, it is important to have knowledge of how to control rates of reaction.
If reaction rates are too low, a manufacturing process will not be economically viable.
However, if reaction rates are too high, there is a risk of thermal explosion.
In recent years, a number of thermal explosions at chemical plants have made the news,
including West Thurrock, Essex and The Imperial Sugar Company, USA.
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Reactions increase their rate at higher temperatures because a higher proportion of
the molecules involved have energy in excess of the activation energy and more
successful collisions can occur. It is observed that a 10◦ C rise is responsible for an
approximate doubling of rate in many reactions.
..........................................
The effect of light energy on kinetic energy
With some chemical reactions, light energy is absorbed and can provide particles with
energy sufficient to raise their kinetic energy above the value needed to cross the
activation energy barrier. Photosynthesis and photography provide two common
examples of photochemical reactions.
(This subject will be considered further in a later topic on skin care products.)
There are many examples of this type of reaction taking place in the atmosphere. High
energy ultraviolet radiation which would be harmful to humans is absorbed by ozone
(O3 ) molecules in the upper atmosphere.
The ozone layer which protects us is fragile, however, as the ozone can be broken
down by certain chemicals. Moves have taken place to reduce the use of these
chemicals and protect the ozone layer.
Key point
•
Temperature is a measure of the average kinetic energy of the particles of
a substance and activation energy is the minimum kinetic energy required
by colliding particles before reaction can occur.
•
Energy distribution diagrams can be used to explain how an increase in
temperature or, in some chemical reactions the energy from light, increases
the number of particles with energy greater than the activation energy.
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TOPIC 2. REACTION RATES - REACTION PROFILES
2.4
Summary
Summary
2.5
•
•
Graphs which use the same axes and place the results for different
experiments in which the concentration or temperature are varied from one
experiment to the next are common and show us how that variable affects
the reaction progress.
•
Temperature is a measure of the average kinetic energy of the particles of
a substance.
•
Activation energy is the minimum kinetic energy required by colliding
particles before reaction can occur.
•
Energy distribution diagrams can be used to explain how an increase in
temperature or, in some chemical reactions, the energy from light increases
the number of particles with energy greater than the activation energy
(Ea /EA ).
•
Reactions increase their rate at higher temperatures because a higher
proportion of the molecules involved have energy in excess of the activation
energy so more successful collisions can occur.
•
It is observed that a 10◦ C rise is responsible for an approximate doubling of
rate in many reactions.
•
The effect of temperature on reaction rate can be explained in terms of an
increase in the number of particles with energy greater than the activation
energy.
Resources
Higher Chemistry for CfE: J Anderson, E Allan and J Harris, Hodder Gibson,
ISBN 978-1444167528
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TOPIC 2. REACTION RATES - REACTION PROFILES
2.6
31
End of topic test
End of Topic 2 test
Q21: A piece of phosphorus ignites when touched with a hot wire, whereas magnesium
ribbon needs strong heating before it will burn. What can you deduce about the
activation energies of the two reactions?
a)
b)
c)
d)
The activation energy is less for magnesium than for phosphorus.
The activation energy is less for phosphorus than for magnesium.
The activation energies are the same for the reactions.
No information about activation energy can be deduced.
..........................................
Q22: A student carries out a reaction and then repeats the experiment at a temperature
ten degrees higher. What is the likely effect on rate of reaction?
a)
b)
c)
d)
No effect.
The reaction rate halves.
The reaction rate doubles.
The reaction rate triples.
..........................................
Q23: Why is this?
a)
b)
c)
d)
The activation energy has been decreased.
The activation energy has increased.
Fewer particles have energy greater than the activation energy.
More particles have energy greater than the activation energy.
..........................................
Q24: Activation energy is:
a)
b)
c)
d)
the minimum kinetic energy required by colliding molecules for a reaction to occur.
the maximum kinetic energy colliding molecules can reach for a reaction to occur.
the temperature at which a reaction will take place.
the temperature at which rate of reaction will be greatest.
..........................................
Q25: The activated complex is:
a)
b)
c)
d)
the initial stage of reaction when the reactants are first put together.
an intermediate stage at the top of the activation energy barrier.
the maximum kinetic energy colliding molecules can reach for a reaction to occur.
the final stage of a reaction when the products have been formed.
..........................................
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TOPIC 2. REACTION RATES - REACTION PROFILES
Q26: The activation energy can be calculated using:
a)
b)
c)
d)
time taken for reaction to complete.
temperature the reaction occurs at.
potential energy diagrams.
the activated complex.
..........................................
Q27: The graph below shows how the concentrations of both a reactant and a product
change during the course of a reaction.
What is the average rate of reaction over the first 15 seconds?
..........................................
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TOPIC 2. REACTION RATES - REACTION PROFILES
Q28: This graph shows the effect of temperature on the rate of a reaction.
How long, in seconds, did the reaction take to complete at 40 ◦ C?
..........................................
Q29: Explain, in terms of the collision theory, why increasing the temperature increases
the reaction rate.
..........................................
Q30: A reaction between A and B forms product C. At a temperature of 40 ◦ C the
reaction is complete after 20 seconds. At what temperature would the rate be expected
to be four times as fast?
a) 50◦ C
b) 60◦ C
c) 80◦ C
d) 160◦ C
..........................................
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TOPIC 2. REACTION RATES - REACTION PROFILES
Q31: A reaction was carried out at four different temperatures. The table shows the
time taken for the reaction to occur.
Temp (◦ C)
20
30
40
50
Time (seconds)
60
30
14
5
The results show that:
a)
b)
c)
d)
the activation energy increases with increasing temperature.
a small rise in temperature results in a large increase in reaction rate.
the rate of the reaction is directly proportional to the temperature.
the reaction is endothermic.
..........................................
..........................................
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Topic 3
Catalysis
Contents
3.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
37
3.2 Introduction to catalysis . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
3.3 Catalyst mechanism . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
37
38
3.4 Catalysts in industry . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
3.4.1 Examples of catalysts . . . . . . . . . . . . . . . . . . . . . . . . . . . .
39
40
3.5 Potential energy diagrams . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
3.5.1 What produces this energy change? . . . . . . . . . . . . . . . . . . . .
43
47
3.5.2 The thermochemical equation . . . . . . . . . . . . . . . . . . . . . . .
3.6 Catalysts and energy . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
49
52
3.6.1 The activated complex . . . . . . . . . . . . . . . . . . . . . . . . . . . .
3.6.2 How catalysts work . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
53
56
3.7 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
3.8 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
60
61
3.9 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
62
36
TOPIC 3. CATALYSIS
Prerequisite knowledge
Before you begin this topic, you should already know that:
•
there are four factors that affect the rate of reaction: temperature, concentration,
catalyst and surface area (National 3, Unit 1);
•
adding a catalyst will speed up a reaction (National 3, Unit 1);
•
reactions can be exothermic or endothermic - this is dependent on the overall
energy change taking place (National 4, Unit 1);
•
the effects of concentration, pressure, surface area (particle size), temperature
and collision geometry on reaction rates can be understood using collision theory
(Higher, Unit 1, Topic 1);
•
a potential energy diagram can be used to show the energy pathway for a reaction
(Higher, Unit 1, Topic 1);
•
enthalpy change is the energy difference between products and reactants - it can
be calculated from a potential energy diagram (Higher, Unit 1, Topic 1);
•
enthalpy change has a negative value for exothermic reactions and a positive value
for endothermic reactions (Higher, Unit 1, Topic 1);
•
activation energy can be calculated from potential energy diagrams (Higher, Unit
1, Topic 1);
Learning objectives
By the end of this topic, you should be able to:
•
state that the enthalpy change for a reaction has a negative value for exothermic
reactions and a positive value for endothermic reactions;
•
explain that the activated complex is an unstable arrangement of atoms formed at
the maximum of the potential energy barrier, during a reaction;
•
state that the activation energy is the energy required by colliding particles to form
an activated complex;
•
explain that activation energy can be calculated from potential energy diagrams;
•
state that a catalyst provides an alternative reaction pathway with a lower activation
energy;
•
explain that a potential energy diagram can be used to show the effect of a catalyst
on activation energy;
•
explain how catalysis works and learn about adsorption;
•
define the terms 'activated complex' and 'activation energy' and hence explain how
catalysts work in terms of potential energy diagrams.
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TOPIC 3. CATALYSIS
3.1
37
Prior knowledge
Test your prior knowledge
Q1: For any chemical, the temperature is a measure of the:
a)
b)
c)
d)
average kinetic energy of all the particles;
minimum kinetic energy required before reaction occurs;
average kinetic energy of the particles which react;
activation energy.
..........................................
Q2: During the course of a decomposition reaction, energy in the form of heat is taken
in. What type of reaction is this?
a) Exothermic
b) Endothermic
..........................................
Q3: What is activation energy?
a)
b)
c)
d)
The minimum kinetic energy required by colliding molecules for a reaction to occur.
The maximum kinetic energy colliding molecules can reach for a reaction to occur.
The temperature at which a reaction will take place.
The temperature at which rate of reaction will be greatest.
..........................................
3.2
Introduction to catalysis
Catalysts are amongst the most important chemicals in the world around us.
Catalysts are essential in:
•
the production of materials such as plastics;
•
the production of fuels to power our transport;
•
the removal of pollutants produced by transport;
•
the production of medicines and food;
•
the functioning and health of our bodies.
It is said that 90% of all manufactured items use catalysts at some stage of their
production.
Catalysts are substances which alter the rates of reactions, but are chemically
unchanged at the end of the reactions. They can be recovered at the end of reactions
chemically unchanged. They are neither reactants nor products and do not appear in
the chemical equation. Often, the catalyst used in a reaction is shown above the arrow
in the equation, as in the Haber Process.
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38
TOPIC 3. CATALYSIS
Figure 3.1: Haber process
..........................................
Catalysts allow chemical reactions to occur more quickly at lower temperatures and so
reduce energy costs.
3.3
Catalyst mechanism
Hydrogenation of ethene
Often, the catalyst is in a different state from the reactants. An example of this is the
addition of hydrogen gas to ethene gas using a solid nickel catalyst.
Addition of hydrogen to ethene
Go online
The atoms within the metal are fully bonded to their neighbours whereas the atoms on
the surface have 'spare' bonds and are capable of forming weak bonds with the reactant
molecules. This is known as adsorption. (This should not be confused with absorption,
which involves a substance penetrating inside a solid - like water being soaked up by a
sponge.) When the molecules leave the surface of the catalyst, the process is known a
desorption.
Q4:
a)
b)
c)
d)
Reactant molecules bond to the surface of the metal. This is known as:
absorption.
desorption.
adsorption.
activation.
..........................................
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Q5: What happens to the bonds within the molecules when they attach to the surface?
a)
b)
c)
d)
They break.
They get weaker.
They get stronger.
They stay the same.
..........................................
Q6: When the product molecules leave the surface of the catalyst, what change has
taken place to the surface of the catalyst?
..........................................
Q7: Which of these statements is false?
a) The catalyst speeds up the reaction by weakening the bonds in the reactant
molecules.
b) The larger the surface area of the catalyst the more effective it will be.
c) The catalyst speeds up the reaction without taking part.
d) The catalyst speeds up the reaction without being used up.
..........................................
Enzymes
Enzymes are biological catalysts that are responsible for the chemical processes that
occur in all living organisms, from the simplest bacteria to the most complex mammals.
They are proteins and are the target for many modern pharmaceutical products used to
control diseases.
Enzymes are studied in more detail in the topic on Proteins later in the course.
Key point
Many catalysts work by adsorbing reactant molecules onto their surface.
3.4
Catalysts in industry
Many important industrial processes involve catalysis.
Figure 3.2: Examples of industrial catalysis
Process
Use
Catalyst
Haber
Production of ammonia
Iron
Contact
Production of sulfuric acid
Vanadium (V) oxide
Oswald
Production of nitric acid
Platinum
Hardening of oils
Production of margarine
Nickel
Catalytic cracking
Production of alkenes
Aluminium oxide
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TOPIC 3. CATALYSIS
..........................................
3.4.1
Examples of catalysts
Catalysis by cobalt(II) chloride solution
The oxidation of aqueous tartrate ions by hydrogen peroxide solution is slow, even when
the solution is heated to 60 ◦ C. The reaction can be catalysed by adding a solution
containing cobalt(II) ions.
Figure 3.3: Oxidation of tartrate ions
5H2 O2 (aq) + C4 H4 O6 2- (aq) → 4CO2 (g) + 6H2 O(l) + 2OH- (aq)
..........................................
Catalysis by cobalt(II) chloride solution
Go online
Study the following diagram, looking closely for any evidence that the cobalt(II) chloride
is acting as a catalyst, and then answer the questions which follow.
Catalysis by cobalt(II) chloride
Q8:
a)
b)
c)
d)
Which species is responsible for the pink colour at the beginning?
Cl- (aq) ions
Co2+ (aq) ions
Co3+ (aq) ions
Tartrate ions
..........................................
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Q9: Suggest a species that could be responsible for the green colour which appears?
a)
b)
c)
d)
Cl- (aq) ions
Co2+ (aq) ions
Co3+ (aq) ions
Tartrate ions
..........................................
Q10: Which species acts as the catalyst?
a)
b)
c)
d)
Cl- (aq) ions
Co2+ (aq) ions
Co3+ (aq) ions
Tartrate ions
..........................................
Q11: A student defines a catalyst as a substance which speeds up a reaction without
taking part. Explain whether or not this is correct, with reference to the above reaction.
..........................................
Key point
Cobalt(II) chloride speeds up the reaction between tartrate ions and hydrogen
peroxide. Colour changes clearly show that the catalyst takes part in the reaction
and is chemically unchanged at the end of the reaction.
Ozone layer destruction
Another example of catalysis is found in the destruction of the ozone layer by CFCs
(chlorofluorocarbons). Ozone, O3 , is constantly being formed in the upper atmosphere
where it provides a protective layer by absorbing a lot of the harmful ultraviolet radiation
which could cause serious skin damage, including skin cancer. Ozone slowly breaks
down by reacting with oxygen atoms to form oxygen molecules.
The breakdown of ozone
This breakdown is accelerated by chlorine atoms, which are produced when sunlight in
the upper atmosphere causes the CFCs to break down.
The catalysed breakdown of ozone
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TOPIC 3. CATALYSIS
Q12: Which species is acting as a catalyst?
a)
b)
c)
d)
O3
O2
ClO
Cl
..........................................
The Monsanto process
A very important industrial example of catalysis is the Monsanto process for the
synthesis of ethanoic acid from methanol.
The Monsanto process
The reaction takes place in solution using a soluble compound containing rhodium.
Industrial uses of catalysts
Go online
Throughout this topic, examples have been given of the use of catalysts in industry. You
do not need to know them all, but should be able to give some examples. Without the
use of catalysts, most industrial processes would not be feasible. You may need to refer
back to the earlier sections of this topic or use the Internet or some of the text books
listed in the section on Resources.
Q13: Use the terms from the 'word bank' to complete the table. Some of the terms may
be required more than once.
Process
Catalyst
Type
Purpose
Haber
Production of
ammonia
Washing powders
Breakdown of protein
stains
Catalytic converters
Pt, Pd, Rb
Brewing
Production of ethanol
Catalytic cracking
Production of alkenes
Food processing
Rennin
Ostwald
Pt gauze
Monsanto
Rhodium complex
Word bank: Al2 O3 , Enzyme, Ethanoic acid from methanol and CO, Iron, Production
of cheese, Production of nitric acid, Protease, Reduction of pollution, Solid, Solution,
Zymase
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..........................................
3.5
Potential energy diagrams
Consider the exothermic reaction between zinc and sulfuric acid.
Zn(s) + H2 SO4 (aq) → ZnSO4 (aq) + H2 (g)
After the reaction, the energy stored in the products is less than the energy originally
stored in the reactants - that is, some of the chemical potential energy stored in
reactants has been transferred to the surroundings. This transferred heat energy is
defined as the enthalpy change for this reaction.
Strictly, the change in energy which occurs during a chemical reaction consists of two
components: heat and work. We live on the surface of the Earth, at the bottom of a 'sea'
of atmosphere which exerts pressure. In the reaction above, the gaseous hydrogen
produced will have to 'work' to displace the atmosphere. This will require use of some
energy, so the heat energy released to the surroundings will be less than the total energy
change. In this reaction, for 1 mole of zinc, the heat (enthalpy) change is 152.4 kJ and
the work against the atmosphere 2.5 kJ, the overall energy change being 154.9 kJ.
Chemists often carry out reactions at constant pressure, hence the use of 'enthalpy
change' which is the heat change at constant pressure.
Let's consider another familiar exothermic reaction, the burning (oxidation) of methane
in a flame. You see this every time you use natural gas.
CH4 + 2O2 → CO2 + 2H2 O
These changes are often shown as a potential energy diagram. An example is shown
for methane oxidation.
The first diagram shows the potential energy (the enthalpy) of the reactants at the start
of the reaction
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TOPIC 3. CATALYSIS
Next, the potential energy (enthalpy) of the products at the end of the reaction is shown.
The arrow shows the energy (enthalpy) change during the reaction.
In the case of methane oxidation, energy is given out to the surroundings.
Figure 3.4: Methane oxidation
..........................................
This diagram shows the potential energy of the reactants, at the start of the reaction. It
also shows the potential energy of the products, at the end of the reaction. The arrow
shows the energy given to the surroundings.
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The easiest measurement to make is the enthalpy change, ΔH, when the reaction
occurs. This is defined as the difference in enthalpy of the products and reactants by
the equation below. The units of energy and enthalpy are joules (J).
Enthalpy change
ΔH is negative for an exothermic reaction. The energy of the products is less than that
of the reactants.
Exothermic reaction
The potential energy diagram for an endothermic reaction, the decomposition of
calcium carbonate is shown below.
Calcium carbonate decomposition
In this case, the energy of the products is greater than that of the reactants.
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TOPIC 3. CATALYSIS
Q14: What will happen to the temperature of the surroundings during an endothermic
reaction?
a) It will increase.
b) It will not change.
c) It will decrease.
..........................................
Q15: What will the value of ΔH be for an endothermic reaction?
a) Negative
b) Zero
c) Positive
..........................................
Endothermic reaction
In order to distinguish endothermic and exothermic reactions, the + and - sign are
always used when giving ΔH values.
Q16: The equation for the 'water gas' reaction is:
C(s) + H2 O(g) → CO(g) + H2 (g)
Note that this equation states that the carbon is in the solid state (s) and the other
materials are gases (g).
This reaction absorbs energy from its surroundings. What type of reaction is this?
a)
b)
c)
d)
Fast
Exothermic
Explosive
Endothermic
..........................................
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Q17: What will be the sign of ΔH?
a)
b)
c)
d)
Positive
Has no sign.
Negative
Insufficient information to know.
..........................................
Q18: Draw a potential energy diagram for this reaction.
..........................................
But how much energy is involved in this reaction? In the case of the water gas reaction,
121 kJ is absorbed for each mole of carbon and steam used. In other words, ΔH =
+121 kJ mol-1 .
The full potential energy diagram for the water gas reaction is as follows.
ΔH = +121 kJ mol-1
Products
CO(g) + H2(g)
H/kJ
-111
-232
Reactants
C(s) + H2O(g)
Reaction progress
Water gas potential energy diagram
Key point
A potential energy diagram can be used to show the energy changes for a
reaction.
The enthalpy change is the energy difference between products and reactants.
The enthalpy change can be calculated from a potential energy diagram.
The enthalpy change has a negative value for exothermic reactions and a positive
value for endothermic reactions.
3.5.1
What produces this energy change?
The chemical energy in molecules is stored in the bonds. When a reaction occurs, the
bonds in the reactant molecules are broken and the atoms rearrange to form new bonds
with different energies in the products.
For example, think of making ammonia from nitrogen and hydrogen.
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TOPIC 3. CATALYSIS
The balanced chemical equation is
N2 + 3H2
2NH3
You can think of this reaction as first breaking the bonds in the nitrogen and hydrogen
molecules. This stage requires an input of energy.
Then these atoms rearrange and combine to form new N - H bonds in ammonia. This
stage releases energy.
The enthalpy change will be given by the total energy in new bonds, minus the energy
present in the original bonds.
It is important to remember that reactions do not normally proceed by such a route, but
the enthalpy change can be calculated as if they did.
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TOPIC 3. CATALYSIS
3.5.2
49
The thermochemical equation
It is usual to express this information in a thermochemical equation such as that for
the water gas reaction (Figure 3.5).
Figure 3.5: Thermochemical equation for the water gas reaction
..........................................
Three points should be noted.
1. The enthalpy value quoted in the balanced equation is measured in kilojoules per
mole (kJ mol-1 ).
Look at Figure 3.5 again. The 121 kJ of energy is absorbed when one mole of
solid carbon and one mole of steam react; two moles of each reacting would
absorb 242 kJ, and so on. (This makes sense when you think that burning 2 kg of
carbon would give out twice the energy from burning 1 kg.)
2. The equation always contains the state symbols of reactants and products.
The enthalpy of reactions will change if the states are different.
Look carefully at these two equations for the reaction of hydrogen and oxygen.
(N.B. the equation has been written for forming 1 mole of water. The " 1 /2 " before
the oxygen means 1 /2 mole of reactant, not 1 /2 a molecule of oxygen.)
The extra enthalpy in the second case is due to the heat released when the water
vapour condenses to liquid water.
3. Since the enthalpy change is independent of the route of the reaction, the
enthalpy change for the reverse reaction will be equal in value, but of opposite
sign.
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TOPIC 3. CATALYSIS
Q19:
For the reaction below:
What is the enthalpy change for the reverse reaction?
CO2 (g) → CO(g) + 1 /2 O2 (g)
a)
b)
c)
d)
ΔH = -566 kJ mol-1
ΔH = -283 kJ mol-1
ΔH = 0 kJ mol-1
ΔH =+283 kJ mol-1
..........................................
Q20: The equation for calculating the enthalpy change (ΔH) for a reaction is:
a)
b)
c)
d)
ΔH = Hproducts + Hreactants
ΔH = Hproducts - Hreactants
ΔH = Hreactants + Hproducts
ΔH = Hreactants - Hproducts
..........................................
Q21: A ΔH value of -500 kJ mol -1 would indicate that the reaction is:
a)
b)
c)
d)
Fast
Reversible
Exothermic
Endothermic
..........................................
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Q22:
The potential energy diagram for the combustion of one mole of ethene is:
What is the enthalpy change for this reaction?
a)
b)
c)
d)
ΔH = +52 kJ mol-1
ΔH = -1362 kJ mol-1
ΔH = -1310 kJ mol-1
ΔH = -1414 kJ mol-1
..........................................
Q23:
The equation for the combustion of ethanol is:
How many kJ of energy will be released to the surroundings if 5 moles of ethanol are
burned?
..........................................
Q24:
Given that:
what would be the enthalpy change for the reaction below? (Hint: don't forget the sign.)
..........................................
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TOPIC 3. CATALYSIS
Q25: The Ostwald process for producing nitric acid involves the oxidation of ammonia
(NH3 ) to form nitric oxide (NO) and water. Write a balanced equation for this reaction.
..........................................
Q26: Given that the enthalpy of the reactants (as written in the balanced equation) is
-184 kJ and the enthalpy of the products is -1090 kJ, calculate the enthalpy change for
this reaction in kJ. Remember the sign.
..........................................
Q27: What would be the enthalpy change per mole of ammonia oxidised? Give your
answer in kJ mol-1 to 1 decimal place.
..........................................
Q28: When 102 g of ammonia is oxidised, how many kJ of heat have to be removed to
maintain a constant temperature?
..........................................
3.6
Catalysts and energy
The effect of a catalyst on a reaction rate can be explained in terms of energy.
Potential energy diagrams are used to show the energy pathway of a chemical reaction.
You will study the following diagrams in much more detail in the next topic in this unit.
Q29: Which diagram shows an exothermic reaction, one in which the products have
less energy than the reactants so heat is released to the surroundings?
a) A
b) B
..........................................
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Q30: Enthalpy change is the difference in energy between products and reactant. What
is the approximate enthalpy change in diagram A?
..........................................
Q31: Enthalpy change is the difference in energy between products and reactant.
Which of these could be the enthalpy change in B?
a)
b)
c)
d)
+1040 kJ
-1040 kJ
+890 kJ
-890 kJ
..........................................
3.6.1
The activated complex
When considering energy changes in reactions, only the starting point and finishing
point are considered. Energy has to be put in to break the old bonds before energy can
be released when the new bonds form. As a result, the potential energy diagram for an
exothermic reaction has the shape shown below.
Figure 3.6: Potential energy diagram for an exothermic reaction
..........................................
There is an energy barrier between the reactants and products. Reactant molecules
must have enough energy when they collide to get over this barrier.
The maximum point, X, in the graph above corresponds to the activated complex. This
is a very unstable arrangement of atoms in which the old bonds are half broken and the
new bonds are half formed. This species can either break down to reform the reactants
('fall' back down the same side of the 'hill') or break down to form the products ('fall' down
the other side of the 'hill').
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TOPIC 3. CATALYSIS
Reaction profile
Go online
Reaction profile with activated complex
..........................................
The activation energy (EA ), is the energy required by the colliding molecules to form
the activated complex, i.e. it is the height of the barrier. The activation energy can be
calculated from potential energy diagrams.
Q32: Which of these could be the activation energy of the reaction shown in Figure
3.6?
a)
b)
c)
d)
-180 kJ
+180 kJ
-110 kJ
+110 kJ
..........................................
Q33: Which of these could be the activation energy of the reaction shown in the
Reaction profile activity?
a)
b)
c)
d)
-22 kJ
-55 kJ
+55 kJ
+90 kJ
..........................................
Q34: Explain what will happen to the rate of the reaction if the activation energy is
decreased.
..........................................
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Earlier, the activation energy was introduced as the minimum kinetic energy required
by colliding molecules before reaction can occur, as shown in the following energy
distribution diagram.
Figure 3.7: Energy distribution graph
..........................................
Q35: Which letter shows the molecules with least energy?
a)
b)
c)
d)
A
B
C
D
..........................................
Q36: Which letter shows molecules with enough energy to react?
a)
b)
c)
d)
A
B
C
D
..........................................
Q37: What will happen to the number of molecules with enough energy to react if the
activation energy is decreased?
a) It will decrease.
b) It will stay the same.
c) It will increase.
..........................................
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TOPIC 3. CATALYSIS
The two definitions of activation energy are quite compatible. As colliding molecules get
closer together, they begin to repel (potential energy increases) and slow down (kinetic
energy decreases) as kinetic energy is converted into potential energy. If the colliding
molecules have enough kinetic energy, then enough potential energy will be produced
to form the activated complex and so cause a reaction. A successful collision will have
taken place.
Key point
The activated complex is an unstable arrangement of atoms formed during a
reaction at the maximum of a potential energy barrier. The activation energy
is energy required by colliding molecules to form the activated complex. The
activation energy can be calculated from potential energy diagrams.
3.6.2
How catalysts work
How catalysts work
Go online
The following activity uses kinetic energy distribution graphs and a potential energy
diagrams to show the effect of a catalyst on the reaction rate.
Without a catalyst
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With a catalyst
Q38: Without a catalyst, which of the following is true?
a)
b)
c)
d)
High activation energy, lots of successful collisions.
High activation energy, few successful collisions.
Low activation energy, lots of successful collisions.
Low activation energy, few successful collisions.
..........................................
Q39: With a catalyst, which of the following is true?
a)
b)
c)
d)
High activation energy, lots of successful collisions.
High activation energy, few successful collisions.
Low activation energy, lots of successful collisions.
Low activation energy, few successful collisions.
..........................................
Q40: What effect does a catalyst have on the reaction rate?
a)
b)
c)
d)
It speeds up the reaction by adding extra energy.
It speeds up the reaction by making the reaction exothermic.
It speeds up the reaction by lowering the activation energy.
It speeds up the reaction by raising the activation energy.
..........................................
Q41: What effect does a catalyst have on the enthalpy change?
a) Increases it.
b) No effect.
c) Decreases it.
..........................................
© H ERIOT-WATT U NIVERSITY
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TOPIC 3. CATALYSIS
The catalyst provides an alternative pathway for the reaction. Often, the catalyst
provides a surface on which the reaction occurs. The bonds in the reactant molecules
are weakened when the molecules are adsorbed onto the surface. This pathway has a
lower activation energy and so the reaction is faster.
Mechanism of catalysis
Go online
This activity shows how a catalyst might speed up the gas phase reaction between two
reactants A and B, as discussed earlier. The equations show the possible steps.
In the uncatalysed reaction, A2 molecules and B2 molecules must collide with sufficient
energy to form the activated complex.
In the catalysed reaction, step 1 involves a lower energy collision between the catalyst
(Y) and an A2 molecule. In step 2, the intermediate, A 2 Y, then undergoes another low
energy collision with B2 to form the products. Y is regenerated in step 2 and is free to
repeat the process.
A potential energy diagram can be drawn for the process.
Figure 3.8: PE diagram with catalyst
© H ERIOT-WATT U NIVERSITY
TOPIC 3. CATALYSIS
59
..........................................
Q42: What indicates the activation energy of the catalysed formation of AB?
a) X
b) W
..........................................
Q43: What indicates the activation energy of the uncatalysed decomposition of AB?
a)
b)
c)
d)
1
2
3
4
..........................................
Q44: Which arrow shows the enthalpy change for the catalysed formation of AB?
a)
b)
c)
d)
1
2
3
4
..........................................
Q45: What effect has the catalyst had on the activation energy of the reverse reaction?
a) Increases it
b) Decreases it
c) No effect
..........................................
Q46: By considering the equations, which letter shows the stage in the uncatalysed
reaction at which you are most likely to find A2 Y + B2 ?
a) X
b) W
..........................................
Q47: By considering the equations which letter in Figure 3.8 shows the stage in the
catalysed reaction at which you are most likely to find the following species?
a) X
b) W
© H ERIOT-WATT U NIVERSITY
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TOPIC 3. CATALYSIS
..........................................
..........................................
Key point
Catalysts speed up reactions by providing an alternative pathway which has a
lower activation energy. A potential energy diagram can be used to show the
effect of a catalyst on activation energy.
3.7
Summary
Summary
•
In general, the lower the activation energy the faster the reaction.
•
Catalysts are substances which speed up chemical reactions without being
used up in the process. They are widely used in industrial processes.
•
Catalysts allow chemical reactions to occur more quickly at lower
temperatures and so reduce energy costs.
•
Catalysts work by the adsorption of reactant molecules onto the surface
of the catalyst with consequent weakening of the bonds in the reactant
molecules.
•
Catalytic converters are fitted to cars to catalyse the conversion of
poisonous carbon monoxide and oxides of nitrogen to carbon dioxide and
nitrogen. Cars with catalytic converters can only use 'lead-free' petrol to
prevent poisoning of the catalyst.
•
Enzymes catalyse the chemical reactions which take place in the living cells
of plants and animals. They are also widely used in industrial processes.
•
Cobalt(II) chloride speeds up the reaction between tartrate ions and
hydrogen peroxide. Colour changes clearly show that the catalyst takes
part in the reaction and is chemically unchanged at the end of the reaction.
•
The activated complex is an unstable arrangement of atoms formed during
a reaction at the maximum of a potential energy barrier.
•
A potential energy diagram can be used to show the energy pathway for a
reaction.
•
The enthalpy change, which can be calculated from the potential energy
diagram, is the energy difference between products and reactants.
•
The enthalpy change has a negative value for exothermic reactions, which
causes heat energy to be released to the surroundings.
© H ERIOT-WATT U NIVERSITY
TOPIC 3. CATALYSIS
Summary continued
3.8
•
•
The enthalpy change has a positive value for endothermic reactions, which
causes absorption of heat energy from the surroundings.
•
The activation energy is the energy required by colliding molecules to form
the activated complex.
•
The activation energy can be calculated from potential energy diagrams.
•
Catalysts speed up reactions by providing an alternative pathway which has
a lower activation energy.
•
A potential energy diagram can be used to show the effect of a catalyst on
activation energy.
Resources
Higher Chemistry for CfE: J Anderson, E Allan and J Harris, Hodder Gibson,
ISBN 978-1444167528
© H ERIOT-WATT U NIVERSITY
61
62
TOPIC 3. CATALYSIS
3.9
End of topic test
End of Topic 3 test
Q48:
Go online
In the shaded area on the graph:
a) molecules always have energy above the activation energy.
b) there are no molecules with kinetic energy above the average.
c) there are no molecules with energy above the activation energy.
d) collisions between molecules are always successful in forming products.
..........................................
Q49: Activation energy is:
a)
b)
c)
d)
the minimum kinetic energy required by colliding molecules for a reaction to occur.
the maximum kinetic energy colliding molecules can reach for a reaction to occur.
the temperature at which a reaction will take place.
the temperature at which rate of reaction will be greatest.
..........................................
© H ERIOT-WATT U NIVERSITY
TOPIC 3. CATALYSIS
The progress of reactions can be followed by energy diagrams.
© H ERIOT-WATT U NIVERSITY
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64
TOPIC 3. CATALYSIS
Q50:
Which graph represents the catalysed version of the reaction in diagram F?
a) A
b) B
c) C
d) D
e) E
..........................................
Q51:
Which graph represents the reaction with the highest activation energy?
a) A
b) B
c) C
d) D
e) E
f) F
..........................................
Q52:
Which diagram represents the reaction with an enthalpy change of -200 kJ mol -1 ?
a) A
b) B
c) C
d) D
e) E
f) F
..........................................
Q53:
What is an activated complex?
a) A very unstable intermediate is formed when bonds within the reactant molecules
begin to break and new bonds begin to form.
b) The energy required for a reaction to take place.
c) When the surface of a catalyst forms weak bonds with reacting molecules.
d) A compound containing a transition metal.
..........................................
© H ERIOT-WATT U NIVERSITY
TOPIC 3. CATALYSIS
65
Use the diagram to answer the following questions.
Q54: What type of reaction does this graph represent?
..........................................
Q55: What is the activation energy value for the forward reaction?
..........................................
Q56: What is the activation energy value for the reverse reaction?
..........................................
Q57: What is the value of the enthalpy change for the forward reaction?
..........................................
Q58: What is the value of the enthalpy change for the reverse reaction?
..........................................
© H ERIOT-WATT U NIVERSITY
66
TOPIC 3. CATALYSIS
Q59:
a) Name the energy changes represented by the numbered intervals on the energy
diagram shown.
b) The reaction shown is exothermic, how can you tell this from the shape of the
graph?
..........................................
..........................................
© H ERIOT-WATT U NIVERSITY
67
Topic 4
The Periodic Table
Contents
4.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
69
4.2 Arrangement of elements in the Periodic Table: Introduction . . . . . . . . . . .
4.3 History of the Periodic Table . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
69
70
4.4 Trends and patterns (periodicity) . . . . . . . . . . . . . . . . . . . . . . . . . .
4.4.1 Melting and boiling points . . . . . . . . . . . . . . . . . . . . . . . . . .
75
75
4.4.2 Atomic size . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
4.5 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
77
79
4.6 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
4.7 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
79
80
Prerequisite knowledge
Before you begin this topic, you should already know that:
•
atoms contain protons, neutrons and electrons, each with a specific charge, mass
and position within the atom - the number of protons defines an element and is
known as the atomic number (National 4, Unit 1);
•
or have knowledge of: sub-atomic particles, their charge, mass and position within
the atom, the structure of the Periodic Table, groups, periods and atomic number
(National 5, Unit 1);
•
all matter is made of atoms - when a substance contains only one kind of atom it
is known as an element (National 4, Unit 1);
•
elements are arranged in the Periodic Table in order of increasing atomic number
- elements with similar chemical properties are grouped together (National 4, Unit
1);
•
or be familiar with the seven diatomic elements (National 5, Unit 1);
•
elements can be categorised as metals and non-metals (National 4, Unit 1);
68
TOPIC 4. THE PERIODIC TABLE
Learning objectives
By the end of this topic, you should be able to:
•
state that elements are arranged in the Periodic Table in order of increasing atomic
number;
•
explain that the Periodic Table allows chemists to make accurate predictions of
physical properties and chemical behaviour for any element based on its position;
•
state that there are periodic variations in the densities, melting points and boiling
points of the elements across a Period and down a Group.
© H ERIOT-WATT U NIVERSITY
TOPIC 4. THE PERIODIC TABLE
4.1
69
Prior knowledge
Test your prior knowledge
Q1: What does the number of protons in an atom determine?
a)
b)
c)
d)
Go online
Atomic number
Mass number
Element name
Atom charge
..........................................
Q2: Name the seven diatomic elements.
..........................................
Q3: Complete the following table;
Particle
Charge
Mass
Location
Proton
Neutron
Electron
..........................................
Q4: What is a row in the Periodic Table is called?
a)
b)
c)
d)
Group
Period
Row
Set
..........................................
Q5: What is a column in the Periodic Table is called?
a)
b)
c)
d)
Column
Group
Period
Set
..........................................
4.2
Arrangement of elements in the Periodic Table:
Introduction
As more and more elements were discovered and their properties investigated, chemists
showed a natural desire to simplify the study of the elements by organising them
according to similarities in their chemical behaviour. Eventually this resulted in the
modern Periodic Table. If you understand how the Periodic Table is constructed, you
will realise that it contains a huge amount of information stored in a very compact form.
This makes it a vital resource for any chemist.
© H ERIOT-WATT U NIVERSITY
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TOPIC 4. THE PERIODIC TABLE
4.3
History of the Periodic Table
Background Information
In 1800, about 33 elements were known but there was no obvious pattern or relationship
between them.
By 1830, a further 20 or so elements had been discovered and some similarities in
properties within small groups of elements was recognised. The German chemist,
Johan Wolfgang Dobereiner, made a tentative connection between chemical behaviour
and the atomic masses of certain groups of elements, each containing three elements,
which he called 'triads'.
Figure 4.1: Triads
..........................................
In each case, the atomic mass of the central element was approximately the mean of
the other two.
The next significant development occurred in 1866 when the English chemist, John
Newlands, published a paper on the 'Law of Octaves'.
Newlands' octaves
Go online
Figure 4.2 shows how Newlands organised the first 14 elements. Study the information
carefully and then answer the questions which follow.
Figure 4.2
..........................................
© H ERIOT-WATT U NIVERSITY
TOPIC 4. THE PERIODIC TABLE
71
Q6: What property did Newlands use to put the elements in order?
a)
b)
c)
d)
Name (i.e. alphabetical order)
Colour
Atomic mass
Atomic number
..........................................
Q7: What property is shared by the elements in the first column (H and F)?
a)
b)
c)
d)
Both are gases.
Both are green.
Both are reactive metals.
None - they are completely different.
..........................................
Q8: What property is shared by the elements in the second column (Li and Na)?
a)
b)
c)
d)
Both are gases.
Both are purple.
Both are reactive metals.
None - they are completely different.
..........................................
Q9: Why do you think Newlands referred to these as 'octaves'? (Hint: Think about
musical scales.)
..........................................
'Soh-fah' so good! Now look at the second arrangement in which the next seven
elements have been added and then answer the questions which follow.
Figure 4.3
..........................................
© H ERIOT-WATT U NIVERSITY
72
TOPIC 4. THE PERIODIC TABLE
Q10: At first sight, the pattern seems to continue. For how many of the further elements
does the pattern work?
..........................................
Q11: Can you think of a reason why chromium, manganese and iron do not fit with the
elements immediately above them?
..........................................
..........................................
Key point
When Newlands arranged the elements in order of increasing atomic mass,
similar chemical properties were repeated with every eighth element, but this only
worked for the first 17 elements.
Newlands' ideas were subjected to ridicule and it was even suggested unkindly that he
would get better agreement if he arranged the elements in alphabetical order. However,
he was on the right lines.
The Modern Periodic Table
The major breakthrough, indeed one of the most important advances in all chemistry,
was provided by Dmitri Ivanovitch Mendeleev in 1869. Like Newlands, he organised the
elements in order of increasing atomic mass, but there were other important differences.
Mendeleev's Periodic Table
Go online
Mendeleev organised the first 45 of the 62 elements known at that time. Figure 4.4
shows his final arrangement.
Figure 4.4: Constructing Mendeleev's table
© H ERIOT-WATT U NIVERSITY
TOPIC 4. THE PERIODIC TABLE
..........................................
Q12: Unlike Newlands, Mendeleev left spaces (shown as *) in his table. Why?
..........................................
© H ERIOT-WATT U NIVERSITY
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74
TOPIC 4. THE PERIODIC TABLE
Q13: Explain why Mendeleev swapped the positions of the elements iodine and
tellurium.
..........................................
Q14: Name the Group of elements which is found in the modern Periodic Table but is
absent from both Newlands' and Mendeleev's tables.
..........................................
Key point
Mendeleev organised the elements in order of increasing atomic mass (mostly) in
conjunction with similar chemical properties, leaving gaps for elements yet to be
discovered.
..........................................
The modern Periodic Table
The relationship between the modern Periodic Table and that of Mendeleev can be seen.
Go online
Figure 4.5
..........................................
Mendeleev was so confident that his table was correct that he predicted the properties
of some of the undiscovered elements. His predictions proved to be very accurate when
these elements were finally isolated, providing startling vindication of his theory.
The modern Periodic Table can be shown in a variety of different ways. The usual form
is the one used in the recommended data booklet. The resources at the end of this topic
contain addresses for a number of web-sites which contain interactive Periodic Tables.
One of the most useful is http://www.webelements.com/.
© H ERIOT-WATT U NIVERSITY
TOPIC 4. THE PERIODIC TABLE
75
Q15: In the modern Periodic Table, the elements are not arranged in order of increasing
atomic mass. What is used instead?
..........................................
..........................................
4.4
Trends and patterns (periodicity)
When the elements are arranged in order of increasing atomic number, many properties
vary in a regular way. As you move across a Period from left to right, a pattern emerges.
A similar pattern appears on crossing the next Period. Properties which behave in this
way are said to be periodic. is the regular recurrence of similar element properties.
Throughout this topic, we will concentrate on the properties of the main group elements
and in particular on the elements in Periods 2 and 3. Much of this work involves the
collection of data, presentation of data in graphical form and the interpretation of such
graphs.
4.4.1
Melting and boiling points
Melting and boiling points across a Period
Figure 4.6: Period 2 and 3 (melting point and boiling point)
Go online
..........................................
© H ERIOT-WATT U NIVERSITY
76
TOPIC 4. THE PERIODIC TABLE
Melting points and boiling points are also periodic properties and depend on the strength
of the forces that exist between the particles which make up a substance.
Q16: Which element in Period 2 which has the strongest forces between its atoms in
the solid state?
..........................................
Q17: Which of the Period 3 elements is the easiest to boil?
..........................................
Q18: Of all the elements shown in the data booklet, which element which has the
weakest forces between its atoms?
..........................................
..........................................
Melting and boiling points down a Group
A more obvious trend can be seen on descending a Group.
Go online
Figure 4.7: Alkali metals (melting point and boiling point)
..........................................
© H ERIOT-WATT U NIVERSITY
TOPIC 4. THE PERIODIC TABLE
77
Q19: Which alkali metal has the strongest forces between its atoms?
..........................................
Q20: Which alkali metal has the weakest forces between its atoms?
..........................................
Q21: What happens to the forces between the atoms on descending Group 1?
..........................................
..........................................
The variation in melting point and boiling point will be explained when the bonding and
structure of the first 20 elements is considered later in this topic.
Key point
There are periodic variations in the densities, melting points and boiling points of
the elements across a Period and down a Group.
4.4.2
Atomic size
The size of an atom is determined by the amount of space taken up by the electrons
and so must be connected to the electron arrangement of the atom. The electron
arrangement is itself a periodic property. Use the following activity (Figure 4.8) to explain
the trends in covalent radius.
Atomic size
Figure 4.8: Trends in nuclear charge and electron arrangement
Go online
© H ERIOT-WATT U NIVERSITY
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TOPIC 4. THE PERIODIC TABLE
..........................................
Q22: Which of the following provides the best reason for the increase in covalent radius
on going down a Group?
a)
b)
c)
d)
The number of protons increases.
The number of electrons increases.
The number of electron shells increases.
The number of neutrons increases.
..........................................
Q23: Which of the following provides the best reason for the decrease in covalent radius
on going from left to right across a Period?
a)
b)
c)
d)
The number of electrons increases.
The number of electron shells increases.
The number of neutrons increases.
The number of protons increases.
..........................................
© H ERIOT-WATT U NIVERSITY
TOPIC 4. THE PERIODIC TABLE
..........................................
Key point
•
The covalent radius decreases across a Period because the increase in
nuclear charge attracts the electrons more strongly.
•
The covalent radius increases on going down a Group as the number of
occupied electron shells increases.
4.5
Summary
Summary
4.6
•
•
Elements are arranged in the Periodic Table in order of increasing atomic
number.
•
The Periodic Table allows chemists to make accurate predictions of physical
properties and chemical behaviour for any element based on its position.
•
There are periodic variations in the densities, melting points and boiling
points of the elements across a Period and down a Group.
Resources
Higher Chemistry for CfE: J Anderson, E Allan and J Harris, Hodder Gibson,
ISBN 978-1444167528
© H ERIOT-WATT U NIVERSITY
79
80
TOPIC 4. THE PERIODIC TABLE
4.7
End of topic test
End of Topic 4 test
Q24:
Go online
Menedeleev is famous for producing the Periodic Table on which the modern version is
based.
Which of the following statements is true?
a)
b)
c)
d)
Mendeleev organised the elements in order of their atomic number.
Mendeleev left gaps because some elements did not fit the pattern of reactivity.
Mendeleev left gaps for elements which had not yet been discovered.
Mendeleev swapped some elements round so that their atomic masses fitted the
pattern.
..........................................
Q25:
Which of the following statements about the Periodic Table are true?
a) There is a steady increase in melting point across a period from left to right.
b) There is a steady decrease in density on going down Group 1.
c) There is a steady decrease in atomic size across a period from left to right.
d) There is a decrease in first ionisation energy on going down Group 0.
e) There is a decrease and then an increase in boiling point on crossing a period from
left to right.
f) There is an increase in electronegativity on going down Group 7.
..........................................
Q26:
What does the number of protons in an atom determine?
a)
b)
c)
d)
Atomic number
Mass number
Element name
Atom charge
..........................................
Q27:
Name the seven diatomic elements.
..........................................
© H ERIOT-WATT U NIVERSITY
TOPIC 4. THE PERIODIC TABLE
Q28:
What is Group 1 in the Periodic Table called?
a)
b)
c)
d)
The halogens.
The alkali metals.
The noble gases.
Transition metals.
..........................................
Q29:
What is Group 7 in the Periodic Table called?
a)
b)
c)
d)
The halogens.
The alkali metals.
The noble gases.
Transition metals.
..........................................
Q30:
What is Group 0 in the Periodic Table called?
a)
b)
c)
d)
The halogens.
The alkali metals.
The noble gases.
Transition metals.
..........................................
Q31:
Elements in the same group in the Periodic Table have the same:
a)
b)
c)
d)
number of occupied energy shells.
density.
number of outer electrons.
number of protons.
..........................................
Q32:
Where in the Periodic Table are non-metals found?
a)
b)
c)
d)
The top.
The bottom.
The right hand side.
The left hand side.
..........................................
Periodicity
© H ERIOT-WATT U NIVERSITY
81
82
TOPIC 4. THE PERIODIC TABLE
© H ERIOT-WATT U NIVERSITY
83
Topic 5
Bonding and structure
Contents
5.1 Prior knowledge . . . . . . . . . . . . .
5.2 Introduction . . . . . . . . . . . . . . .
5.3 Metallic bonds . . . . . . . . . . . . . .
5.3.1 Metallic structures . . . . . . .
5.4 Covalent bonds . . . . . . . . . . . . .
5.4.1 Covalent molecular structures
5.4.2 Covalent network structures .
5.5 Ionic bonds . . . . . . . . . . . . . . .
5.5.1 Ionic lattice structures . . . . .
5.6 Summary . . . . . . . . . . . . . . . .
5.7 Resources . . . . . . . . . . . . . . . .
5.8 End of topic test . . . . . . . . . . . . .
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85
86
90
91
92
94
98
102
105
107
107
108
Prerequisite knowledge
Before you begin this topic, you should already know that:
•
elements can be categorised as metals and non-metals (National 4, Unit 1);
•
experimental procedures are required to confirm the type of bonding present in a
substance (National 5, Unit 1);
•
metallic bonding can explain the conductivity of metals (National 5, Unit 3);
•
covalent compounds form when non-metal atoms form covalent bonds by sharing
their outer electrons (National 4, Unit 1);
•
covalent molecular compounds have low melting and boiling points - as a result,
they can be found in any state at room temperature (National 4, Unit 1);
•
in a covalent bond, the shared pair of electrons is attracted to the nuclei of the two
bonded atoms (National 5, Unit 1);
•
more than one bond can be formed between atoms leading to double and triple
covalent bonds (National 5, Unit 1);
•
covalent substances can form either discrete molecular or giant network structures
(National 5, Unit 1);
84
TOPIC 5. BONDING AND STRUCTURE
•
diagrams show how outer electrons are shared to form the covalent bond(s) in a
molecule and the shape of simple two-element compounds (National 5, Unit 1);
•
covalent molecular substances have low melting and boiling points due to only
weak forces of attraction between molecules being broken (National 5, Unit 1);
•
giant covalent network structures have very high melting and boiling points
because the network of strong covalent bonds must be broken (National 5, Unit 1);
•
when there is an imbalance in the number of positive protons and electrons, the
particle is known as an ion (National 5, Unit 1);
•
ionic bonds are the electrostatic attraction between positive and negative ions ionic compounds form lattice structures of oppositely charged ions (National 5,
Unit 1);
•
ionic compounds have high melting and boiling points because strong ionic bonds
must be broken in order to break down the lattice - dissolving also breaks down
the lattice structure (National 5, Unit 1);
•
ionic compounds have high melting and boiling points - as a result, they are found
in the solid state at room temperature (National 4, Unit 1);
•
ionic compounds form when metal atoms join to non-metal atoms by transferring
electron(s) from the metal to the non-metal - the resulting charged particles are
called ions and an ionic bond is the attraction of the oppositely charged ions
(National 4, Unit 1);
•
ionic compounds conduct electricity, only when molten or in solution due to the
breakdown of the lattice resulting in the ions being free to move (National 5, Unit
1).
Learning objectives
By the end of this topic, you should be able to:
•
state that the first 20 elements in the Periodic Table can be categorised according
to bonding and structure:
◦
metallic Li, Be, Na, Mg, Al, K, Ca;
◦
covalent molecular H2 , N2 , O2 , F2 , Cl2 , P4 , S8 and fullerenes (eg C 60 );
◦
covalent network B, C, Si (diamond, graphite and the element silicon and
silicon dioxide) monatomic (noble gases);
explain each of these types of substance in terms of bonding and structure.
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TOPIC 5. BONDING AND STRUCTURE
5.1
85
Prior knowledge
Test your prior knowledge
Q1: Covalent bonding involves:
a)
b)
c)
d)
a shared pair of electrons.
transfer of electrons.
delocalised electrons.
gaining electrons.
..........................................
Q2: Metals can conduct electricity in their solid state because of their:
a)
b)
c)
d)
ions.
positive cores.
lattice structure.
delocalised electrons.
..........................................
Q3: Ionic bonding involves:
a)
b)
c)
d)
a shared pair of electrons.
transfer of electrons.
delocalised electrons.
protons only.
..........................................
Q4: One property of covalent networks is that they:
a)
b)
c)
d)
have high melting and boiling points.
have low melting and boiling points.
are soluble in water.
conduct electricity when molten or in solution.
..........................................
Q5: One property of ionic substances is that they:
a)
b)
c)
d)
are all white in colour.
have low melting and boiling points.
are insoluble in water.
conduct electricity when molten or in solution.
..........................................
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TOPIC 5. BONDING AND STRUCTURE
5.2
Introduction
Most elements are found on Earth as compounds.
Q6:
Can you think of any elements that are found 'free', not as compounds?
..........................................
Comparing the elements iron and tantalum
Sometimes ores can be simple compounds (e.g. haematite, an oxide of iron, Fe 2 O3 );
some are complex minerals (e.g. columbite, containing tantalum, (Fe,Mn)(Ta,Nb) 2O6 ).
Haematite (http:/ / bit.ly/ 2aK AD0W by http:/ / commons.w ikimed ia.or g/ w iki/ User :
Sailko, is licensed under http:/ / cr eativecommons.or g/ licenses/ by/ 2.0/ d eed .en)
Columbite is a black mineral group that is an ore of niobium (http:/ / commons.w ikime
d ia.or g/ w iki/ F ile:Columbite-7 5444.jpg, by http:/ / w w w .ir ocks.com, is licensed
under http:/ / cr eativecommons.or g/ licenses/ by/ 2.0/ d eed .en)
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TOPIC 5. BONDING AND STRUCTURE
Iron is also abundant in the Earth's crust (5%); tantalum is rare (1 - 2 parts per million).
The extraction of iron from its ores is an established, well-understood process; the
extraction of tantalum is complex, mainly because it occurs with the very similar metal,
niobium.
Q7: Niobium is used for a very special purpose in space rockets. Can you find out
what it is for?
..........................................
Iron is used for a vast variety of purposes; tantalum is almost exclusively used to make
high-performance capacitors in electronic equipment, for example, mobile phones.
Despite these contrasts, you should always remember that all the 'chemicals' we
humans use have to be extracted from the finite resources present in the Earth.
A variety of chemical processes are used to extract elements from their compounds,
but the choice depends on the bonding and structure of the materials containing the
element.
Bonding and structure in elements and simple compounds
Everything we see around us is made from fewer than 100 different types of atoms
chemically bonded in various ways to produce a multitude of different molecules.
The different types of chemical bonding determine the structure that elements and
compounds adopt. In turn the structure, together with the size of intermolecular forces
mainly determines the physical and chemical properties possessed by these materials.
Bonding and structure will be studied in this topic. Intermolecular forces (the forces that
exist between molecules) and properties are studied in a later topic.
Examples of some substances with different properties, depending on different bonding
and structures are shown in the images below.
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TOPIC 5. BONDING AND STRUCTURE
Iron screws - a typical metal
Water - a volatile liquid
Nitrogen dioxide - a brown gas
Sodium chloride - a white crystalline solid
No atoms, except those of the noble gases, exist in isolation under normal conditions.
They all interact in one way or another to form more stable structures.
Atoms consist of a tiny positively charged nucleus, with different numbers of electrons,
in shells of increasing energy levels, around it.
The following activity shows the final electronic structure of a sodium atom.
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89
Construction of a sodium atom
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Construction of a sodium atom
In the case of the atoms of noble gases, all shells are completely filled with electrons.
This arrangement is particularly stable, which makes noble gases unreactive.
Atoms of other elements combine in ways which try to achieve this stable noble gas
arrangement of electrons, in other words, to become isoelectronic with a noble gas.
Q8: Look at the Periodic Table in the SQA data booklet (page 2). Which noble gas has
an electronic structure closest to chlorine?
..........................................
Q9: What must happen to a chlorine atom to get the noble gas electronic
arrangement?
..........................................
Q10: What must happen to the sodium atom for it to have a noble gas electronic
structure?
..........................................
Q11: What will the sodium atom become?
..........................................
Q12: Which noble gas has the same electronic structure as a sodium ion?
..........................................
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5.3
Metallic bonds
Metals occur to the left and centre of the Periodic Table (see Topic 4). When atoms of
metals come together, the most stable condition is for them to 'pool' their outer electrons
and become, in effect, a regular arrangement of fixed metal ions in a 'sea' of delocalised
electrons shared by all the ions.
For example, in metallic sodium each atom will donate an electron to the delocalised
pool, and will achieve the stable sodium ion structure (isoelectronic with neon).
Metallic bonds
The following shows the structure of metallic sodium.
Go online
Such an array of positively charged ions would normally split apart by mutual
repulsions, but the influence of the negative electrons holds the particles in the
structure so well that most metals are hard solids with high melting and boiling points.
Q13: There are exceptions. Can you think of soft metals, and one that is liquid under
normal conditions?
..........................................
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5.3.1
91
Metallic structures
In metallic bonding, the delocalised electrons are able to migrate freely throughout the
metal (making them good conductors of electricity), but the positive ions have a regular
3D structure known as a lattice.
Part of the structure of the giant lattice for copper is shown below.
Copper lattice
Go online
Q14: The electrons in a metallic bond are said to be:
a)
b)
c)
d)
loose.
ionised.
delocalised.
hard.
..........................................
Q15: What do the ions in a metallic bond form into?
a)
b)
c)
d)
Gel
Grid
Bar
Lattice
..........................................
Q16: Suggest a reason why aluminium is a better conductor of electricity than sodium.
..........................................
Key point
Metallic bonding is the electrostatic force of attraction between positively charged
ions and delocalised outer electrons.
A metallic structure consists of a giant lattice of positively charged ions and
delocalised outer electrons.
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5.4
Covalent bonds
You should already have studied covalent bonds, which are the most common type of
bond. This section revises your knowledge.
Covalent bonds form when atoms share two electrons, enabling both atoms to complete
their valency shells.
The activity below illustrates two separate chlorine atoms and the sharing of electrons
to form a chlorine molecule.
Covalent bond formation
Go online
..........................................
This sharing of two electrons (one from each atom) to complete an outer shell of
electrons is called a covalent bond.
In most cases the outer shell will contain 8 electrons, but for hydrogen only two are
present.
The following diagram shows the electrostatic forces in a hydrogen molecule. The
positively charged nuclei will repel each other, as will the negatively charged electrons;
but these forces are more than balanced by the attraction between the nuclei and
electrons.
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Attractive and repulsive forces in a hydrogen molecule
When atoms require more than one electron to complete their outer shell, they can
share two or three electrons to make a double or triple covalent bond, or share with
more than one other atom.
The diagrams show the shared electrons in a molecule of oxygen (O 2 ) and ammonia
(NH3 ). Notice the easier way to show a covalent bond with a line for each bonding
(shared) pair of electrons.
Covalent bonding in oxygen and ammonia
The dark pairs of electrons in above diagram are the bonding electrons, shared
between the atoms. The light pairs of electrons, completing the shells around the
atoms, are called lone pairs.
Q17: How many electrons are involved in the double bond in an oxygen molecule?
..........................................
Q18: How many lone pairs are there in an ammonia molecule?
..........................................
Q19: Which two words complete this sentence?
A covalent bond is formed when two atoms
configuration.
achieve a noble gas
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a pair of electrons, so that each can
94
TOPIC 5. BONDING AND STRUCTURE
a)
b)
c)
d)
share, proton
share, electron
donate, atomic
donate, ionic
..........................................
Q20: The dominant attractive force in a covalent bond is between:
a)
b)
c)
d)
the positively charged nuclei.
the negatively charged shared electrons.
negatively charged nuclei and positively charged shared electrons.
positively charged nuclei and negatively charged shared electrons.
..........................................
Q21: Draw the structure of iodine chloride. Look at the SQA data booklet page 2 to get
the electron structures of the atoms.
..........................................
Key point
Atoms in a covalent bond are held together by electrostatic forces of attraction
between positively charged nuclei and negatively charged shared electrons.
5.4.1
Covalent molecular structures
Most covalent substances (for example all those mentioned in the previous section) exist
as discrete molecules. There are strong covalent bonds binding the atoms together in
the molecule, but much weaker forces between these molecules. You will study these
forces in detail later (refer to the 'Intermolecular Forces' topic). Consequently, many
small covalent molecules are gases (e.g. fluorine, oxygen, nitrogen, carbon dioxide,
sulfur dioxide etc). Larger molecules make structures which are liquids or low melting
point solids (think of candle wax, with molecular mass around 500, but a low melting
point about 70 ◦ C).
The molecules of nitrogen, oxygen, and the halogens consist of two atoms, they are
diatomic. (You always write them N2 , O2 etc.)
The non-metals phosphorus and sulfur form larger molecules as described in the
following activities.
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White phosphorus
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White phosphorus
Phosphorus (in the common form of 'white' phosphorus) forms P 4 molecules, with each
phosphorus atom at the corner of a tetrahedron.
Tetrahedron of four atoms in white phosphorus
..........................................
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TOPIC 5. BONDING AND STRUCTURE
Red phosphorus
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Red phosphorus
Phosphorus can also form a much less reactive (and less toxic) form, 'red' phosphorus,
which is used in making matches. Its structure consists of chains of these tetrahedra.
The structure is shown below.
These P4 tetrahedra are quite stable, so that it is usual to write "P 4 " in equations where
phosphorus is involved. It is even retained in some compounds - the oxides are
molecules of formula P4 O6 and P4 O10 .
Red phosphorus structure
..........................................
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Sulfur
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Sulfur
Molecules of sulfur (in the normal solid state as rhombic sulfur) consist of puckered
rings of eight sulfur atoms, written S8 .
Rings of eight sulfur atoms
In this case, using S8 in equations would be rather cumbersome so sulfur in equations
is written 'S'.
Allotropes of sulfur
In addition to the eight-membered rings in the rhombohedral, α-sulfur found in common
'flowers of sulfur', there are many other allotropes of sulfur. At temperatures above
95◦ C, the S8 rings pack to form monoclinic crystals of β-sulfur.
Sulfur can also form rings and chains with any number of S atoms from 2 to 20.
S2 is a gaseous form, an analogue of O 2 ; the blue colour of burning sulfur is due to the
emission of light by S2 molecules produced in the flame.
S3 is cherry red and has a structure like ozone, O3 .
N.B. The recently discovered molecular structures of carbon (the fullerenes) are
discussed in the next section, after the other common forms of carbon.
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..........................................
5.4.2
Covalent network structures
In some elements and compounds the covalent bonds are not limited only to those within
the molecules, but all the atoms are held to others by strong covalent bonds.
An example of this is the structure of diamond. It consists solely of a network of carbon
atoms held together by covalent bonds. Each carbon atom makes covalent bonds to
four different carbon atoms, which each bond to three more carbons, and so on,
binding the whole structure together, as shown in the following activity.
This covalent network structure has the characteristic properties of hardness and a
high melting point, both due to the strong bonding throughout the structure.
Diamond
Go online
Arrangement of carbon atoms in diamond
Q22: How many pairs of electrons are around each carbon atom?
..........................................
Q23: What type of bond holds the carbon atoms together in diamond?
a) Single covalent.
b) Double covalent.
c) Triple covalent.
..........................................
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The following activities allow you explore other covalent network structures.
Silicon dioxide
Look at the following model, which compares the structure of silica (quartz, SiO 2 ) with
diamond. Again, this consists of a network of covalent bonds, with each silicon atom
bonded to four oxygen atoms which in turn bond to another silicon atom, to form a
strongly-bound network. Quartz is hard and has a high melting point, like diamond.
Go online
Silica and diamond
..........................................
Graphite structure
Carbon has another interesting structure, graphite. In this case, carbon atoms
covalently bond with only three other atoms to form a sheet of hexagonal rings.
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TOPIC 5. BONDING AND STRUCTURE
The distance between carbon atoms within these sheets is 142 pm.
Q24: This is less than the bond length in diamond (154 pm) and suggests that the bond
strength is:
a) weaker.
b) the same.
c) stronger.
..........................................
Q25: In these sheets, each carbon atom is bonded to only three others. How many
bonding electrons surround each carbon atom in this case?
..........................................
Graphene
It is only recently that single layers of carbon atoms as found in graphite have been
isolated in sufficient quantities to study their properties. The Nobel Prize for physics for
2010 was awarded to Andre Geim and Konstantin Novoselov for work on graphene, as
it was called, which was performed at Manchester University in 2004. They found that
these isolated single layers have very unusual electrical and physical properties, which
have started research into new electronic applications.
Graphite
The carbon atoms have four electrons in their outer shell. Only three are used to bond
with three other atoms, leaving a fourth electron. These sheets are layered together
and the carbon atoms donate their extra electron to a delocalised pool which holds the
sheets in a metallic type bonding.
Layers of carbon atoms in graphite
The distance between these sheets is 335 pm which indicates a weak bonding. These
weaker forces produce a material which is greasy (as the layers slip over each other)
and can be used as a lubricant, and the delocalised electron pool makes graphite a
good conductor of electricity.
..........................................
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101
Fullerenes
Another form of carbon (discovered in the 1980s) exists as covalent molecular
structures. These are the fullerenes.
The most stable of these has 60 carbon atoms arranged in 5- and 6-membered rings
forming a large sphere, looking like a football. It is called buckminsterfullerene after
the architect (Robert Buckminster Fuller) who creates geodesic domes resembling the
structure found in these carbon molecules.
Buckminsterfullerene
These fullerenes have interesting properties, which are currently being researched. For
example, the 'buckyball', as illustrated, will dissolve in benzene to form a red solution.
Q26: Why do you think buckminsterfullerene is soluble when diamond and graphite are
not?
..........................................
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TOPIC 5. BONDING AND STRUCTURE
Long 'nanotubes' of carbon have tensile strengths 50 to 100 times that of steel!
Why are carbon nanotubes so strong? The diagram shows only a small part of a tube.
When very many of these long molecules are combined, the result consists of very
large molecules of carbon atoms bonded to each other with strong bonds, similar to
those in sheets of graphite. These bonds are much stronger than metallic bonds in
steel.
..........................................
Key point
A covalent molecular structure consists of discrete molecules held together by
intermolecular forces (see next topic).
A covalent network structure consists of a giant lattice of covalently bonded
atoms.
5.5
Ionic bonds
Polar covalent bonds (you will study 'Polar covalent bonding' in a later topic) are formed
when the atoms involved in a bond have different attractions for the bonding electrons.
When two atoms have a large difference in their attraction for electrons (e.g. sodium and
chlorine), it is most energetically favourable for an electron in the metal to be donated
completely to the non-metal. This is the concept of electronegativity and will be fully
explained in the next topic. Both then achieve a noble gas structure as shown in the
following activity.
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103
Ionic bond formation
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Q27: How many outer electrons does a sodium atom have?
..........................................
Q28: How many outer electrons does a chlorine atom have?
..........................................
Q29: When a sodium atom loses its outer electron to a chlorine atom, how many outer
electrons do both ions then have?
..........................................
Q30: Why is this special?
..........................................
The original atoms change to ions (the metal forms a positive ion, the non-metal a
negative ion) which are held together by electrostatic attraction.
You can think of covalent, polar covalent and ionic bonds as forming a spectrum: bonds
become more polar, then ionic as the difference in electronegativity increases.
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TOPIC 5. BONDING AND STRUCTURE
The extent of covalent or ionic character depends mainly on the difference in
electronegativity. The properties of the different bonds are summarised in the following
table.
Bond
Covalent
Polar covalent
Ionic
Electrons
Equally shared
Unequally shared
Totally transferred
Charge
Distribution
None
Partial +ve and -ve
Fully charged +ve
and -ve ions
Other ionic compounds
Metals in Group 2, for example calcium, need to lose two electrons from their atoms to
achieve a stable noble gas structure. This is reflected in the relatively low first and
second ionisation energies. These elements will require two chlorine atoms to receive
these two electrons and maintain electrical neutrality so the ratio of metal to non metal
is 1:2 (e.g. [Ca2+ ][ Cl- ]2 ).
Metals from other groups in the Periodic Table can form ions with 3+ charge and other
non-metal ions with 2- and 3- charges. Higher charges than these are rare. When
elements form ionic bonds, the ratio of ions is such that positive and negative charges
always balance.
Further consideration of the nature of the ions produced by different elements will be
discussed in a later section.
Q31: Ionic bonds are formed between atoms which have a
electronegativity.
a)
b)
c)
d)
difference in
zero
small
large
0.2
..........................................
Q32: Caesium chloride is an ionic solid consisting of:
a)
b)
c)
d)
negative caesium ions and negative chloride ions.
negative caesium ions and positive chloride ions.
positive caesium ions and negative chloride ions.
positive caesium ions and positive chloride ions.
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..........................................
Q33: An ionic crystal consists of M2+ and X3- ions. What will be the ratio of ions (M 2+ :
X3- ) in the solid?
a)
b)
c)
d)
1:
2:
3:
3:
1
3
2
3
..........................................
5.5.1
Ionic lattice structures
The mutual electrostatic attraction of positive and negative ions is completely non
directional (unlike covalent bonds which are highly directed). When solid ionic
compounds form, a positive ion will be surrounded by several negative ions which in
turn will attract more positive ions. This process results in the formation of a lattice of
regularly arranged ions all held together by electrostatic forces. These forces act equally
in all directions , so that no one ion is attracted particularly to any other specific ion. For
example, in sodium chloride there are no molecules formed. The ionic lattice for sodium
chloride is shown in the following activity.
Sodium chloride ionic lattice
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Sodium chloride lattice
This extensively bonded ionic solid is relatively hard, though not as hard as the
covalent network solids, but is brittle. (Think of how easily salt (sodium chloride)
crystals can be crushed.)
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TOPIC 5. BONDING AND STRUCTURE
Rock salt cellar
Q34: What are caesium chloride (CsCl) crystals formed from?
a)
b)
c)
d)
Covalent molecules.
A covalent lattice.
An ionic lattice.
Ionic molecules.
..........................................
Now that we have looked at the three types of bonding in elements and compounds,
see the table below roughly compares their strengths. The table also gives a figure for
the weaker bonds between molecules, which will be studied in detail in the next topic.
Bond type
Strength/kJ mol-1
Metallic
80 - 600
Covalent
100 - 500
Ionic
100 - 450
Between covalent molecules
1 - 30
Strengths of bonds
You can see that metallic bonds have the greatest range of strengths. (This is sensible
if you think of hard metals like tungsten and chromium, and soft metals like sodium.)
The strengths of covalent and ionic bonds are similar, and much greater than the bonds
which hold covalent molecules together. (These are the intermolecular forces, which
are studied in the next topic.)
Key point
Ionic bonding is the electrostatic force of attraction between positively and
negatively charged ions.
An ionic structure consists of a giant lattice of oppositely charged ions.
© H ERIOT-WATT U NIVERSITY
TOPIC 5. BONDING AND STRUCTURE
5.6
Summary
Summary
•
Metallic bonding is the electrostatic force of attraction between positively
charged ions and delocalised outer electrons.
•
A metallic structure consists of a giant lattice of positively charged ions and
delocalised outer electrons.
•
Atoms in a covalent bond are held together by electrostatic forces of
attraction between positively charged nuclei and negatively charged shared
electrons.
•
A covalent molecular structure consists of discrete molecules held together
by intermolecular forces.
•
A covalent network structure consists of a giant lattice of covalently bonded
atoms.
•
Ionic bonding is the electrostatic force of attraction between positively and
negatively charged ions.
•
An ionic structure consists of a giant lattice of oppositely charged ions.
•
Elements can be categorised into four classes according to their bonding
and structure:
1. Metallic
2. Covalent molecular
3. Covalent network
4. Monatomic
5.7
•
Resources
Higher Chemistry for CfE: J Anderson, E Allan and J Harris, Hodder Gibson,
ISBN 978-1444167528
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TOPIC 5. BONDING AND STRUCTURE
5.8
End of topic test
End of Topic 5 test
Go online
Q35: In which of the following compounds do both ions have the same electron
arrangement as argon?
a)
b)
c)
d)
Calcium sulfide
Magnesium oxide
Sodium sulfide
Calcium bromide
..........................................
Q36: Which element is a solid at room temperature and consists of discrete molecules?
a)
b)
c)
d)
Sulfur
Neon
Silicon
Boron
..........................................
Q37: Graphite, a form of carbon, conducts electricity because it has:
a)
b)
c)
d)
pure covalent bonding.
delocalised electrons.
metallic bonding.
van der Waals' bonding.
..........................................
Q38: Which of the following can be applied to lithium but not to carbon?
a)
b)
c)
d)
e)
f)
Covalent
Metallic
Made up of discrete molecules.
Made up of diatomic molecules.
Gas
Solid
..........................................
Q39: Which of the following can be applied to both fluorine and phosphorus?
a) Covalent
b) Metallic
c) Made up of discrete molecules.
d) Made up of diatomic molecules.
e) Gas
f) Solid
..........................................
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109
Q40: Identify the element which exists as a covalent network solid.
a)
b)
c)
d)
e)
f)
Boron
Chlorine
Nitrogen
Phosphorus
Sodium
Sulfur
..........................................
Q41: Identify the two elements which exist as discrete covalent molecular solids.
a) Boron
b) Chlorine
c) Nitrogen
d) Phosphorus
e) Sodium
f) Sulfur
..........................................
Q42: Identify the two elements which react to form a compound with the most ionic
character.
a) Boron
b) Chlorine
c) Nitrogen
d) Phosphorus
e) Sodium
f) Sulfur
..........................................
Q43: Which of the elements is most likely to have a covalent network structure?
Element
Melting Point
(K)
Boiling Point
(K)
Density (g
cm-3 )
Conduction
when solids
A
317
553
1.82
No
B
933
2740
2.7
Yes
C
1683
2628
2.32
No
D
387
457
4.93
No
a)
b)
c)
d)
A
B
C
D
..........................................
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Q44: Which type of structure is found in a substance melting at 771 K which conducts
electricity when molten, but not when solid?
a)
b)
c)
d)
Ionic
Covalent (network structure)
Metallic
Covalent (discrete molecules)
..........................................
Q45: Identify the covalent network substance.
a)
b)
c)
d)
e)
f)
NH4 Cl (s)
CH3 OH (l)
C6 H14 (l)
SO2 (g)
Na2 CO3 (s)
SiO2 (s)
..........................................
Q46: Identify the two substances which are ionic.
a) NH4 Cl (s)
b) CH3 OH (l)
c) C6 H14 (l)
d) SO2 (g)
e) Na2 CO3 (s)
f) SiO2 (s)
..........................................
Q47: Diamond and graphite are forms of carbon with very different properties. Graphite
can mark paper, is a lubricant and is a conductor of electricity. Diamond has none of
these properties.
Draw a diagram to show the structure of diamond.
..........................................
Q48: Diamond and graphite are forms of carbon with very different properties. Graphite
can mark paper, is a lubricant and is a conductor of electricity. Diamond has none of
these properties.
Why is graphite an effective lubricant?
..........................................
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TOPIC 5. BONDING AND STRUCTURE
Q49: Boron nitride can form a similar structure to graphite. The boron and nitrogen
atoms alternate throughout the structure as shown.
Why is this substance a non-conductor, while graphite conducts electricity?
..........................................
Q50: Suggest why the bonds between the layers in boron nitride are stronger than the
bonds between the layers in graphite.
..........................................
Q51: Which of the following structures is never found in compounds?
a)
b)
c)
d)
Ionic
Monatomic
Covalent Molecular
Covalent Network
..........................................
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Topic 6
Periodic Table trends
Contents
6.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . .
6.2 Covalent radius . . . . . . . . . . . . . . . . . . . . . . . .
6.2.1 The trends in covalent radius . . . . . . . . . . . .
6.3 Ionisation energies . . . . . . . . . . . . . . . . . . . . . .
6.3.1 Explanation of the trends in first ionisation energy
6.4 Electronegativity . . . . . . . . . . . . . . . . . . . . . . .
6.5 Summary of trends in the Periodic Table . . . . . . . . . .
6.6 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . .
6.7 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . .
6.8 End of topic test . . . . . . . . . . . . . . . . . . . . . . . .
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115
116
117
119
122
125
128
129
129
130
Prerequisite knowledge
Before you begin this topic, you should already know that:
•
atoms contain protons, neutrons and electrons each with a specific charge, mass
and position within the atom - the number of protons defines an element and is
known as the atomic number (National 4, Unit 1);
•
or have knowledge of: sub-atomic particles, their charge, mass and position within
the atom, the structure of the Periodic Table, groups, periods and atomic number
(National 5, Unit 1);
•
all matter is made of atoms - when a substance contains only one kind of atom it
is known as an element (National 4, Unit 1);
•
elements are arranged in the Periodic Table in order of increasing atomic number
- elements with similar chemical properties are grouped together (National 4, Unit
1);
•
or be familiar with the seven diatomic elements (National 5, Unit 1);
•
covalent compounds form when non-metal atoms form covalent bonds by sharing
their outer electrons (National 4, Unit 1);
•
covalent molecular compounds have low melting and boiling points - as a result,
they can be found in any state at room temperature (National 4, Unit 1);
114
TOPIC 6. PERIODIC TABLE TRENDS
•
in a covalent bond, the shared pair of electrons is attracted to the nuclei of the two
bonded atoms (National 5, Unit 1);
•
more than one bond can be formed between atoms leading to double and triple
covalent bonds (National 5, Unit 1);
•
covalent substances can form either discrete molecular or giant network structures
(National 5, Unit 1);
•
diagrams show how outer electrons are shared to form the covalent bond(s) in a
molecule and the shape of simple two-element compounds (National 5, Unit 1);
•
when there is an imbalance in the number of positive protons and electrons the
particle is known as an ion (National 5, Unit 1).
Learning objectives
By the end of this topic, you should be able to:
•
state that covalent radius is a measure of the size of an atom;
•
state that trends in covalent radius across periods and down groups can be
explained in terms of the number of occupied shells, and the nuclear charge;
•
state that trends in ionisation energies across periods and down groups can be
explained in terms of the atomic size, nuclear charge and the screening effect due
to inner shell electrons;
•
explain that atoms of different elements have different attractions for bonding
electrons;
•
state that electronegativity is a measure of the attraction an atom involved in a
bond has for the electrons of the bond;
•
state that electronegativity values increase across a period and decrease down a
group;
•
state that electronegativity trends can be rationalised in terms of nuclear charge,
covalent radius and the presence of 'screening' inner electrons.
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TOPIC 6. PERIODIC TABLE TRENDS
6.1
115
Prior knowledge
Test your prior knowledge
Q1: Complete the following table:
Particle
Charge
Mass
Proton
Neutron
Electron
..........................................
Q2: Elements in the same group in the Periodic Table have the same:
a)
b)
c)
d)
number of occupied energy shells.
density.
number of outer electrons.
number of protons.
..........................................
Q3: Elements in the same period in the Periodic Table have the same:
a)
b)
c)
d)
number of occupied energy shells.
density.
number of outer electrons.
number of protons.
..........................................
Q4: What determines the number of protons in an atom?
a)
b)
c)
d)
Atomic number
Mass number
Name of element
Charge of atom
..........................................
Q5: Covalent bonding involves:
a)
b)
c)
d)
a shared pair of electrons.
transfer of electrons.
delocalised electrons.
gaining electrons.
..........................................
..........................................
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116
TOPIC 6. PERIODIC TABLE TRENDS
6.2
Covalent radius
The size of individual atoms is difficult to measure since the size is determined by the
space taken up by the constantly moving electrons. However, the distance between the
nuclei of atoms in the solid state can be measured by a technique called X-ray diffraction.
Consequently, the size of an atom is usually described in terms of its covalent radius.
This is defined as half the distance between the nuclei of two bonded atoms of the
element (see Figure 6.1).
Figure 6.1: Definition of covalent radius
..........................................
In Figure 6.1, the distance between the nuclei is shown as 2r and so the covalent radius
in each case is r.
Covalent radius - relative sizes
The covalent radius is another periodic property.
Go online
Figure 6.2: The covalent radius
..........................................
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TOPIC 6. PERIODIC TABLE TRENDS
117
Q6: What happens to the size of the atoms on crossing a Period from left to right?
a) An increase
b) A decrease
c) Nothing
..........................................
Q7: What happens to the size of the atoms on descending a group?
a) An increase
b) A decrease
c) Nothing
..........................................
Q8: Can you suggest why are there no covalent radii quoted for the noble gases
(Group 0)?
..........................................
..........................................
6.2.1
The trends in covalent radius
The size of an atom is determined by the amount of space taken up by the electrons
and so must be connected to the electron arrangement of the atom. The electron
arrangement is itself a periodic property.
Use the following activity to explain the trends in covalent radius.
The trends in covalent radius
Figure 6.3: The trends in covalent radius
Go online
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TOPIC 6. PERIODIC TABLE TRENDS
..........................................
Q9: Which of the following provides the best reason for the increase in covalent radius
on going down a group?
a)
b)
c)
d)
The number of protons increases.
The number of electrons increases.
The number of electron shells increases.
The number of neutrons increases.
..........................................
Q10: Which of the following provides the best reason for the decrease in covalent radius
on going from left to right across a Period?
a)
b)
c)
d)
The number of electrons increases.
The number of electron shells increases.
The number of neutrons increases.
The number of protons increases.
..........................................
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TOPIC 6. PERIODIC TABLE TRENDS
..........................................
Key point
The covalent radius decreases across a Period because the increase in nuclear
charge attracts the electrons more strongly.
The covalent radius increases on going down a group as the number of occupied
electron shells increases.
6.3
Ionisation energies
The electrons orbiting the nucleus of an atom are held by the electrostatic attraction
between the negative electrons and the positive nucleus. Some atoms lose electrons
relatively easily, whereas some lose electrons only with great difficulty. No atoms simply
give away electrons; to remove an electron always requires energy to overcome the force
of attraction between the electron and the nucleus. The electron lost always comes from
the outer shell of electrons.
The energy required to remove an electron from an atom can be measured. It is normal
to quote values not for an individual atom but for one mole of atoms.
Moles
At National 5 level, the mole was introduced as the gram formula mass of a substance
(the formula mass expressed in grams). For an element, one mole is normally the gram
formula mass, e.g.
•
12 g of carbon is 1 mole
•
4 g of helium is 1 mole
•
40 g of calcium is 1 mole
These three quantities are different masses but they have one very important feature
in common: they all contain the same number of atoms. This arises because calcium
atoms are ten times as heavy as helium atoms while carbon atoms are three times as
heavy as helium atoms.
This number is also known as one mole which refers to a particular number of particles
(this will be investigated more fully in Unit 3). In fact, one mole of a substance contains
a huge number of particles.
Ionisation energy
The first ionisation energy (IE) is defined as the energy required to remove one mole
of electrons from one mole of gaseous atoms (one electron from each atom). The units
used are kilojoules per mole (kJ mol-1 ).
Consider the element carbon. The value quoted for the 1 st ionisation energy of carbon
is 1090 kJ mol-1 . In other words, 1090 kJ of energy is required to remove one electron
from each atom in one mole (12 g) of gaseous carbon.
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TOPIC 6. PERIODIC TABLE TRENDS
This can be represented in the following way:
•
the first ionisation energy for an element E refers to the reaction
E(g) → E+ (g) + e- ;
•
the second ionisation energy refers to the reaction
E+ (g) → E2+ (g) + e- .
Figure 6.4: First ionisation energy of carbon
..........................................
The C+ ion is smaller than the C atom because the remaining five electrons in the ion
still feel the full attractive force from six protons and so they are more tightly held. As a
result, it is even more difficult to remove the next electron.
Figure 6.5: Second ionisation energy of carbon
..........................................
The removal of one mole of electrons from one mole of gaseous C + ions is known as
the second ionisation energy of carbon. This has a value of 2360 kJ mole -1 . The figure
below shows equations representing the first four ionisation energies of carbon.
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TOPIC 6. PERIODIC TABLE TRENDS
121
Figure 6.6: Successive ionisation energies
..........................................
Ionisation energies
Q11: Using a data booklet, plot a graph of first ionisation energy against atomic number
for the first 20 elements.
..........................................
Q12: Is the first ionisation energy a Periodic property?
..........................................
Q13: From the shape of the graph, explain your answer to the previous question.
..........................................
Q14: Name the group of elements which appears at the peaks in the graph.
..........................................
Q15: Name the group of elements which have the lowest ionisation energies.
..........................................
Q16: What is the general trend in first ionisation energy, going across a Period from left
to right?
a)
b)
c)
d)
A steady increase.
A steady decrease.
An increase followed by a decrease.
A decrease followed by an increase.
..........................................
Q17: What is the general trend in first ionisation energy, going down a group?
a)
b)
c)
d)
A steady increase.
A steady decrease.
An increase followed by a decrease.
A decrease followed by an increase.
..........................................
Q18: Confirm your answer to the previous question by listing the first ionisation energies
of the alkali metals in order of atomic number.
..........................................
..........................................
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TOPIC 6. PERIODIC TABLE TRENDS
6.3.1
Explanation of the trends in first ionisation energy
Going down a group
In general, the first ionisation energy decreases on going down a group. For this
exercise, we will concentrate on Group 1 since there is only one outer electron.
The higher the ionisation energy, the more difficult it is to remove the electron, i.e. the
stronger are the attractive forces between the electron and the nucleus. The strength of
this force of attraction will depend on:
1. the size of the nuclear charge;
2. the distance between the electron and the nucleus;
3. the number of other electrons between the electron and the nucleus (i.e. the
number of inner-shell electrons). The inner electrons cause a screening effect
which prevents the outer electron from feeling the full effect of the nuclear charge.
First ionisation energy - Going down a group
Consider the electron arrangements of lithium, sodium and potassium (see Figure 6.7).
Figure 6.7: Electron arrangements of alkali metals
..........................................
Now discuss the following points with a partner, a group or even your tutor.
•
What effect will an increase in nuclear charge have on the ionisation energy?
•
What effect will an increase in covalent radius have on the ionisation energy?
•
What will happen to the number of inner shell electrons on going down a group?
•
What effect will this have on the ionisation energy?
After discussion, you should be able to answer the following question.
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TOPIC 6. PERIODIC TABLE TRENDS
123
Q19: Explain fully why the first ionisation energy decreases on going down Group 1.
You should mention nuclear charge, covalent radius and screening effect.
(There would be about 3 marks allocated to this type of 'explain' question.)
..........................................
..........................................
Going across a Period
The first ionisation energy increases on going from left to right across a Period.
The same factors, as described previously, will be involved in explaining this trend.
Concentrate on Period 2 for this part.
First ionisation energy - Going across a Period
Discuss the same points with a partner as before.
Go online
•
What effect will an increase in nuclear charge have on the ionisation energy?
•
What effect will an increase in covalent radius have on the ionisation energy?
•
What will happen to the number of inner shell electrons on going down a group?
•
What effect will this have on the ionisation energy?
Now answer the following questions.
Q20: Explain fully why the first ionisation energy increases on going across Period 2
from left to right.
(There would be about three marks allocated to this type of 'explain' question.)
..........................................
Q21: What happens to the nuclear charge?
a) It decreases.
b) It increases.
c) It stays the same.
..........................................
Q22: What effect will this have on the first ionisation energy?
a) It increases.
b) It decreases.
c) It stays the same
..........................................
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TOPIC 6. PERIODIC TABLE TRENDS
Q23: What happens to the number of inner shell electrons?
a) It decreases.
b) It increases.
c) It stays the same.
..........................................
Q24: What happens to the screening effect?
a) It decreases.
b) It increases.
c) It stays the same.
..........................................
Q25: As a result, what happens to the size of the atoms?
a) It decreases.
b) It increases.
c) It stays the same.
..........................................
•
Use your data book to calculate the energy required to carry out the following
reactions:
Al(g) → Al+ (g) + e- Mg(g) → Mg2+ (g) + 2e- -
•
Why is there such a large increase in the energy required to remove a fourth
electron from aluminium compared to removing the first, second or third electrons?
..........................................
Key point
The first ionisation energy is the energy required to remove one mole of electrons
from one mole of gaseous atoms. The second and subsequent ionisation
energies refer to the energies required to remove further moles of electrons.
First ionisation energies increase across a Period and decrease down a group.
This can be explained in terms of atomic size, nuclear charge and the screening
effect due to inner shell electrons.
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TOPIC 6. PERIODIC TABLE TRENDS
6.4
125
Electronegativity
The first ionisation energy involves the removal of electrons from gaseous atoms and so
is a measure of how strongly an isolated atom holds on to its outermost electrons. In the
world around us, isolated atoms are very rare and the vast majority of atoms are found
bonded to one or more other atoms. How and why atoms bond together is the basis of
chemistry and electrons are the fundamental particles involved in bonding.
In all theories of bonding, different types of bond arise because atoms of different
elements have different attractions for electrons. Atoms which tend to attract the
electrons within a bond are said to be electronegative.
An atom with a high electronegativity will tend to attract bonded electrons towards
it whereas an atom with a low electronegativity will have a very weak attraction for
electrons. There will be a 'tug of war' between the different atoms for the electrons.
Electronegativity
In Figure 6.8, the darker shading shows the electrons being pulled towards the more
electronegative atom.
Figure 6.8: Electronegativity
..........................................
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126
TOPIC 6. PERIODIC TABLE TRENDS
Q26: If X is hydrogen, Y could be:
a) Beryllium
b) Bromine
c) Boron
..........................................
Q27: The bond X-Y could be:
a) C-Cl
b) H-S
c) C-S
..........................................
Q28: If X is in Group 5, Y could be in:
a) Group 3
b) Group 6
c) Group 7
..........................................
..........................................
There are several different methods for estimating electronegativity values. The one
most commonly used is that devised by double Nobel Prize winner, Linus Pauling. He
assigned a value to each of the elements most commonly found in bonds. Lithium
was assigned a value of 1.0 while fluorine was assigned the highest value of 4.0. The
electronegativity values produced by Pauling are quoted in the SQA data booklet (page
11). Study these values carefully and look for any patterns.
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TOPIC 6. PERIODIC TABLE TRENDS
127
Q29: What is the Group number of the elements which have the lowest
electronegativities?
..........................................
Q30: What is the Group number of the elements with the highest electronegativities?
..........................................
Q31: What is the name of the Group for which no values are quoted?
..........................................
Q32: Can you suggest a reason why no values are quoted for these elements?
..........................................
In general, the electronegativity increases from left to right across a Period. This is
because as you move from left to right, there are more protons and so the nuclear charge
is increased. As you move down a Group, atomic size is increasing so outer electrons
are further from the positively charged nucleus. As a result, the electronegativity
decreases.
Q33: Which of the following statements describes the trends in electronegativity values
in the Periodic Table?
a) Electronegativity values increase on going from left to right
increase on going down a Group.
b) Electronegativity values decrease on going from left to right
decrease on going down a Group.
c) Electronegativity values increase on going from left to right
decrease on going down a Group.
d) Electronegativity values decrease on going from left to right
increase on going down a Group.
across a Period and
across a Period and
across a Period and
across a Period and
..........................................
Key point
Electronegativity is a measure of the attraction an atom in a bond has for the
electrons of the bond. Electronegativity values increase across a Period and
decrease down a Group.
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TOPIC 6. PERIODIC TABLE TRENDS
6.5
Summary of trends in the Periodic Table
Summary: Periodic Table trends
Go online
Q34: Complete the table by selecting the appropriate arrow symbol to indicate the
increasing size of each property as you move across a Period and up/down Group 1
and 7.
..........................................
© H ERIOT-WATT U NIVERSITY
TOPIC 6. PERIODIC TABLE TRENDS
6.6
Summary
Summary
6.7
•
•
The covalent radius decreases across a Period because the increase in
nuclear charge attracts the electrons more strongly.
•
The covalent radius increases on going down a Group as the number of
occupied electron shells increases.
•
The first ionisation energy is the energy required to remove one mole of
electrons from one mole of gaseous atoms.
•
The second and subsequent ionisation energies refer to the energies
required to remove further moles of electrons.
•
First ionisation energies increase across a Period and decrease down a
Group.
•
This can be explained in terms of atomic size, nuclear charge and the
screening effect due to inner shell electrons.
•
Electronegativity is a measure of the attraction an atom in a bond has for
the electrons of the bond.
•
Electronegativity values increase across a Period and decrease down a
Group.
•
This can be explained in terms of atomic size, nuclear charge and the
screening effect due to inner shell electrons.
Resources
Higher Chemistry for CfE, J Anderson, E Allan and J Harris, Hodder Gibson,
ISBN 978-1444167528
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TOPIC 6. PERIODIC TABLE TRENDS
6.8
End of topic test
End of Topic 6 test
Go online
Q35: Menedeleev is famous for producing the Periodic Table on which the modern
version is based. Which of the following statements is true?
a) Mendeleev organised the elements in order of their atomic number.
b) Mendeleev swapped some elements round so that their atomic masses fitted the
pattern.
c) Mendeleev left gaps for elements which had not yet been discovered.
d) Mendeleev left gaps because some elements did not fit the pattern of reactivity.
..........................................
Q36: Which of the following statements about the Periodic Table are true?
1. There is a steady decrease in density on going down Group 1.
2. There is an increase in electronegativity on going down Group 7.
3. There is a decrease in first ionisation energy on going down Group 0.
4. There is a decrease and then an increase in boiling point on crossing a Period from
left to right.
5. There is a steady increase in melting point across a Period from left to right.
6. There is a steady decrease in atomic size across a Period from left to right.
..........................................
Q37: Which property of the Group 1 elements could be represented by the following
graph?
Arbitrary
Scale
Li
a)
b)
c)
d)
Na
K
Rb
Electronegativity
Atomic size
Melting point
First ionisation energy
..........................................
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131
Q38: Some information about four atoms A, B, C and D is presented below.
D
C
Covalent
Radius
B
A
Number of occupied electron shells
A
3
B
4
C
3
D
4
Which atom will have the largest nuclear charge?
a)
b)
c)
d)
A
B
C
D
..........................................
Q39: Calcium has a larger covalent radius than magnesium because calcium has:
a)
b)
c)
d)
a smaller first ionisation energy?
a larger nucleus?
a larger nuclear charge?
more occupied electron shells?
..........................................
Q40: The difference in atomic size between sodium and chlorine is mainly due to the
number of:
a)
b)
c)
d)
electron shells?
neutrons?
protons?
electrons?
..........................................
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Q41: Which of the following equations represents the first ionisation energy of
magnesium?
a) Mg(s) → Mg+ (g) + eb) Mg(g) → Mg2+ (g) + 2ec) Mg(g) → Mg+ (g) + 2ed) 12 Mg(g) → 12 Mg2+ (g) + e−
..........................................
Q42: Which of the following equations represents the first ionisation energy of fluorine?
a) F- (g) → F(g) + eb) 12 F2 (g) → F + (g) + e−
c) F(g) + e- → F- (g)
d) F(g) → F+ (g) + e..........................................
Q43: Identify the statements which correctly describe a trend in the Periodic Table.
1. The covalent radius increases from Li to F.
2. The boiling point increases from Li to F.
3. The first ionisation energy decreases from Na to Cl.
4. The first ionisation energy decreases from Li to Cs.
5. Electronegativity decreases from Li to Cs.
6. The number of electrons in the outer shell increases from Li to Cs.
..........................................
Q44: The change in first ionisation energy from Li to F is mainly due to:
a)
b)
c)
d)
increasing number of electron shells?
increased screening effect?
increasing number of electrons?
increasing nuclear charge?
..........................................
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133
Q45: The bar graph shows the variation in the first ionisation energy with atomic number
for sixteen consecutive elements in the Periodic Table. The element at which the bar
graph starts is not specified.
First
Ionisation
Energy
(kJ mol-1)
Z
Atomic Number
Element Z is identified in the bar graph. In which Group of the table is element Z?
a)
b)
c)
d)
1
3
5
6
..........................................
Q46: Which of the following graphs shows the variation in first ionisation energy on
going from left to right across period 2?
a)
Atomic Number
b)
Atomic Number
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TOPIC 6. PERIODIC TABLE TRENDS
c)
Atomic Number
d)
Atomic Number
..........................................
Q47: Which of the following will not affect the first ionisation energy of an element?
a)
b)
c)
d)
Screening effect
Atomic size
Atomic number
Atomic mass
..........................................
Q48:
bonding electrons.
is a measure of the ability of an atom in a bond to attract the
..........................................
Q49:
has the highest attraction for bonding electrons.
..........................................
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TOPIC 6. PERIODIC TABLE TRENDS
Q50: Atoms of different elements have different attractions for the electrons in a bond.
This property shows Periodic variation.
Which of the following shows the correct trends within the Periodic Table?
a)
b)
c)
d)
..........................................
Q51: Explain fully why the first ionisation energy increases on going across Period 2
from left to right. (3 marks)
..........................................
Q52: Explain fully why the first ionisation energy decreases on going down Group 1.
You should mention nuclear charge, covalent radius and screening effect. (3 marks)
..........................................
Q53: Why are no electronegativity values are quoted for the Group 0 elements?
a)
b)
c)
d)
They have a value of zero.
They generally do not form bonds with other elements.
They are unreactive non-metals.
Their electronegativities are too high for the scale.
..........................................
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Topic 7
Bonding continuum and polar
covalent bonding
Contents
7.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
139
7.2 Polar covalent bonds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
7.3 Predicting bonding type using electronegativity . . . . . . . . . . . . . . . . . .
140
144
7.4 Polar molecules . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
7.4.1 Polar molecules and molecular polarity . . . . . . . . . . . . . . . . . .
146
147
7.5 The bonding continuum . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
7.6 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
150
152
7.7 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
7.8 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
152
153
Prerequisite knowledge
Before you begin this topic, you should already know that:
•
covalent compounds form when non-metal atoms form covalent bonds by sharing
their outer electrons (National 4, Unit 1);
•
covalent molecular compounds have low melting and boiling points - as a result,
they can be found in any state at room temperature (National 4, Unit 1);
•
in a covalent bond, the shared pair of electrons is attracted to the nuclei of the two
bonded atoms (National 5, Unit 1);
•
covalent substances can form either discrete molecular or giant network structures
(National 5, Unit 1);
•
diagrams show how outer electrons are shared to form the covalent bond(s) in a
molecule and the shape of simple two-element compounds (National 5, Unit 1);
•
the first 20 elements in the Periodic Table can be categorised according to bonding
and structure:
◦
metallic (Li, Be, Na, Mg, Al, K, Ca);
◦
covalent molecular (H2 , N2 , O2 , F2 , Cl2 , P4 , S8 and fullerenes (eg C 60 ));
138
TOPIC 7. BONDING CONTINUUM AND POLAR COVALENT BONDING
◦
covalent network (B, C (diamond, graphite), Si) monatomic (noble gases);
•
electronegativity is a measure of the attraction an atom in a bond has for the
electrons of the bond (Higher, Unit 1);
•
electronegativity values increase across a Period and decrease down a Group
(Higher, Unit 1);
•
trends in electronegativity can be explained in terms of atomic size, nuclear charge
and the screening effect due to inner shell electrons (Higher, Unit 1).
Learning objectives
By the end of this topic, you should be able to:
•
state that, in a covalent bond, atoms share pairs of electrons;
•
explain that a covalent bond is a result of two positive nuclei being held together
by their common attraction for the shared pair of electrons;
•
explain that polar covalent bonds are formed when the attraction of the atoms for
the pair of bonding electrons is different;
•
explain that delta positive and delta negative notation can be used to indicate the
partial charges on atoms, which give rise to a dipole;
•
state that pure covalent bonding and ionic bonding can be considered as being at
opposite ends of a bonding continuum with polar covalent bonding lying between
these two extremes;
•
state that the larger the difference in electronegativities between bonded atoms,
the more polar the bond will be;
•
state that, if the difference is large, then the movement of bonding electrons from
the element of lower electronegativity to the element of higher electronegativity is
complete, resulting in the formation of ions;
•
state that compounds formed between metals and non-metals are often, but not
always ionic.
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TOPIC 7. BONDING CONTINUUM AND POLAR COVALENT BONDING
7.1
139
Prior knowledge
Test your prior knowledge
Q1: Covalent bonding involves:
a)
b)
c)
d)
transfer of electrons.
a shared pair of electrons.
delocalised electrons.
gaining electrons.
..........................................
Q2: Why are no electronegativity values quoted for the Group 0 elements?
a)
b)
c)
d)
They have a value of zero.
They are unreactive non-metals.
They generally do not form bonds with other elements.
Their electronegativities are too high for the scale.
..........................................
Q3: Which of the following elements has the greatest attraction for the shared pair of
electrons in a bond?
a)
b)
c)
d)
Fluorine
Carbon
Hydrogen
Chlorine
..........................................
Q4: Which of the following elements has the least attraction for the shared pair of
electrons in a bond?
a)
b)
c)
d)
Fluorine
Carbon
Hydrogen
Chlorine
..........................................
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TOPIC 7. BONDING CONTINUUM AND POLAR COVALENT BONDING
7.2
Polar covalent bonds
Previously, you saw how the electronegativity of atoms is a measure of their attraction
for electrons in a bond. Only when the two atoms bonded by a covalent bond are the
same (e.g. Cl2 in chlorine gas, C-C in diamond) will the electrons be exactly shared
equally. In other cases there will be an unequal distribution of the electrons and the
atom with the higher electronegativity will have a greater share of the electrons.
Since electrons carry a negative charge, an unequal distribution will result in the bond
having a partial negative charge (called 'delta negative' and drawn δ-) where there is an
excess of electrons around the most electronegative atom, and a partial positive charge
(δ+) where there is a deficiency. This is not to be confused with a full charge as found
on ions.
Figure 7.1: Polar bond
..........................................
In the case of HCl (Figure 7.1), chlorine, with electronegativity 3.0, becomes negative
and hydrogen, at 2.2, positive. The δ partial charge is about 0.17 of a full charge.
This type of bond is called a polar covalent bond (sometimes abbreviated to a polar
bond). The greater the difference between the electronegativities of the atoms, the
greater the distortion of the electrons in the bond, and the greater the charge
distribution. This effect is shown in the following activity.
Electronegativity
Go online
In Figure 7.2, the darker shading shows the electrons being pulled towards the more
electronegative atom.
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TOPIC 7. BONDING CONTINUUM AND POLAR COVALENT BONDING
Figure 7.2: Electronegativity and charge
..........................................
Q5: How much ionic character could this bond have?
a) 0%
b) 50%
c) 100%
..........................................
Q6: This bond has almost zero ionic character. It could be:
a) Li-F
b) Mg-O
c) C-S
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TOPIC 7. BONDING CONTINUUM AND POLAR COVALENT BONDING
..........................................
Q7:
This bond could be:
a) X is Na and Y is O
b) X is S and Y is K
c) X is N and Y is Cl
..........................................
Use the table of electronegativities, where necessary, (SQA data booklet, page 11) to
answer the following questions.
Q8:
a)
b)
c)
d)
Which molecule contains polar bonds?
Hydrogen (H2 )
Chlorine (Cl2 )
Hydrogen chloride (HCl)
Nitrogen (N2 )
..........................................
Q9:
a)
b)
c)
d)
Which molecule contains the most polar bonds?
Hydrogen fluoride (HF)
Hydrogen chloride (HCl)
Hydrogen bromide (HBr)
Hydrogen iodide (HI)
..........................................
Q10: Which molecule has the bonds that are most polar?
a)
b)
c)
d)
Hydrogen sulfide (H2 S)
Phosphine (PH 3 )
Methane (CH4 )
Ammonia (NH 3 )
..........................................
can be used to show the direction of the dipole, the arrow
The symbol
pointing to the negative side.
..........................................
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TOPIC 7. BONDING CONTINUUM AND POLAR COVALENT BONDING
143
Direction of charge
Use the electronegativity values (page 11 of the data booklet) to help show the direction
of partial charge (or no charge) in some bonds.
Q11: Complete the table by selecting the correct sign for the bond shown.
..........................................
Key point
Polar covalent bonds occur when the atoms of the bond attract the bonding
electrons unequally causing the atoms to have partial positive and negative
charges.
The polarity of a covalent bond depends on the difference in electronegativity
between the bonded atoms, the most electronegative becoming more negative.
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TOPIC 7. BONDING CONTINUUM AND POLAR COVALENT BONDING
7.3
Predicting bonding type using electronegativity
One of the main factors determining the type of bond formed between two elements in
a compound is the difference in electronegativity. As previously discussed,
electronegativity is a Periodic property, with increasing electronegativity as you move
across a Period from left to right, and decreasing electronegativity as you move down a
Group.
The relationship between electronegativity and bond type is summarised below.
Figure 7.3: Relationship between electronegativity and bond type
..........................................
Electronegativity difference is only one predictor of bond type. For a more complete
description, the properties of the substances (as described in another topic) need to be
considered.
Most covalent bonded compounds generally exist as discrete molecular structures.
However, two important covalent network compounds are silicon dioxide (SiO 2 , ) and
silicon carbide (SiC, carborundum), which has a similar structure to diamond. Both of
these are used as abrasives on account of their hardness.
Use the electronegativity values in the SQA data booklet to help you answer the
following questions.
© H ERIOT-WATT U NIVERSITY
TOPIC 7. BONDING CONTINUUM AND POLAR COVALENT BONDING
Q12: Hydrogen, bromine and potassium have electronegativities of 2.2, 2.8 and 0.8
respectively. What type of bonding would you expect:
•
between two bromine atoms in a Br 2 molecule?
•
between hydrogen and bromine in HBr?
•
between potassium and bromine in KBr?
..........................................
Q13: What type of bonding would you expect between the Group 1 metal caesium and
the Group 6 element sulfur?
a)
b)
c)
d)
Pure covalent
Polar covalent
Ionic
Metallic
..........................................
Q14: What type of bonding would you expect between phosphorus and hydrogen in
phosphine (PH 3 ) ?
a)
b)
c)
d)
Pure covalent
Polar covalent
Ionic
Metallic
..........................................
Q15: Which of the elements below will bond with oxygen, to produce a polar bond with
a partial positive charge on the oxygen atom?
a)
b)
c)
d)
Carbon
Fluorine
Hydrogen
Lithium
..........................................
Key point
The type of bonding between two atoms depends mainly on the difference in
electronegativity between the atoms.
When the difference is zero, the bond will be covalent. With a small difference, a
polar covalent bond is likely. When the difference is large, fully charged ions are
produced and ionic bonding will be predicted.
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TOPIC 7. BONDING CONTINUUM AND POLAR COVALENT BONDING
7.4
Polar molecules
When two atoms of different electronegativity are bonded together by sharing electrons,
one atom attracts the electrons more than the other and a polar bond results. In simple
molecules like hydrogen chloride and iodine chloride this polar bond has a permanent
dipole.
Polar molecules are attracted to one another by forces called permanent
dipole-permanent dipole interactions (see Figure 7.4) as well as the London
dispersion attractions caused by the movement of electrons observed in the last
section.
Permanent dipole interactions
Figure 7.4: Permanent dipoles in iodine chloride
Go online
..........................................
How do permanent dipole-permanent dipole interactions compare to London
dispersion forces of attraction in terms of strength?
A comparison of strengths can be made by looking at the table below.
Figure 7.5: Boiling points of halogen containing compounds
A
B
C
Cl - Cl
I - Cl
Br - Br
-35o C
97o C
59o C
..........................................
Any comparison of boiling points has to be between molecules of similar size and
shape so that the London dispersion forces are similar.
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TOPIC 7. BONDING CONTINUUM AND POLAR COVALENT BONDING
Q16: Which molecule has permanent dipole-permanent dipole interactions?
a) A
b) B
c) C
..........................................
Q17: Which other molecule would have almost the same London dispersion forces?
a) A
b) B
c) C
..........................................
Q18: Which of the molecules has the strongest intermolecular forces?
a) A
b) B
c) C
..........................................
..........................................
Key point
Permanent dipole-permanent dipole interactions act in addition to London
dispersion electrostatic attractions between polar molecules and are stronger
than these attractions for molecules of equivalent size.
7.4.1
Polar molecules and molecular polarity
Not all covalent molecules with polar bonds result in polar molecules. Those molecules
which are highly symmetrical prove to be non-polar (see Figure 7.6).
Figure 7.6: Non-polar molecules
..........................................
In each of the two molecules shown, although there is a difference in electronegativity
in the polar bonds, the charge is distributed around the central carbon atoms with the
positive and negative charges balancing out. The molecules have no overall dipole.
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148
TOPIC 7. BONDING CONTINUUM AND POLAR COVALENT BONDING
Notice that the central carbon on each of the molecules is shown with only one (δ + )
sign for clarity, even though each atom has sufficient charge to balance out the
negative charge.
In molecules with polar bonds which are not symmetrical (see Figure 7.7), the dipoles
cannot cancel each other out. If the molecule has a permanent slight positive charge at
one side and a negative charge at the other, then it is a polar molecule. Bond polarity
can thus be predicted from electronegativity differences and molecular polarity can be
predicted from electronegativity differences if we also take into account the shape of
the molecule.
can be used to show the direction of the dipole, the arrow
The symbol
pointing to the negative side. The molecules shown in the figure below are both polar
molecules.
Figure 7.7: Polarity of molecules
..........................................
Predicting molecular polarity
Go online
Electronegativity values are used to predict the polarity of bonds, and the shapes of
molecules are used to predict molecular polarity. Page 11 in the data booklet may be
helpful.
Q19: Complete the table by using the items below it. Dipole symbols may be repeated.
..........................................
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TOPIC 7. BONDING CONTINUUM AND POLAR COVALENT BONDING
149
Detecting polar molecules
An experiment is performed to show the attractions between polar molecules and
electrostatically charged rods.
Figure 7.8
Paraffin
Water
Charged
Rods
..........................................
Q20: Which liquid is affected by the charged rod?
a) Water
b) Paraffin
..........................................
Q21: Which liquid is polar?
a) Water
b) Paraffin
..........................................
Q22: Which of these statements is correct when a jet of polar molecules passes a
charged rod?
a)
b)
c)
d)
It is not affected by the rod.
It is attracted by the rod.
It is repelled by the rod.
It is sped up by the rod.
..........................................
Q23: Which of these statements is correct when a jet of non-polar molecules passes a
charged rod?
a)
b)
c)
d)
It is not affected by the rod.
It is attracted by the rod.
It is repelled by the rod.
It is sped up by the rod.
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TOPIC 7. BONDING CONTINUUM AND POLAR COVALENT BONDING
..........................................
..........................................
Key point
An electrostatically charged rod can be used to detect the presence of polar
molecules in a liquid. Polar molecules are attracted to both a negative and positive
rod.
7.5
The bonding continuum
Covalent bonding usually involves elements close to each other in the Periodic Table.
Ionic bonding generally involves metals from the left side and non-metals from the
opposite side of the table. Most chemical bonds are somewhere between the two
extremes and it is best to think of ionic and covalent bonding as being at opposite ends
of a bonding continuum with varying degrees of polar covalent bonding lying between
the extremes (see Figure 7.9).
Figure 7.9: Bonding spectrum
..........................................
In most cases, the bigger the difference in electronegativity between the atoms, the
more polar the bond and the greater the ionic character. However, other factors make a
contribution and care needs to be taken before jumping to conclusions. Identical atoms
share the electrons in a covalent bond equally. In the hydrogen molecule there is no
charge distribution. Non-identical atoms like hydrogen and chlorine attract the bonding
electrons unequally, since chlorine attracts electrons more strongly. A polar bond
results (see Figure 7.10).
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TOPIC 7. BONDING CONTINUUM AND POLAR COVALENT BONDING
Figure 7.10: Polar bond
..........................................
In some cases, the bond polarity results in a polar molecule. In certain symmetrical
molecules the polar nature of the bonds tends to cancel out (see Figure 7.11).
Figure 7.11: Polar and non-polar molecules
..........................................
In water, the molecule is asymmetrical, so the δ+ on the hydrogen atoms will add and
the molecule will have a resultant charge distribution; it will be a polar molecule.
The tetrachloromethane molecule is highly symmetrical, so the charge distribution on
the C-Cl bonds will cancel out and the molecule will be non-polar, despite having polar
bonds.
Q24: Will trichloromethane (CHCl 3 , chloroform) be a polar or non-polar molecule?
a) Polar
b) Non-polar
..........................................
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TOPIC 7. BONDING CONTINUUM AND POLAR COVALENT BONDING
7.6
Summary
Summary
7.7
•
•
Polar covalent bonds occur when the atoms of the bond attract the bonding
electrons unequally, causing the atoms to have partial positive and negative
charges.
•
The polarity of a covalent bond depends on the difference in
electronegativity between the bonded atoms, the most electronegative
becoming more negative.
•
There are polar covalent bonds between pure covalent and pure ionic
bonds.
•
The type of bonding between two atoms depends mainly on the difference
in electronegativity between the atoms:
◦
when the difference is zero, the bond will be covalent;
◦
with a small difference, a polar covalent bond is likely;
◦
when the difference is large, fully charged ions are produced and ionic
bonding will be predicted.
•
Permanent dipole-permanent dipole interactions act in addition to London
dispersion electrostatic attractions between polar molecules and are
stronger than these attractions for molecules of equivalent size.
•
Not all covalent molecules with polar bonds result in polar molecules.
•
Molecules which are highly symmetrical tend to be non-polar.
•
An electrostatically charged rod can be used to detect the presence of polar
molecules in a liquid. Polar molecules are attracted to both a negative and
positive rod.
•
There is a complete range of bond types leading to a bonding spectrum
mainly based on electronegativity.
Resources
Higher Chemistry for CfE, J Anderson, E Allan and J Harris, Hodder Gibson,
ISBN 978-1444167528
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TOPIC 7. BONDING CONTINUUM AND POLAR COVALENT BONDING
7.8
153
End of topic test
End of Topic 7 test
Q25: In which of the following compounds do both ions have the same electron
arrangement as argon?
a)
b)
c)
d)
Calcium bromide
Magnesium oxide
Sodium sulfide
Calcium sulfide
..........................................
Q26: Which element is a solid at room temperature and consists of discrete molecules?
a)
b)
c)
d)
Sulfur
Carbon
Silicon
Boron
..........................................
Q27: Graphite, a form of carbon, conducts electricity because it has:
a)
b)
c)
d)
metallic bonding.
pure covalent bonding.
delocalised electrons.
London dispersion forces.
..........................................
Q28: Which of the following is true of lithium but not of carbon?
a)
b)
c)
d)
e)
f)
Covalent
Metallic
Made up of discrete molecules.
Made up of diatomic molecules.
Gas
Solid
..........................................
Q29: Which of the following is true of both fluorine and phosphorus?
a) Covalent
b) Metallic
c) Made up of discrete molecules.
d) Made up of diatomic molecules.
e) Gas
f) Solid
..........................................
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TOPIC 7. BONDING CONTINUUM AND POLAR COVALENT BONDING
Q30: Identify the element which exists as a covalent network solid.
a)
b)
c)
d)
e)
f)
Boron
Chlorine
Nitrogen
Phosphorus
Sodium
Sulfur
..........................................
Q31: Identify the two elements which exist as discrete covalent molecular solids.
a) Boron
b) Chlorine
c) Nitrogen
d) Phosphorus
e) Sodium
f) Sulfur
..........................................
Q32: Identify the element which exists as a covalent network solid.
a) Boron
b) Chlorine
c) Nitrogen
d) Phosphorus
e) Sodium
f) Sulfur
..........................................
Q33: In which molecule will the chlorine atom carry a partial positive charge (δ + )?
a)
b)
c)
d)
Cl-F
Cl-Br
Cl-Cl
Cl-I
..........................................
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155
Q34: Which of the following elements is most likely to have a covalent network
structure?
Element
Melting point
(K)
Boiling point
(K)
Density ( g
cm-3 )
Conduction
when solid
A
317
553
1.82
No
B
933
2740
2.7
Yes
C
1683
2628
2.32
No
D
387
457
4.93
No
..........................................
Q35: Which of the following chlorides is likely to have least ionic character?
a)
b)
c)
d)
LiCl
CsCl
BeCl2
CaCl2
..........................................
Q36: Which type of structure is found in a substance melting at 771 K which conducts
electricity when molten, but not when solid?
a)
b)
c)
d)
Covalent (network structure)
Covalent (discrete molecules)
Ionic
Metallic
..........................................
Q37: Identify the covalent network substance.
a)
b)
c)
d)
e)
f)
NH4 Cl (s)
CH3 OH (l)
C6 H14 (l)
SO2 (g)
Na2 CO3 (s)
SiO2 (s)
..........................................
Q38: Identify the two substances which are ionic.
a) NH4 Cl (s)
b) CH3 OH (l)
c) C6 H14 (l)
d) SO2 (g)
e) Na2 CO3 (s)
f) SiO2 (s)
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..........................................
Q39: Identify the covalent network substance.
a)
b)
c)
d)
e)
f)
NH4 Cl (s)
CH3 OH (l)
C6 H14 (l)
SO2 (g)
Na2 CO3 (s)
SiO2 (s)
..........................................
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Topic 8
Intermolecular forces
Contents
8.1
8.2
8.3
8.4
8.5
Prior knowledge . . . . . . . . . . . . . . . . . . . . . .
Introduction . . . . . . . . . . . . . . . . . . . . . . . . .
London dispersion forces . . . . . . . . . . . . . . . . .
Hydrogen bonding . . . . . . . . . . . . . . . . . . . . .
Relating properties to bonding . . . . . . . . . . . . . .
8.5.1 Boiling point . . . . . . . . . . . . . . . . . . . .
8.5.2 Density . . . . . . . . . . . . . . . . . . . . . . .
8.6 Viscosity . . . . . . . . . . . . . . . . . . . . . . . . . .
8.7 Predicting solubilities from solute and solvent polarities
8.8 Summary . . . . . . . . . . . . . . . . . . . . . . . . . .
8.9 Resources . . . . . . . . . . . . . . . . . . . . . . . . .
8.10 End of topic test . . . . . . . . . . . . . . . . . . . . . .
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160
160
162
166
169
169
170
172
174
178
179
180
Prerequisite knowledge
Before you begin this topic, you should already know that:
•
the first 20 elements in the Periodic Table can be categorised according to bonding
and structure:
◦
metallic Li, Be, Na, Mg, Al, K, Ca;
◦
covalent molecular H2 , N2 , O2 , F2 , Cl2 , P4 , S8 and fullerenes (eg C 60 );
◦
covalent network B, C (diamond, graphite), S i ;
◦
monatomic (noble gases).
•
electronegativity is a measure of the attraction an atom in a bond has for the
electrons of the bond (Higher, Unit 1);
•
electronegativity values increase across a Period and decrease down a Group
(Higher, Unit 1);
•
this can be explained in terms of atomic size, nuclear charge and the screening
effect due to inner shell electrons (Higher, Unit 1);
•
in a covalent bond, atoms share pairs of electrons (National 4, Unit 1);
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TOPIC 8. INTERMOLECULAR FORCES
•
the covalent bond is a result of two positive nuclei being held together by their
common attraction for the shared pair of electrons (National 4, Unit 1);
•
polar covalent bonds are formed when the attraction of the atoms for the pair of
bonding electrons is different (Higher, Unit 1);
•
delta positive and delta negative notation can be used to indicate the partial
charges on atoms, which give rise to a dipole (Higher, Unit 1);
•
pure covalent bonding and ionic bonding can be considered as being at opposite
ends of a bonding continuum with polar covalent bonding lying between these two
extremes (Higher, Unit 1);
•
the larger the difference in electronegativities between bonded atoms, the more
polar the bond will be (Higher, Unit 1).
Learning objectives
By the end of this topic, you should be able to:
•
state that all molecular elements and compounds and monatomic elements
condense and freeze at sufficiently low temperatures - for this to occur, some
attractive forces must exist between the molecules or discrete atoms;
•
state that any ‘intermolecular’ forces acting between molecules are known as van
der Waals’ forces;
•
state that there are several different types of van der Waals’ forces such as
London dispersion forces and permanent dipole-permanent dipole interactions,
which include hydrogen bonding;
•
state that London dispersion forces are forces of attraction that can operate
between all atoms and molecules;
•
state that these forces are much weaker than all other types of bonding;
•
explain that London dispersion forces are formed as a result of electrostatic
attraction between temporary dipoles and induced dipoles, caused by movement
of electrons in atoms and molecules;
•
state that the strength of London dispersion forces is related to the number of
electrons within an atom or molecule;
•
state that a molecule is described as polar if it has a permanent dipole - the spatial
arrangement of polar covalent bonds can result in a molecule being polar;
•
state that permanent dipole-permanent dipole interactions are additional
electrostatic forces of attraction between polar molecules;
•
state that permanent dipole-permanent dipole interactions are stronger than
London dispersion forces for molecules with similar numbers of electrons;
•
state that bonds consisting of a hydrogen atom bonded to an atom of a strongly
electronegative element such fluorine, oxygen or nitrogen are highly polar;
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TOPIC 8. INTERMOLECULAR FORCES
•
state that hydrogen bonds are electrostatic forces of attraction between molecules
which contain these highly polar bonds;
•
state that a hydrogen bond is stronger than other forms of permanent dipolepermanent dipole interaction, but weaker than a covalent bond;
•
explain that melting points, boiling points and viscosity can all be rationalised in
terms of the nature and strength of the intermolecular forces which exist between
molecules;
•
explain that, by considering the polarity and number of electrons present in
molecules, it is possible to make qualitative predictions of the strength of the
intermolecular forces;
•
state that the melting and boiling points of polar substances are higher than
the melting and boiling points of non-polar substances with similar numbers of
electrons;
•
state that the anomalous boiling points of ammonia, water and hydrogen fluoride
are a result of hydrogen bonding;
•
explain that boiling points, melting points, viscosity and solubility/miscibility in
water are properties of substances which are affected by hydrogen bonding;
•
explain that hydrogen bonding between molecules in ice results in an expanded
structure which causes the density of ice to be less than that of water at low
temperatures.
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8.1
Prior knowledge
Test your prior knowledge
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Q1: Which of these is a measure of the ability of an atom in a bond to attract the
bonding electrons?
a)
b)
c)
d)
Ionisation Energy
Periodicity
Electronegativity
Polarity
..........................................
Q2: Which of the following elements has the greatest attraction for the shared pair of
electrons in a bond?
a)
b)
c)
d)
Fluorine
Nitrogen
Phosphorus
Lithium
..........................................
Q3:
a)
b)
c)
d)
As you move across the Periodic Table, the electronegativity:
stays the same.
decreases.
increases.
decreases then increases.
..........................................
8.2
Introduction
Simple molecules like nitrogen N 2 , methane CH4 and water H2 O, have their atoms held
together by covalent bonds within the molecule.
Bonds within a molecule are called intramolecular forces (intra: on the inside, as in
intramuscular, intravenous).
Nitrogen and methane are both gases at room temperature, but can be made liquid by
cooling to a very low temperature. Water is a liquid at room temperature and other
covalent molecules like candle wax are solid. Forces of attraction must exist between
all these covalent molecules or it would never be possible to have covalent molecular
solids at all. The molecules would fly apart and all discrete covalent molecules would
be gases even at the lowest temperatures possible.
The forces between molecules are called intermolecular forces (inter: between, as in
inter-city, international).
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Figure 8.1 shows both intermolecular and intramolecular forces in a sample of water.
Figure 8.1
..........................................
The way that the charge is distributed in a water molecule is not symmetrical. There is
more negative charge on one side than the other. It is said to have a dipole. The
Greek letter delta, δ, is used to show a 'small amount'.
Van der Waals' forces
Johannes Diderik van der Waals recognised that relatively weak forces were responsible
for the change of state from gas to liquid. He was awarded the Nobel prize for his work
in 1910.
There are several different forms of these forces and these will be discussed in the
following sections.
•
London dispersion forces.
•
Permanent dipole-permanent dipole interactions.
•
Hydrogen bonding.
In this topic the forces of attraction caused by dipoles will be explored.
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8.3
London dispersion forces
The strength of the intermolecular forces between covalent substances can be
estimated by looking at the boiling points. The lower the boiling point, the weaker the
forces must be, since it is the intermolecular forces which have to be overcome to change
the liquid into gas.
Steam
Water
Heat energy added
Liquid to gas
The covalent diatomic elements like hydrogen, H 2 gas, and bromine, Br 2 , which is a
liquid at room temperature, are completely non polar and would seem to have no dipole
at all. As well as these, even monatomic gases like helium and neon can be liquified.
How can intermolecular forces exist in these substances?
The work of Fritz London (in his 1930 paper) led to suggestions that an atom or
molecule could have a temporary dipole at a particular instant in time and that if there
were other atoms or molecules nearby; this temporary dipole might affect them and
induce a dipole in the nearby atom or molecule. The resultant electrostatic attraction
between the temporary dipole and the induced dipole might, although very weak, be
able to hold the substance together.
These forces of attraction are known as London dispersion forces.
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Induced dipoles
Figure 8.2 shows how the temporary dipole in a helium atom can exist at a single brief
instant in time. This can induce a dipole in its neighbour. Hydrogen molecules are
shown forming the same temporary dipole-induced dipole pair. The attractions
between the dipole pairs form the same type of bonds; London dispersion forces.
Figure 8.2: Temporary dipoles and induced dipoles
..........................................
Q4: When the electron movement is paused, as summarised in Figure 8.2, the
electron distribution is:
a) evenly spread.
b) unevenly spread.
..........................................
Q5: This distribution of electrons causes:
a) a permanent dipole.
b) a temporary dipole.
..........................................
Q6: What effect does this have on the neighbouring particle?
a)
b)
c)
d)
The permanent dipole causes an induced dipole.
The temporary dipole causes an induced dipole.
The permanent dipole induces a temporary dipole.
The temporary dipole induces a permanent dipole.
..........................................
Q7: The London dispersion forces of attraction formed can be said to be:
a) intramolecular.
b) intermolecular.
..........................................
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..........................................
London dispersion forces of attraction can operate between all atoms and molecules.
They are much weaker than other types of bonding, at around 2 kJ mol -1 compared to
covalent bonds of between 150-500 kJ mol -1 . In many substances, London dispersion
forces are not large enough to influence the behaviour of substances. If they are the
only type of intermolecular force present, however, as in the noble gases, they must be
considered. One example of this relates the London dispersion forces to the size of
atoms or molecules.
This diagram shows the relative sizes of atoms and boiling points ( ◦ C) of the first four
noble gases.
Noble gas boiling points
Q8:
Name the noble gas shown which has the largest atom.
..........................................
Q9:
Which noble gas shown has the highest boiling point?
..........................................
Q10: Which noble gas shown must have the strongest London dispersion forces holding
the atoms together?
..........................................
Q11: Look at the relationship between the London dispersion forces and size. Which of
these statements is true?
a)
b)
c)
d)
The size of the atom doesn't affect the London dispersion forces.
The larger the atom the weaker the London dispersion forces.
The larger the atom the stronger the London dispersion forces.
The smaller the atom the stronger the London dispersion forces.
..........................................
The lowest temperature it is possible to reach is called Absolute Zero. This occurs at
-273◦ C.
Solid helium is changed into liquid helium at approximately -272 ◦ C and liquid helium to
gas at -269◦ C, both by breaking London dispersion forces of attraction.
Q12: What does this suggest about the strength of London dispersion forces?
..........................................
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Molecular substances which have only London dispersion forces between the
molecules, i.e. intermolecular, also show a relationship between size and strength of
the London dispersion forces. The following table shows the boiling points of the first
four alkane molecules.
Name
Formula
Boiling point (◦ C)
Methane
CH4
-164
Ethane
C2 H6
-89
Propane
C3 H8
-42
Butane
C4 H10
-1
Alkane boiling points
Q13: Describe the relationship between size and boiling point and explain why it occurs.
(There would be about three or four marks allocated to this type of 'describe' and 'explain'
question and these require practice. If you are in any way unsure about being able to
explain your answer fully, work through the next four questions before attempting the
complete answer in the repeated question. If you feel confident, check your answer.)
..........................................
Q14: Name the alkane shown which has the largest molecules.
..........................................
Q15: Name the alkane shown which has the highest boiling point.
..........................................
Q16: Name the alkane shown which must have the strongest London dispersion forces.
..........................................
Q17: Which relationship between molecule size and London dispersion forces is true?
a)
b)
c)
d)
The size of the molecule doesn't affect the London dispersion forces.
The larger the molecule the weaker the London dispersion forces.
The larger the molecule the stronger the London dispersion forces.
The smaller the molecule the stronger the London dispersion forces.
..........................................
Key point
London dispersion forces of attraction operate between all atoms and molecules
and are weaker than all other types of bonding. The strength of London dispersion
forces is related to the size of the atoms or molecules.
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8.4
Hydrogen bonding
Hydrogen bonding is the name given to an intermolecular force which is actually a type
of permanent dipole-permanent dipole attraction. These hydrogen bonds are
separated from other examples of this type of interaction because they are unusually
strong (about 15-25 kJ mol -1 ) as shown below.
For hydrogen bonding to occur, the hydrogen atom involved needs to:
•
be the positive end of a strong dipole (estimated from the difference in
electronegativity).
•
have a small, highly electronegative atom on a neighbouring molecule.
Only the three elements fluorine, oxygen and nitrogen are considered able to satisfy
these conditions.
Hydrogen bonding elements
Notice that the bonds called hydrogen bonds are the forces of electrostatic attraction
between the molecules which contain these highly polar bonds. i.e. the hydrogen
bonds are the intermolecular forces.
Key point
Bonds consisting of a hydrogen atom bonded to an atom of a strongly
electronegative element such as fluorine, oxygen or nitrogen are highly polar and
the electrostatic attractions between molecules which contain these highly polar
bonds are called hydrogen bonds.
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Strength of hydrogen bonds
This activity shows the formation of hydrogen bonds in a sample of water, and the
following diagram shows the boiling points and relative sizes of the first four hydrides
of the Group 6 elements.
Strength of hydrogen bonds
A comparison of the strength of hydrogen bonds with other permanent
dipole-permanent dipole interactions and London dispersion forces can be made by
considering the data in the following diagram.
Boiling points of the Group 6 hydrides
Remember that London dispersion forces of attraction will operate between all atoms
and molecules, and that some of the hydride molecules will have permanent
dipole-permanent dipole attractions since they are polar molecules.
Q18: Use the boiling point data to compare the intermolecular forces present in water
with those present in the other Group 6 hydrides and explain fully how hydrogen bonds
compare in strength to others present.
(There would be about three or four marks allocated to this type of 'describe' and 'explain'
question and these require practice. If you are in any way unsure about being able to
explain your answer fully, work through the next four questions before attempting the
complete answer in the repeated question. If you feel confident, check your answer.)
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TOPIC 8. INTERMOLECULAR FORCES
..........................................
Q19: Name the type of weak intermolecular force of attraction which increases as size
increases.
..........................................
Q20: Name the substance which shows boiling point evidence which goes against this
trend.
..........................................
Q21: Name the substance which has the strongest intermolecular force of attraction
holding the molecules together.
..........................................
Q22: How do the hydrogen bonds in water compare in strength to the London dispersion
forces present and the other permanent dipole-permanent dipole interactions present in
some of the other molecules?
a) The hydrogen bonds are weaker than London dispersion forces and the other
permanent dipole-permanent dipole interactions.
b) The hydrogen bonds are stronger than London dispersion forces and the other
permanent dipole-permanent dipole interactions.
c) The hydrogen bonds are stronger than London dispersion forces but weaker than
the other permanent dipole-permanent dipole interactions.
..........................................
A comparison of the strength of hydrogen bonds with intramolecular bonds like the
covalent bond between hydrogen and oxygen can be made by remembering that the
hydrogen bond has a value of about 15-25 kJ mol -1 and the strength of the covalent
O-H bond can be found in the data booklet (page 9). The table of 'Mean Bond
Enthalpies' gives the value.
Q23: Which statement is true about hydrogen bonds in comparison to covalent bonds?
a)
b)
c)
d)
Covalent bonds are much stronger than hydrogen bonds.
Hydrogen bonds are much stronger than covalent bonds.
Covalent bonds are the same strength as hydrogen bonds.
Hydrogen bonds are twice as strong as covalent bonds.
..........................................
Key point
A hydrogen bond is stronger than other forms of permanent dipole-dipole
interaction but weaker than a covalent bond.
..........................................
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8.5
Relating properties to bonding
Intermolecular bonds are the relatively weak bonds which attract molecules to each
other. They are not as strong as the bonds, for example covalent bonds, that bind the
atoms together into a molecule.
Q24: Which of the following properties of a substance are a reflection of the
intermolecular forces between the molecules of the substance?
•
Melting point
•
Boiling point
•
Density
•
Viscosity
..........................................
Water is extremely important for life on Earth; in fact, the search for extraterrestrial life
involves looking for water or evidence that water was present at some time. Water is the
commonest liquid on Earth and also one of the most unusual! The next sections explore
some properties of water from the point of view of hydrogen bonding.
8.5.1
Boiling point
The boiling points of the hydrides along the Periods for Groups 4 to 7 elements are:
Boiling points of several hydrides
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Hydrides of non-metals
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Q25: The trend in boiling points in Group 4 from CH 4 to SnH4 shows a regular increase
with formula mass. Is there any evidence of polar attractions between molecules?
a) Yes
b) No
..........................................
Q26:
Three of the second period hydrides have higher boiling points than expected by
comparison with the trends seen for other elements. (Remember that higher formula
mass usually means higher boiling point.)
Which are the three hydrides?
..........................................
Q27: What is the main force between molecules in group 4 hydrides, from CH 4 to SnH4 ?
a)
b)
c)
d)
e)
Polar-polar attractions
Van der Waals attractions
Covalent bonds
Ionic
Hydrogen bonds
..........................................
Key point
The higher-than-expected boiling points of ammonia, water and hydrogen
fluoride result from additional intermolecular forces, in this case, hydrogen
bonds.
8.5.2
Density
Most substances contract when they are cooled and solids are normally denser than
their own liquids. This causes the majority of solids to sink when placed in their own
liquid. The structure of ice, however, is very unusual because the solid is less dense
then the liquid water.
As water freezes, the intermolecular hydrogen bonding spreads out the water
molecules into a strong 'open' structure with large spaces in it (see the following
diagram). This makes ice less dense and able to float on water.
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The arrangement of water molecules in ice maximises the hydrogen bonding between
them and leads to an open structure. As it melts the structure collapses into the spaces
and the liquid becomes less dense.
The fact that ice floats on water is good news for fish living in ponds. The ice forms on
the surface and the fish and other aquatic life can exist below this layer. The layer of ice
thermally insulates the water beneath.
A disadvantage of water expanding as it forms ice is that water trapped in pipes over a
cold spell causes the pipes to crack open as it freezes, resulting in leaks when it thaws.
Key point
Hydrogen bonding between molecules in ice results in an expanded structure
which causes the density of ice to be less than that of water at low temperatures.
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8.6
Viscosity
The stronger the intermolecular forces are between molecules in a liquid, the more
viscous the liquid. The viscosity is a measure of how easily it flows (how thick or
syrupy a liquid is). The molecules have to be able to move past each other to flow.
Testing viscosity
Figure 8.3: Testing the viscosity of three different liquids
Go online
..........................................
Q28: In which liquid does the marble fall fastest?
..........................................
Q29: Which chemical structure has the least hydrogen bonding?
..........................................
Q30: Which chemical structure has the most hydrogen bonding?
..........................................
Q31: Given that the marble in water took 2 seconds to reach the bottom, calculate its
rate of descent in m s-1 .
..........................................
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Q32: Predict the rate of descent of a marble in an identical experiment using
ethoxyethane.
a)
b)
c)
d)
0.8 m s-1
0.6 m s-1
0.4 m s-1
0.1 m s-1
..........................................
Q33: Assuming your answer to the last question is correct, how long would it take a
marble to travel through the ethoxyethane?
..........................................
..........................................
Viscosity is a term which applies to fluids (both liquids and gases), and since these are
covalent substances, being able to move past each other involves breaking the London
dispersion forces, dipole-dipole interactions or hydrogen bonds. Of these, the hydrogen
bonds are the strongest.
Key point
The viscosity of a liquid is related to the strength of the intermolecular bonding.
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8.7
Predicting solubilities from solute and solvent polarities
When you add a spoon of sugar to your cup of coffee and stir, the sugar dissolves. In
fact, sugar will dissolve in water to a considerable extent to produce treacle or 'golden
syrup'.
However, if you added a spoon of sugar to a cup of petrol, no amount of stirring would
make it dissolve.
A polar solvent is one whose molecules exhibit strong permanent dipoles. Because of
this, polar solvents like water can generally dissolve polar substances, like sucrose, and
ionic solids, for example sodium chloride. The expression often used to describe this is:
Key point
Like dissolves like.
In the same way, because the intermolecular attractions are of a similar type, non-polar
solvents are more likely to dissolve non-polar substances: like dissolves like.
For example, wax dissolved in white spirit in a furniture polish.
If, however, a mixture of polar and non-polar substances is used, no dissolving occurs.
Opposites do not dissolve.
Polar and ionic substances tend to be insoluble in non-polar solvents. Non-polar
substances tend to be insoluble in polar solvents. The attractions are not sufficiently
strong to allow dissolving.
When an ionic compound dissolves in water, the ions need to be separated from the
lattice. The polar water molecules can sometimes pull the ions into solution and
surround them.
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Dissolving solid sodium chloride
The way the water molecules direct themselves depends on the charge the ion carries.
The ions are said to be hydrated.
Dissolving of sodium chloride can be shown in an equation.
Dissolving sodium chloride
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A polar molecule, like hydrogen chloride, dissolves in water as the water 'interacts' with
the molecules.
Hydrogen chloride dissolving
Hydrogen chloride dissolving can also be shown in an equation.
Hydrogen chloride dissolving (equation).
In a similar way, a polar substance, such as sucrose (C12 H22 O11 ), containing several
-OH groups, can bond with polar water molecules.
Q34: Ammonia (NH 3 ) is a gas which is very soluble in water, whereas nitrogen (N 2 ) is
almost insoluble. Can you explain why?
..........................................
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Key point
Ionic compounds and polar molecular compounds tend to be soluble in polar
solvents such as water and insoluble in non-polar solvents. Non-polar molecular
substances tend to be soluble in non-polar solvents and insoluble in polar
solvents.
Miscibility
When two liquids mix thoroughly with no visible boundary between them, they are said
to be miscible, e.g. water and methanol. Both of these molecules have hydrogen
bonds shown in Figure 8.4 (a).
Miscibility arises when the intermolecular attractions between two types of substances
are fairly similar. They are then able to mix easily.
If one substance has different intermolecular attractions from the other, its molecules
stay grouped around each other and form a separate layer, e.g. water (which is hydrogen
bonded), and tetrachloromethane (which is non-polar) shown in Figure 8.4 (b).
Figure 8.4: Miscible and immiscible liquids
..........................................
Whether two liquids are miscible or immiscible can only be determined by experiment,
but the general rule of like dissolves like can be useful in trying to make a rough
prediction.
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8.8
Summary
Summary
•
All molecular elements, compounds and monatomic elements condense
and freeze at sufficiently low temperatures. For this to occur, some attractive
forces must exist between the molecules or discrete atoms.
•
Intermolecular forces acting between molecules are known as van der
Waals' forces.
•
There are several different types of van der Waals' forces such as London
dispersion forces and permanent dipole-permanent dipole interactions,
which include hydrogen bonding.
•
London dispersion forces are forces of attraction that can operate between
all atoms and molecules.
•
These forces are much weaker than all other types of bonding.
•
They are formed as a result of electrostatic attraction between temporary
dipoles and induced dipoles caused by movement of electrons in atoms
and molecules.
•
The strength of London dispersion forces is related to the number of
electrons within an atom or molecule.
•
Bonds consisting of a hydrogen atom bonded to an atom of a strongly
electronegative element, such fluorine, oxygen or nitrogen, are highly polar.
•
Hydrogen bonds are electrostatic forces of attraction between molecules
which contain these highly polar bonds.
•
A hydrogen bond is stronger than other forms of permanent dipolepermanent dipole interaction, but weaker than a covalent bond.
•
Melting points, boiling points and viscosity can all be rationalised in terms
of the nature and strength of the intermolecular forces which exist between
molecules.
•
By considering the polarity and number of electrons present in molecules,
it is possible to make qualitative predictions of the strength of the
intermolecular forces.
•
The melting and boiling points of polar substances are higher than the
melting and boiling points of non-polar substances with similar numbers of
electrons.
•
The anomalous boiling points of ammonia, water and hydrogen fluoride are
a result of hydrogen bonding.
•
Boiling points, melting points, viscosity and solubility/miscibility in water are
properties of substances which are affected by hydrogen bonding.
© H ERIOT-WATT U NIVERSITY
TOPIC 8. INTERMOLECULAR FORCES
Summary continued
8.9
•
•
Hydrogen bonding between molecules in ice results in an expanded
structure which causes the density of ice to be less than that of water at
low temperatures.
•
Ionic compounds and polar molecular compounds tend to be soluble in polar
solvents such as water and insoluble in non-polar solvents.
•
Non-polar molecular substances tend to be soluble in non-polar solvents
and insoluble in polar solvents.
Resources
Higher Chemistry for CfE: J Anderson, E Allan and J Harris, Hodder Gibson,
ISBN 978-1444167528
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TOPIC 8. INTERMOLECULAR FORCES
8.10
End of topic test
End of Topic 8 test
Q35: London dispersion forces of attraction are a result of:
Go online
a)
b)
c)
d)
an induced dipole causing an temporary dipole.
a permanent dipole causing an permanent dipole.
an induced dipole causing an permanent dipole.
a temporary dipole causing an induced dipole.
..........................................
Q36: Which of these is the strongest?
a)
b)
c)
d)
Hydrogen bonds
Covalent bonds
London dispersion forces
Permanent dipole-permanent dipole interactions
..........................................
Q37: What is the connection between the size of molecules and the London dispersion
forces of attraction?
a)
b)
c)
d)
The size makes no difference.
The smaller the molecule the weaker the force.
The smaller the molecule the stronger the force.
The weakest force is between the largest molecules.
..........................................
Q38: Which of these molecules has the largest permanent dipole?
a)
b)
c)
d)
HF
HCl
HBr
HI
..........................................
Q39: Identify the trends which would occur in descending the elements of the halogen
group.
a) The electronegativity increases.
b) The London dispersion forces become stronger.
c) The boiling point decreases.
d) The covalent radius decreases.
e) The relative atomic mass increases.
..........................................
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Q40: Look at the following table which shows the boiling points of some substances.
Substance
Boiling Point (K) Bonds (or forces) broken at the boiling point
Sodium
610
Metallic
Neon
27
A
Water
373
B
Which word(s) should be inserted at A?
..........................................
Q41: Which word(s) should be inserted at B?
..........................................
Q42: Identify the substance which has neither covalent or metallic bonding.
a)
b)
c)
d)
e)
f)
NH3 (l)
CCl4 (l)
He (g)
Hg (l)
H2 (g)
HBr (l)
..........................................
Q43: Identify the substance which has hydrogen bonding between the molecules.
a)
b)
c)
d)
e)
f)
NH3 (l)
CCl4 (l)
He (g)
Hg (l)
H2 (g)
HBr (l)
..........................................
Q44: Identify the substances which are polar molecules.
a) NH3 (l)
b) CCl4 (l)
c) He (g)
d) Hg (l)
e) H2 (g)
f) HBr (l)
..........................................
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TOPIC 8. INTERMOLECULAR FORCES
Q45: Identify the substance which has polar bonds but is a non-polar molecule.
a)
b)
c)
d)
e)
f)
NH3 (l)
CCl4 (l)
He (g)
Hg (l)
H2 (g)
HBr (l)
..........................................
Q46: The following diagram shows the structure of kevlar, which is strong because of
the intermolecular bonding between neighbouring molecules.
Name the type of intermolecular bonding involved.
..........................................
Q47: Which area of the diagram shows the position of one such bond?
..........................................
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Topic 9
End of unit test
184
TOPIC 9. END OF UNIT TEST
End of Unit 1 test
Q1:
Go online
Ethanol is formed industrially by of ethene.
a) increases.
b) decreases.
..........................................
Q2:
a)
b)
c)
d)
When the enthalpy change has a positive sign, the reaction is:
exothermic.
endothermic.
large.
small.
..........................................
Q3: Ionic bonding is most likely when the electronegativity difference between the
elements is:
a)
b)
c)
d)
exothermic.
endothermic.
large.
small.
..........................................
Q4:
a)
b)
c)
d)
Which of the following exist as diatomic molecules?
Ethane
Potassium bromide
Carbon monoxide
Neon
..........................................
Q5:
a)
b)
c)
d)
When an atom X of an element in Group 2 reacts to become X 2+ :
the mass number of X increases.
the charge on the nucleus increases.
the atomic number of X decreases.
the number of occupied energy levels decreases.
..........................................
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TOPIC 9. END OF UNIT TEST
185
Q6: The same reaction was carried out at four different temperatures. The following
table shows the time taken for the reaction to occur.
Temp (◦ C)
20
30
40
50
Time (seconds)
60
30
14
5
The results show that:
a)
b)
c)
d)
the rate of the reaction is directly proportional to the temperature.
a small rise in temperature results in a large increase in reaction rate.
the reaction is endothermic.
the activation energy increases with increasing temperature.
..........................................
The potential energy diagram below refers to the reversible reaction involving reactants
R and products P.
Q7: What is the value of the activation energy for the forward reaction (reactants to
products)?
a)
b)
c)
d)
10 kJ mol-1
20 kJ mol-1
40 kJ mol-1
60 kJ mol-1
..........................................
Q8: What is the enthalpy change for the forward reaction (reactants to products)?
a)
b)
c)
d)
-30 kJ mol-1
-20 kJ mol-1
+20 kJ mol-1
+30 kJ mol-1
..........................................
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TOPIC 9. END OF UNIT TEST
Q9:
a)
b)
c)
Which of the following molecules may be described as polar?
C1 Be
H
C1
C1
d)
C1
C1
C1
C
C1
C1
C1
..........................................
Q10: The melting points of Group 7 elements increase on descending the group
increase:
because the
a)
b)
c)
d)
mean bond energies
nuclear charges
covalent bond lengths
London dispersion forces
..........................................
Q11: The difference between the covalent radii of lithium and carbon is mainly due to
the difference in the:
a)
b)
c)
d)
number of electrons.
number of protons.
mass of each atom.
number of neurons.
..........................................
Q12: In general, covalent substances have lower melting points than ionic substances
because:
a)
b)
c)
d)
covalent bonds have no electrostatic attractive forces.
bonds between molecules are weaker than bonds between ions.
covalent compounds are composed of non-metals which have low melting points.
ionic bonds are stronger than covalent bonds.
..........................................
Q13: Carbon dioxide is a gas at room temperature while silicon dioxide is a solid
because:
a) London dispersion forces are much weaker than covalent bonds.
b) carbon dioxide contains double covalent bonds and silicon dioxide contains single
covalent bonds.
c) carbon-oxygen bonds are less polar than silicon-oxygen bonds.
d) the relative formula mass of carbon dioxide is less than that of silicon dioxide.
..........................................
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187
The apparatus shown can be used to prepare iron(III) chloride.
Q14: Which substance in the diagram contains metallic bonding?
..........................................
Q15: By considering the method of production, predict the type of bonding in iron(III)
chloride.
..........................................
Q16: Which of the following describes the bonding in the chlorides across Period 3 from
NaCl to SCl2 ?
a)
b)
c)
d)
Ionic ⇒ polar covalent
Ionic ⇒ polar covalent ⇒ pure covalent
Polar covalent ⇒ ionic
Ionic ⇒ pure covalent ⇒ polar covalent
..........................................
Q17: The elements in the third row of the Periodic Table are shown below.
Na
Mg
Al
Si
P
S
Cl
Ar
Why does the atomic size decrease crossing the period from sodium to argon?
..........................................
Q18: Atoms of different elements have different attractions for bonded electrons. What
term is used as a measure of these differing attractions?
..........................................
Q19: Use a table of these values from the SQA data booklet to explain why nitrogen
chloride contains pure covalent bonds.
..........................................
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TOPIC 9. END OF UNIT TEST
Q20: Atoms of different elements have different ionisation energies. Explain clearly why
the first ionisation energy of sodium is less than the first ionisation energy of lithium.
..........................................
Q21: The ability of an atom to form a negative ion is measured by its electron affinity.
The electron affinity is defined as the energy change when one mole of gaseous atoms
of an element combines with one mole of electrons to form gaseous negative ions. Write
the equation, showing state symbols, that represents the electron affinity of chlorine.
..........................................
Graph A shows the volume of hydrogen gas produced against time when an excess of
magnesium is added to 50 cm3 of hydrochloric acid of concentration 1 mol l-1 at 20 ◦ C.
Volume of Hydrogen
188
B
A
Time (s)
Graph B was obtained when the reaction was repeated with excess magnesium and
hydrochloric acid of the same concentration but at a different temperature.
Q22: Which reaction was at the higher temperature?
a) A
b) B
..........................................
Q23: How many cm3 of acid was required to produce graph B?
..........................................
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189
The potential energy diagram below represents the decomposition of hydrogen iodide:
2HI (g) → H2 (g) + I2 (g)
Q24: What is the value for the activation energy (E A ) for the above reaction?
..........................................
Q25: What would be the enthalpy change (ΔH) for the reaction H 2 (g) + I2 (g) → 2HI (g)?
..........................................
Q26: The use of a catalyst would
the activation energy in part 1.
a) increase
b) decrease
c) have no effect on
..........................................
Q27: The use of a catalyst would
the enthalpy change in part 2.
a) increase
b) decrease
c) have no effect on
..........................................
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TOPIC 9. END OF UNIT TEST
Q28: The enthalpy change for the reaction
enthalpy of formation of hydrogen iodide.
1
2 H2 (g)
+ 12 I2 (s) → HI (g) is known as the
It differs from the above equation in two respects:
•
it involves the formation of one mole of HI;
•
the reactants are in their standard states, so iodine is a solid.
Given that the enthalpy of sublimation of iodine I 2 (s) → I2 (g) has ΔH = +62 kJ, calculate
the enthalpy of formation of hydrogen iodide.
..........................................
Q29: If the enthalpy of formation of hydrogen chloride is -92 kJ mol -1 , predict a value
for the enthalpy of formation of hydrogen bromide.
..........................................
Q30: Match the following chemical terms to their correct definitions:
Definition
Chemical term
(randomised)
For a chemical reaction to occur, particles must collide.
Enthalpy change
The minimum kinetic energy required by colliding particles
before a reaction will occur.
Activated complex
Intermediate stage at the top of the energy barrier when
reactants change to products.
Catalyst
The difference in potential energy between products and
reactants.
Collision theory
A substance that alters the rate of a reaction without being
used up.
Activation energy
..........................................
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191
Q31: Match the following chemical terms to their correct definitions:
Chemical term
(randomised)
Definition
Regular arrangement of positively charged ions surrounded
by delocalised electrons.
Covalent bonding
The electrostatic force of attraction of a shared pair of
electrons for two positive nuclei.
Metallic bonding
The electrostatic attraction between negative and positive
ions.
Intramolecular forces
Dipole
Weak bonds between molecules.
A slight negative or positive charge caused by the uneven
distribution of electrons.
Intermolecular forces
Bonds holding atoms together within molecules.
Ionic bonding
..........................................
Q32: Match the following chemical terms to their correct definitions:
Chemical term
(randomised)
Definition
Half the distance between the nuclei of two covalently
bonded atoms of an element.
Ionisation energy
The strength of the attraction of an element for the electrons
of its bonding electrons.
Covalent radius
The energy required to remove an electron from a gaseous
atom to form an ion with a single positive charge.
Electronegativity
The temperature at which a substance changes from a solid
to a liquid.
Boiling point
The temperature at which a substance changes from a
liquid to a gas.
Melting point
..........................................
Q33: Sort the forces of attraction into order of strength from strongest to weakest:
•
Hydrogen bonding
•
Permanent dipole - permanent dipole interactions
•
London dispersion forces
•
Covalent bonding
..........................................
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TOPIC 9. END OF UNIT TEST
Q34: Complete each statement using 'increases' or 'decreases' as appropriate.
•
Going across a period, atomic size . . . increases / decreases.
•
Going down a group, atomic size . . . increases / decreases.
•
Going down a group, electronegativity . . . increases / decreases.
•
Going across a period, electronegativity . . . increases / decreases.
•
Going across a period, boiling point . . . increases / decreases.
•
Going down a group, boiling point . . . increases / decreases.
..........................................
A student writes the following two statements. Both are incorrect. In each case, explain
the mistake in the student's reasoning.
Q35: All ionic compounds are solids at room temperature. Many covalent compounds
are gases at room temperature. This proves that ionic bonds are stronger than covalent
bonds.
..........................................
Q36: The formula for magnesium chloride is MgCl2 because, in solid magnesium
chloride, each magnesium ion is bonded to chloride ions.
..........................................
Q37: Although propane and ethanol have similar molecular masses, the alkane is a gas
at room temperature while the alcohol is a liquid. Explain why propane is a gas at room
temperature whereas ethanol is a liquid. (3 marks)
..........................................
..........................................
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GLOSSARY
Glossary
Activated complex
a very unstable arrangement of atoms formed at the maximum of the potential
energy barrier, during a chemical reaction
Activation energy
the minimum kinetic energy required by colliding particles before reaction will
occur, since a high energy activated complex must be formed
Adsorption
occurs when molecules become bonded to the surface of a catalyst
Allotropes
one of two or more existing forms of an element, for example, graphite and
diamond are allotropes of carbon
Bonding electrons
shared pairs of electrons from both atoms forming the covalent bond
Chemical bonding
describes the mechanism by which atoms are held together
Chemical structure
describes the way in which atoms, ions or molecules are arranged
Collision theory
suggests that particles must collide for a chemical reaction to occur
Covalent bond
formed when two atoms share electrons in their outer shell to complete the filling
of that shell
Covalent radius
half the distance between the nuclei of two bonded atoms of an element
Delocalised
in metallic bonding, delocalised electrons are free from attachment to any one
metal ion and are shared amongst the entire structure
Desorption
occurs when the bonds between the molecules and the surface break and the
molecules leave the surface of the catalyst
Diatomic
molecules with only two atoms (e.g. oxygen, O 2 , and carbon monoxide, CO)
Dipole
an atom or molecule in which a concentration of positive charges is separated from
a concentration of negative charge
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GLOSSARY
Electronegativity
a measure of the attraction that an atom involved in a bond has for the electrons
of the bond
Enthalpy change
the name given to the amount of heat evolved or absorbed in a reaction carried
out at constant pressure. It is given the symbol ΔH, read as "delta H"
Fullerenes
molecules of pure carbon constructed from 5- and 6-membered rings combined
into hollow structures; the most stable contains 60 carbon atoms in a shape
resembling a football
Hydrogen bonds
electrostatic forces of attraction between molecules containing a hydrogen atom
bonded to an atom of a strongly electronegative element such as fluorine, oxygen
or nitrogen, and a highly electronegative atom on a neighbouring molecule
Intermolecular forces
forces of attraction which exist between molecules; they are weaker than chemical
bonds
Intramolecular forces
forces of attraction which exist within a molecule
Ionisation energy
the energy required to remove one mole of electrons from one mole of atoms in
the gaseous state
Isoelectronic
molecules which have the same arrangement of electrons. For example, the noble
gas neon, a sodium ion (Na + ) and a magnesium ion (Mg 2+ ) are isoelectronic
Lattice
a regular 3D arrangement of particles in space; the term is applied to metal ions
in a solid, and to positive and negative ions in an ionic solid
London dispersion forces
the forces of attraction which result from the electrostatic attraction between
temporary dipoles and induced dipoles caused by movement of electrons in atoms
and molecules
Lone pairs
pairs of electrons in the outer shell of an atom which take no part in bonding
Miscible
fluids which mix with or dissolve in each other in all proportions
Periodicity
the regular recurrence of similar properties when the elements are arranged in
order of increasing atomic number
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GLOSSARY
Polar covalent bond
a covalent bond between atoms of different electronegativity, which results in an
uneven distribution of electrons and a partial charge along the bond
Potential energy diagram
shows the enthalpy of reactants and products, and the enthalpy change during a
chemical reaction
Properties
physical and chemical characteristics of a substance; these are often a reflection
of the chemical bonding and structure of the material
Thermochemical equation
states the enthalpy change for the reaction defined, with reactants and products in
the states shown
Viscosity
the resistance to flow that is exhibited by all liquids
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ANSWERS: TOPIC 1
Answers to questions and activities
1 Reaction rates - collision theory
Test your prior knowledge (page 2)
Q1:
•
Concentration;
•
particle size / surface area;
•
temperature.
Q2:
Any two from:
•
mass of products;
•
mass of reactants;
•
mass of gas given off;
•
concentration of reactants;
•
concentration of products;
•
volume of reactants;
•
volume of products;
•
pressure;
•
conductivity;
•
pH;
•
colour Intensity.
Q3:
0.01 g s-1 (to 2 decimal places)
Working:
Change in mass
Rate =
Change in time
0.7
=
60
= 0.01 g s - 1
Measuring rate (page 5)
Q4:
7.0 cm3
Q5:
1.4 cm3 s-1
Q6:
0.7 cm3 s-1
Q7: The rate of reaction is dropping as it progresses, being relatively fast at the start
and slowing towards the end. This is because the reactants are being used up.
Q8:
0.52 cm3 s-1
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ANSWERS: TOPIC 1
Collisions and concentration (page 7)
Q9: 3
Q10: 0.3
Q11: 0.6
Q12: The rate would be 1.2 collisions per second. It would seem that doubling the
concentration of acid has doubled the rate, and in some reactions this is correct; the
rate can be directly proportional to the concentration of a reactant. On the other hand,
sometimes doubling the concentration of a reactant only increases rate a little, or even
not at all. There is no simple way to predict the relationship in advance and each
concentration/rate relationship must be investigated experimentally to determine any
relationship. It can be said however that for many reactions increase in concentration
increases rate.
Collisions and particle size (page 8)
Q13: Small
Q14: 6
Q15: 0.6
Q16: 0.8
Collisions and temperature (page 9)
Q17: 4
Q18: 0.4
Q19: 0.8
Q20: c) It doubles.
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ANSWERS: TOPIC 1
End of Topic 1 test (page 10)
Q21: a) average kinetic energy of all the particles.
Q22: a) Time taken for reaction to complete.
Q23: a) increasing the concentration of the acid.
Q24: 0.2 g s-1
Q25: 0.08 g s-1
Q26: d) Curve E
Q27: c) Curve D
Q28: a) a small rise in temperature results in a large increase in reaction rate.
Q29: a) decreasing the mass of copper carbonate.
Q30: a)
Q31: 0.25 moles
Q32: 0.1 moles
Q33: 0.44 mol -1-1
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ANSWERS: TOPIC 2
2 Reaction rates - reaction profiles
Test your prior knowledge (page 19)
Q1: Exothermic
Q2: Endothermic
Q3:
•
Particles must collide to react (although not all collisions are successful).
•
Sufficient energy is needed and orientation must be correct.
Measuring rate at different concentrations (page 21)
Q4: c) 2.0 mol -1
Q5: c) 2.0 mol -1
Q6: 20 seconds
Q7: 26 cm3
Q8:
•
The steepness of the slope for the 1.0 mol -1 is less because the reaction rate is
lower (fewer acid particles).
•
The point at which the 1.0 mol -1 line becomes horizontal takes longer to reach
(25 seconds rather than 20 seconds) because the reaction is slower (fewer acid
particles).
•
The final level of the 1.0 mol -1 horizontal line is lower by half because the acid is
present in only half the quantity.
Q9: 6.5 cm3
Answers from page 23.
Q10: 20 cm3
Q11: 10.4 cm3 s-1
Answers from page 26.
Q12: a) A
Q13: d) D
Q14: b) B
Q15: d) D
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ANSWERS: TOPIC 2
Controlling the rate - temperature and kinetic energy (page 27)
Q16: c) Not changed
Q17: a) Increased
Q18: b) C and D
Q19: a) Particles with sufficient energy to react.
Q20: c) Roughly doubled.
End of Topic 2 test (page 31)
Q21: b) The activation energy is less for phosphorus than for magnesium.
Q22: c) The reaction rate doubles.
Q23: d) More particles have energy greater than the activation energy.
Q24: a) the minimum kinetic energy required by colliding molecules for a reaction to
occur.
Q25: b) an intermediate stage at the top of the activation energy barrier.
Q26: c) potential energy diagrams.
Q27: 0.025
Either curve can be used:
Rate =
=
0.375
15
Change in concentration
Change in time
( 12 mark)
( 12 mark)
= 0.025 ( 12 ) mol l-1 s-1 ( 12 mark)
Allow ±0.001 only. Two significant figures must be shown. (- 12 otherwise)
Q28: At 40◦ C, rate = 0.032 s-1 ( 12 mark)
1
1
Rate ( 2 mark)
1
( 12 mark)
= 0.032
= 31.25 seconds ( 12 mark
Time =
- units must be given.)
Q29: Increasing temperature increases kinetic energy (or speed of movement) of
molecules ( 12 mark)
The (potential) energy generated in collisions increases (collisions have more force) ( 12
mark)
More collisions will result in the activation energy being attained (1 mark)
Q30: b) 60◦ C
Q31: b) a small rise in temperature results in a large increase in reaction rate.
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ANSWERS: TOPIC 3
3 Catalysis
Test your prior knowledge (page 37)
Q1: a) average kinetic energy of all the particles;
Q2: b) Endothermic
Q3: a) The minimum kinetic energy required by colliding molecules for a reaction to
occur.
Addition of hydrogen to ethene (page 38)
Q4: c) adsorption.
Q5: b) They get weaker.
Q6: None
Q7: c) The catalyst speeds up the reaction without taking part.
Catalysis by cobalt(II) chloride solution (page 40)
Q8: b) Co2+ (aq) ions
Q9: c) Co3+ (aq) ions
Q10: b) Co2+ (aq) ions
Q11: This is a poor definition. It is correct to say that the Co2+ (aq) ions do speed up
the reaction, but it is incorrect to say that they do not take part in the reaction. The
pink Co2+ (aq) ions are changed into something green (probably Co 3+ (aq) ions) which
are later changed back into pink Co 2+ (aq) ions. So the Co2+ (aq) ions can be recovered
at the end.
Answers from page 42.
Q12: d) Cl
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ANSWERS: TOPIC 3
Industrial uses of catalysts (page 42)
Q13:
Catalyst
Process
Type
Purpose
Haber
Iron
Solid
Production of
ammonia
Washing powders
Protease
Enzyme
Breakdown of protein
stains
Catalytic converters
Pt, Pd, Rb
Solid
Reduction of pollution
Brewing
Zymase
Enzyme
Production of ethanol
Catalytic cracking
Al2 O3
Solid
Production of alkenes
Food processing
Rennin
Enzyme
Production of cheese
Ostwald
Pt gauze
Solid
Production of nitric
acid
Monsanto
Rhodium complex
Solution
Ethanoic acid from
methanol and CO
Answers from page 45.
Q14: c) It will decrease.
Q15: c) Positive
Answers from page 46.
Q16: d) Endothermic
Answers from page 47.
Q17: a) Positive
Q18:
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ANSWERS: TOPIC 3
203
Answers from page 49.
Q19: d) ΔH =+283 kJ mol-1
Q20: b) ΔH = Hproducts - Hreactants
Q21: c) Exothermic
Q22: d) ΔH = -1414 kJ mol-1
Q23: 6835 kJ
Q24: +150 kJ
Q25: 4NH3 (g) + 5O2 (g) → 4NO(g) + 6H2 O(g)
Q26: -906 kJ
Q27: -226.5 kJ mol-1
Q28: 1359 kJ
Answers from page 52.
Q29: b) B
Q30:
The enthalpy change (ΔH) is given by the expression:
ΔH(reaction) = ΔH(products) - ΔH(reactants)
From the graph:
Enthalpy of reactants = −1200 kJ
Enthalpy of products = −630 kJ
So, ΔH = (−630) − (−1200) kJ
= +570 kJ
Q31: d) -890 kJ
Answers from page 54.
Q32: d) +110 kJ
Q33: c) +55 kJ
Q34:
If the activation energy is less, the barrier will be lower and it will be easier for colliding
molecules to get over the barrier. There will be more molecules with enough energy to
collide successfully and so the reaction rate will increase.
In general, the lower the activation energy, the faster the reaction.
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ANSWERS: TOPIC 3
Answers from page 55.
Q35: a) A
Q36: d) D
Q37: c) It will increase.
How catalysts work (page 56)
Q38: b) High activation energy, few successful collisions.
Q39: c) Low activation energy, lots of successful collisions.
Q40: c) It speeds up the reaction by lowering the activation energy.
Q41: b) No effect.
Mechanism of catalysis (page 58)
Q42: a) X
Q43: c) 3
Q44: d) 4
Q45: b) Decreases it
Q46: b) W
Q47: a) X
End of Topic 3 test (page 62)
Q48: a) molecules always have energy above the activation energy.
Q49: a) the minimum kinetic energy required by colliding molecules for a reaction to
occur.
Q50: d)
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ANSWERS: TOPIC 3
Q51: e)
Q52: c)
Q53: a) A very unstable intermediate is formed when bonds within the reactant
molecules begin to break and new bonds begin to form.
Q54: Exothermic
Q55: 10 kJ
Q56: 35 kJ
Q57: -25 kJ
Q58: +25 kJ
Q59:
a)
1. Energy of reactants / Reactant molecules collide with correct geometry.
2. Energy of activated complex / The collision has enough energy to overcome
the Ea .
3. Activation energy of the forward reaction / Activated Complex: bonds in the
reactant molecules break and new bonds begin to form.
4. Enthalpy change of the forward reaction
b) The product molecule(s) has less energy than the reactants.
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ANSWERS: TOPIC 4
4 The Periodic Table
Test your prior knowledge (page 69)
Q1:
a) Atomic number
Q2:
•
Bromine
•
Chlorine
•
Fluorine
•
Hydrogen
•
Iodine
•
Nitrogen
•
Oxygen
Q3:
Particle
Charge
Mass
Location
Proton
+1
1
Nucleus
Neutron
0
1
Nucleus
Electron
-1
0
Orbiting nucleus or
in shells
Q4:
b) Period
Q5:
b) Group
Newlands' octaves (page 70)
Q6:
c) Atomic mass
Q7:
a) Both are gases.
Q8:
c) Both are reactive metals.
Q9:
When listed in order of increasing atomic mass, similar properties appear with every
eighth element. If the elements are numbered in order, element 1 has properties in
common with element 8, element 2 has properties in common with element 9, etc.
This is similar to musical notation. Notes in a scale are described as the letters A to G.
The note after G is A and so on. A to G is said to be one octave. The same applies if
the notes are described as 'doh-ray-me-fah-soh- etc'.
Q10: 3
Q11: Chromium, manganese and iron are metals.
Newlands' table are non-metals.
The elements above them in
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ANSWERS: TOPIC 4
Mendeleev's Periodic Table (page 72)
Q12: Mendeleev believed that elements, as yet undiscovered, would fit in these spaces.
Q13: Iodine has similar chemical properties to bromine and so fits better in the column
which contains bromine.
Q14: Noble gases
The modern Periodic Table (page 74)
Q15: Atomic number
Melting and boiling points across a Period (page 75)
Q16: Carbon
Q17: Argon
Q18: Helium
Melting and boiling points down a Group (page 76)
Q19: Lithium
Q20: Caesium
Q21: Melting point and boiling point both decrease and so the forces must be getting
weaker.
Atomic size (page 77)
Q22: c) The number of electron shells increases.
Q23: d) The number of protons increases.
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ANSWERS: TOPIC 4
End of Topic 4 test (page 80)
Q24: c) Mendeleev left gaps for elements which had not yet been discovered.
Q25: c) There is a steady decrease in atomic size across a period from left to right and
d) There is a decrease in first ionisation energy on going down Group 0.
Q26: a) Atomic number
Q27:
•
Bromine
•
Chlorine
•
Fluorine
•
Hydrogen
•
Iodine
•
Nitrogen
•
Oxygen
Q28: b) The alkali metals.
Q29: a) The halogens.
Q30: c) The noble gases.
Q31: a) number of occupied energy shells.
Q32: c) The right hand side.
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5 Bonding and structure
Test your prior knowledge (page 85)
Q1: a) a shared pair of electrons.
Q2: d) delocalised electrons.
Q3: b) transfer of electrons.
Q4: a) have high melting and boiling points.
Q5: d) conduct electricity when molten or in solution.
Answers from page 86.
Q6:
You have probably thought of gold, nuggets of which are found in rocks and as gold
particles in some river beds. Platinum and other platinum group metals (rhodium,
iridium, ruthenium and osmium) are also found native.
Sulfur is also found as yellow crystals in rocks and around volcanos. Copper is also
found in elemental form, as well as in other ores.
Gold
Sulfur
Native copper
Answers from page 87.
Q7: Niobium is used in steel superalloys which can resist high temperatures and are
used in jet engine components and heat shields.
Construction of a sodium atom (page 89)
Q8: Argon
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ANSWERS: TOPIC 5
Q9:
Gain an outer electron.
Q10: Lose the outer electron.
Q11: A positive sodium ion (Na + )
Q12: Neon
Metallic bonds (page 90)
Q13: The group 1 metals sodium and potassium are easily cut with a knife (which is
made from hard metals - iron and chromium). Mercury is liquid (melting point -39 ◦ C)
and gallium melts in the hand (at 30 ◦ C).
Copper lattice (page 91)
Q14: c) delocalised.
Q15: d) Lattice
Q16: Each aluminium atom will ionise to produce three electrons and a Al 3+ ion. Sodium
will produce only one electron (to carry the electric current) from each atom.
Answers from page 93.
Q17: 4
Q18: 1
Q19: b) share, electron
Q20: d) positively charged nuclei and negatively charged shared electrons.
Q21:
Diamond (page 98)
Q22: 4
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ANSWERS: TOPIC 5
Q23: a) Single covalent.
Graphite structure (page 99)
Q24: c) stronger.
Q25: 6
Fullerenes (page 101)
Q26: Diamond and graphite are network covalently bonded, so that it would require a
very large input of energy to separate the atoms in a solution. 'Buckyballs' are covalent
molecular bonded, with strong bonds binding the C 60 atoms into a molecule, but much
weaker bonds between the molecules in the solid. These can be easily broken to form
a solution.
Ionic bond formation (page 103)
Q27: 1
Q28: 7
Q29: 8
Q30: Both ions have a stable noble gas electron configuration.
Answers from page 104.
Q31: c) large
Q32: c) positive caesium ions and negative chloride ions.
Q33: c) 3 : 2
Sodium chloride ionic lattice (page 105)
Q34: c) An ionic lattice.
End of Topic 5 test (page 108)
Q35: a) Calcium sulfide
Q36: a) Sulfur
Q37: b) delocalised electrons.
Q38: b) Metallic
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ANSWERS: TOPIC 5
Q39: a) Covalent and
c) Made up of discrete molecules.
Q40: a) Boron
Q41: d) Phosphorus and
f) Sulfur
Q42: b) Chlorine and
e) Sodium
Q43: c) C
Q44: a) Ionic
Q45: f) SiO2 (s)
Q46: a) NH4 Cl (s) and
e) Na2 CO3 (s)
Q47:
Q48: The bonds between the layers of carbon in graphite are comparatively weak,
which means that the layers can easily slide over each other providing lubrication.
Q49: The carbon atoms in graphite have electrons not involved on the covalent bonds
between atoms within the layers. These delocalised electrons permit conduction of
electricity.
The boron atoms in boron nitride have no free electrons once they have formed three
covalent bonds, so there is no delocalised pool of electrons to conduct electricity.
Q50: The nitrogen atoms in the layers have two free electrons which can form strong
covalent bonds between adjacent layers in boron nitride.
Q51: b) Monatomic
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ANSWERS: TOPIC 6
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6 Periodic Table trends
Test your prior knowledge (page 115)
Q1:
Particle
Charge
Mass
Location
Proton
+1
1
Nucleus
Neutron
0
1
Nucleus
Electron
-1
0
Orbiting nucleus or
in shells
Q2: c) number of outer electrons.
Q3: a) number of occupied energy shells.
Q4: a) Atomic number
Q5: a) a shared pair of electrons.
Covalent radius - relative sizes (page 116)
Q6: b) A decrease
Q7: a) An increase
Q8: The covalent radius is defined as half the distance between the nuclei of bonded
atoms. Noble gases do not form bonds because they are so unreactive.
The trends in covalent radius (page 117)
Q9: c) The number of electron shells increases.
Q10: d) The number of protons increases.
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ANSWERS: TOPIC 6
Ionisation energies (page 121)
Q11:
Q12: Yes
Q13: The shape of the first part of the graph (from atomic number 3-10) is repeated for
elements 11-18
Q14: Noble gases
Q15: Alkali metals
Q16: a) A steady increase.
Q17: b) A steady decrease.
Q18:
Metal
1st Ionisation energy / kJ mol-1
Lithium
526
Sodium
502
Potassium
425
Rubidium
409
Caesium
382
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ANSWERS: TOPIC 6
First ionisation energy - Going down a group (page 122)
Q19:
The nuclear charge increases on going down a group which should make the outer
electron more difficult to remove. (1 mark.)
However, the covalent radius increases and so the outer electron is further away from
the nucleus. The electrostatic forces get weaker as the distance between the charges
increases. This should make the outer electron easier to remove. (1 mark)
On going down the group, there are more and more inner electrons which prevent the
outer electron from experiencing the full effect of the nuclear charge. These inner
electrons increasingly screen the outer electron from the nucleus. Consequently, the
outer electron becomes easier to remove as the atom gets bigger, i.e. the first ionisation
energy decreases on going down a group. (1 mark)
First ionisation energy - Going across a Period (page 123)
Q20:
On going from left to right across Period 2, the nuclear charge increases. The electrons
are held more tightly making it more difficult to remove one of the outer electrons. (1
mark)
Each additional electron goes into the second shell. This makes little difference to the
screening effect. Consequently, the increased nuclear charge pulls in the electrons and
the covalent radius decreases. (1 mark)
Because the outer electrons are closer to the nucleus, they are more strongly held and
so the first ionisation increases across the Period. (1 mark)
Q21: b) It increases.
Q22: a) It increases.
Q23: c) It stays the same.
Q24: c) It stays the same.
Q25: a) It decreases.
Electronegativity (page 125)
Q26: b) Bromine
Q27: c) C-S
Q28: a) Group 3
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ANSWERS: TOPIC 6
Answers from page 126.
Q29: 1
Q30: 7
Q31: Noble gases
Q32: The noble gases do not generally form bonds with other elements.
Answers from page 127.
Q33: c) Electronegativity values increase on going from left to right across a Period and
decrease on going down a Group.
Summary: Periodic Table trends (page 128)
Q34:
End of Topic 6 test (page 130)
Q35: c) Mendeleev left gaps for elements which had not yet been discovered.
Q36: 3. There is a decrease in first ionisation energy on going down Group 0, and
6. There is a steady decrease in atomic size across a Period from left to right.
Q37: b) Atomic size
Q38: b) B
Q39: d) more occupied electron shells?
Q40: c) protons?
Q41: c) Mg(g) → Mg+ (g) + 2eQ42: d) F(g) → F+ (g) + eQ43: 4. The first ionisation energy decreases from Li to Cs, and
5. Electronegativity decreases from Li to Cs.
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ANSWERS: TOPIC 6
Q44: d) increasing nuclear charge?
Q45: d) 6
Q46: b)
Atomic Number
Q47: d) Atomic mass
Q48: Electronegativity
Q49: Fluorine
Q50: d)
Q51: On going from left to right across Period 2, the nuclear charge increases. The
electrons are held more tightly making it more difficult to remove one of the outer
electrons. (1 mark)
Each additional electron goes into the second shell. This makes little difference to the
screening effect. Consequently, the increased nuclear charge pulls in the electrons and
the covalent radius decreases. (1 mark)
Because the outer electrons are closer to the nucleus, they are more strongly held and
so the first ionisation increases across the Period. (1 mark)
Q52: The nuclear charge increases on going down a group which should make the
outer electron more difficult to remove. (1 mark)
However, the covalent radius increases and so the outer electron is further away from
the nucleus. The electrostatic forces get weaker as the distance between the charges
increases. This should make the outer electron easier to remove. (1 mark)
On going down the Group, there are more and more inner electrons which prevent
the outer electron from experiencing the full effect of the nuclear charge. These inner
electrons increasingly screen the outer electron from the nucleus. Consequently, the
outer electron becomes easier to remove as the atom gets bigger, i.e. the first ionisation
energy decreases on going down a Group. (1 mark)
Q53: b) They generally do not form bonds with other elements.
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ANSWERS: TOPIC 7
7 Bonding continuum and polar covalent bonding
Test your prior knowledge (page 139)
Q1:
b) a shared pair of electrons.
Q2:
c) They generally do not form bonds with other elements.
Q3:
a) Fluorine
Q4:
c) Hydrogen
Electronegativity (page 140)
Q5:
b) 50%
Q6:
c) C-S
Q7:
a) X is Na and Y is O
Q8:
c) Hydrogen chloride (HCl)
Q9:
a) Hydrogen fluoride (HF)
Q10: d) Ammonia (NH 3 )
Direction of charge (page 143)
Q11:
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Answers from page 145.
Q12:
•
Pure covalent - electronegativity difference 0; Br - Br
•
Polar covalent - electronegativity difference 0.6;
•
Ionic - electronegativity difference 2.0; K+ ...... Br-
Q13: c) Ionic
Q14: a) Pure covalent
Q15: b) Fluorine
Permanent dipole interactions (page 146)
Q16: b) B
Q17: c) C
Q18: b) B
Predicting molecular polarity (page 148)
Q19:
Detecting polar molecules (page 149)
Q20: a) Water
Q21: a) Water
Q22: b) It is attracted by the rod.
Q23: a) It is not affected by the rod.
Answers from page 151.
Q24: a) Polar
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- Brδ-
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ANSWERS: TOPIC 7
End of Topic 7 test (page 153)
Q25: d) Calcium sulfide
Q26: a) Sulfur
Q27: c) delocalised electrons.
Q28: b) Metallic
Q29: a) Covalent, and
b) Made up of discrete molecules.
Q30: a) Boron
Q31: d) Phosphorus, and
f) Sulfur
Q32: b) Chlorine, and
e) Sodium
Q33: a) Cl-F
Q34: C
Q35: c) BeCl2
Q36: c) Ionic
Q37: f) SiO2 (s)
Q38: a) NH4 Cl (s), and
e) Na2 CO3 (s)
Q39: b) CH3 OH (l)
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ANSWERS: TOPIC 8
8 Intermolecular forces
Test your prior knowledge (page 160)
Q1: c) Electronegativity
Q2: a) Fluorine
Q3: c) increases.
Induced dipoles (page 163)
Q4: b) unevenly spread.
Q5: b) a temporary dipole.
Q6: b) The temporary dipole causes an induced dipole.
Q7: b) intermolecular.
Answers from page 164.
Q8: Krypton
Q9: Krypton
Q10: Krypton
Q11: c) The larger the atom the stronger the London dispersion forces.
Answers from page 164.
Q12: The energy required to break forces which are only effective at such low
temperatures is very small, and in fact, London dispersion forces of attraction are weaker
than all other types of bonding.
Answers from page 165.
Q13:
As the alkane molecules in the family get bigger, from methane to butane, the boiling
point increases (butane at -1 ◦ C is a higher temperature than methane at -164 ◦ C). (1
mark)
Since boiling point depends on the strength of the London dispersion forces, butane
must have stronger London dispersion forces than methane. (1 mark)
Therefore the larger the molecule the stronger the London dispersion forces. (1 mark)
Q14: Butane
Q15: Butane
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ANSWERS: TOPIC 8
Q16: Butane
Q17: c) The larger the molecule the stronger the London dispersion forces.
Strength of hydrogen bonds (page 167)
Q18: As size increases, London dispersion forces increase, so one might expect
water (the smallest) to have the lowest boiling point. However, the boiling point of
water goes against the trend and is much higher than might be expected. This
is because water molecules exhibit hydrogen bonding between the molecules, due
to the highly polar character of the water molecule. Although the other hydrides
have some permanent dipole-permanent dipole interactions, the hydrogen bonded
permanent dipole-permanent dipole interactions in water are stronger than other forms
of permanent dipole-permanent dipole interactions and London dispersion forces, as
evidenced by the high boiling point.
Q19: London dispersion forces
Q20: Water
Q21: Water
Q22: b) The hydrogen bonds are stronger than London dispersion forces and the other
permanent dipole-permanent dipole interactions.
Q23: a) Covalent bonds are much stronger than hydrogen bonds.
Answers from page 169.
Q24: All of the above.
Hydrides of non-metals (page 170)
Q25: b) No
Q26: NH3 , H2 O and HF
Q27: b) Van der Waals attractions
Testing viscosity (page 172)
Q28: Methanol
Q29: Methanol
Q30: Glycerol
Q31: 0.5 m s-1
Q32: a) 0.8 m s-1
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ANSWERS: TOPIC 8
Q33: 1.25 seconds
Answers from page 176.
Q34: Ammonia is a polar substance which can form hydrogen bonds with water;
nitrogen is non-polar and unable to bond to water.
End of Topic 8 test (page 180)
Q35: d) a temporary dipole causing an induced dipole.
Q36: b) Covalent bonds
Q37: b) The smaller the molecule the weaker the force.
Q38: a) HF
Q39: b) The London dispersion forces become stronger.
e) The relative atomic mass increases.
Q40: London dispersion forces.
Q41: Hydrogen bonding (a form of permanent-permanent dipole interaction) between
the molecules.
Q42: c) He (g)
Q43: a) NH3 (l)
Q44: a) NH3 (l)
f) HBr (l)
Q45: b) CCl4 (l)
Q46: Hydrogen bonding
Q47:
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ANSWERS: TOPIC 9
9 End of unit test
End of Unit 1 test (page 184)
Q1:
b) decreases.
Q2:
b) endothermic.
Q3:
c) large.
Q4:
c) Carbon monoxide
Q5:
d) the number of occupied energy levels decreases.
Q6:
b) a small rise in temperature results in a large increase in reaction rate.
Q7:
b) 20 kJ mol-1
Q8:
a) -30 kJ mol-1
Q9:
b)
H
C1
Q10: d) London dispersion forces
Q11: b) number of protons.
Q12: b) bonds between molecules are weaker than bonds between ions.
Q13: a) London dispersion forces are much weaker than covalent bonds.
Q14:
Fe(s)
Cl2(g)
Heat
Iced waterr
FeCl3(g)
Q15: Polar covalent
Q16: b) Ionic ⇒ polar covalent ⇒ pure covalent
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ANSWERS: TOPIC 9
225
Q17: The number of protons increases OR greater nuclear charge OR greater nuclear
attraction.
Q18: Electronegativity
Q19: Same electronegativity values.
Q20: Bigger atom so electron is further from the nucleus; inner electrons shield (screen)
the outer electron from attraction of the nucleus.
Q21: Cl(g) + e- → Cl- (g)
Q22: b) B
Q23: 100 cm3
Q24: 190 kJ
Q25: -10 kJ
Q26: b) decrease
Q27: c) have no effect on
Q28: 26 kJ mol-1
Q29: -33 kJ mol-1
Q30:
Definition
Chemical term
For a chemical reaction to occur, particles must collide.
Collision theory
The minimum kinetic energy required by colliding particles
before a reaction will occur.
Activation energy
Intermediate stage at the top of the energy barrier when
reactants change to products.
Activated complex
The difference in potential energy between products and
reactants.
Enthalpy change
A substance that alters the rate of a reaction without being
used up.
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Catalyst
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ANSWERS: TOPIC 9
Q31:
Definition
Chemical term
Regular arrangement of positively charged ions surrounded
by delocalised electrons.
Metallic bonding
The electrostatic force of attraction of a shared pair of
electrons for two positive nuclei.
Covalent bonding
The electrostatic attraction between negative and positive
ions.
Weak bonds between molecules.
Ionic bonding
Intermolecular forces
A slight negative or positive charge caused by the uneven
distribution of electrons.
Bonds holding atoms together within molecules.
Dipole
Intramolecular forces
Q32:
Definition
Chemical term
Half the distance between the nuclei of two covalently
bonded atoms of an element.
Covalent radius
The strength of the attraction of an element for the electrons
of its bonding electrons.
Electronegativity
The energy required to remove an electron from a gaseous
atom to form an ion with a single positive charge.
Ionisation energy
The temperature at which a substance changes from a solid
to a liquid.
Melting point
The temperature at which a substance changes from a
liquid to a gas.
Boiling point
Q33:
1. Covalent bonding
2. Hydrogen bonding
3. Permanent dipole - permanent dipole interactions
4. London dispersion forces
Q34:
•
Going across a period, atomic size decreases.
•
Going down a group, atomic size increases.
•
Going down a group, electronegativity decreases.
•
Going across a period, electronegativity increases.
•
Going across a period, boiling point decreases.
•
Going down a group, boiling point increases.
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ANSWERS: TOPIC 9
Q35: Covalent bonds not being broken OR Intermolecular bonds that are breaking. Any
alternative wording that recognises that covalent bonds are not broken when covalent
substances melt or boil.
Q36: Formula refers to the ratio of Mg 2+ :Cl- ions (in lattice) OR alternative wording ie
in the lattice there are twice as many chloride ions as magnesium ions OR Mg 2+ ions
surrounded by > 2 Cl - ions OR Cl- surrounded by >1 Mg 2+ . Any reference to 'chlorine
ions' is not acceptable.
Q37: As a general rule, award yourself a half mark, up to a maximum of three marks
in total, for each point you have successfully made (NB this is a general rule only, there
are no half marks awarded at Higher):
•
Propane molecules are held together by weak intermolecular forces / ethanol
molecules are held together by strong intermolecular forces.
•
The only intermolecular forces in propane are London Dispersion forces.
•
These are weak forces are due to momentary displacement of electrons between
atoms creating temporary dipoles.
•
In response to these temporary dipoles an induced dipole can occur.
•
The intermolecular forces in ethanol are hydrogen bonds.
•
Hydrogen bonding arises because the O-H bond is highly polar (there is a large
difference in the electronegativities of O and H).
•
The small positive dipole on H and small negative dipole on O strongly attract.
•
This causes ethanol to have a higher boiling point than propane as more energy
is needed to overcome the stronger hydrogen bonds than the weak London
Dispersion forces.
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