SCHOLAR Study Guide Higher Chemistry Unit 1: Chemical ... › 2020 › 04 › scholar-unit-1 … · 5, Unit 1); • all matter is made of atoms - when a substance contains only one

SCHOLAR Study Guide

Higher ChemistryUnit 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 2018 by Heriot-Watt University.

This edition published in 2018 by Heriot-Watt University SCHOLAR.

Copyright © 2018 SCHOLAR Forum.

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Distributed by the SCHOLAR Forum.

SCHOLAR Study Guide Higher Chemistry: Unit 1

Higher Chemistry Course Code: C813 76

ISBN 978-1-911057-38-3

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AcknowledgementsThanks are due to the members of Heriot-Watt University's SCHOLAR team who planned and created thesematerials, and to the many colleagues who reviewed the content.

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v

Contents

1 The Periodic Table 11.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 31.2 Arrangement of elements in the Periodic Table: Introduction . . . . . . . . . . . . 31.3 History of the Periodic Table . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 41.4 Trends and patterns (periodicity) . . . . . . . . . . . . . . . . . . . . . . . . . . . . 91.5 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 131.6 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 131.7 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14

2 Bonding and structure 172.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 202.2 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 212.3 Metallic bonds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 252.4 Covalent bonds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 272.5 Ionic bonds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 372.6 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 482.7 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 482.8 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 49

3 Periodic Table trends 533.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 553.2 Covalent radius . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 563.3 Ionisation energies . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 593.4 Electronegativity . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 643.5 Summary of trends in the Periodic Table . . . . . . . . . . . . . . . . . . . . . . . 673.6 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 683.7 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 683.8 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 69

4 Bonding continuum and polar covalent bonding 754.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 774.2 Polar covalent bonds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 784.3 Predicting bonding type using electronegativity . . . . . . . . . . . . . . . . . . . . 814.4 Polar molecules . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 834.5 The bonding continuum . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 884.6 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 904.7 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 904.8 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 91

5 Intermolecular forces 955.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 98

vi CONTENTS

5.2 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 985.3 London dispersion forces . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1005.4 Hydrogen bonding . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1045.5 Relating properties to bonding . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1075.6 Viscosity . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1105.7 Predicting solubilities from solute and solvent polarities . . . . . . . . . . . . . . . 1115.8 The solubility of flavour molecules . . . . . . . . . . . . . . . . . . . . . . . . . . . 1155.9 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1165.10 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1185.11 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 118

6 Oxidising or reducing agents 1216.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1246.2 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1256.3 Elements as oxidising and reducing agents . . . . . . . . . . . . . . . . . . . . . . 1266.4 Molecules and group ions as oxidising and reducing agents . . . . . . . . . . . . 1306.5 Uses for strong oxidising agents . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1336.6 Ion-electron half equations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1366.7 Combining ion-electron equations . . . . . . . . . . . . . . . . . . . . . . . . . . . 1396.8 Complex ion-electron equations . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1416.9 Summary exercise . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1456.10 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1476.11 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1486.12 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 149

7 End of unit test 153

Glossary 162

Answers to questions and activities 165

© HERIOT-WATT UNIVERSITY

1

Topic 1

The Periodic Table

Contents1.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3

1.2 Arrangement of elements in the Periodic Table: Introduction . . . . . . . . . . . . . . . . . . . 3

1.3 History of the Periodic Table . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4

1.4 Trends and patterns (periodicity) . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 9

1.4.1 Melting and boiling points . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 9

1.4.2 Atomic size . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11

1.5 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13

1.6 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13

1.7 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14

Prerequisites

Before you begin this topic, you should already know that:

• atoms contain protons, neutrons and electrons, each with a specific charge, mass andposition within the atom - the number of protons defines an element and is known asthe atomic number (National 4, Unit 1);

• or have knowledge of: sub-atomic particles, their charge, mass and position within theatom, the structure of the Periodic Table, groups, periods and atomic number (National5, Unit 1);

• all matter is made of atoms - when a substance contains only one kind of atom it isknown 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);

• there are seven diatomic elements: hydrogen, nitrogen, oxygen, fluorine, chlorine,bromine and iodine (National 5, Unit 1);

• elements can be categorised as metals and non-metals (National 4, Unit 1);

2 TOPIC 1. THE PERIODIC TABLE

Learning objective

By the end of this topic, you should be able to:

• state that elements are arranged in the Periodic Table in order of increasing atomicnumber;

• explain that the Periodic Table allows chemists to make accurate predictions of physicalproperties and chemical behaviour for any element based on its position;

• features of the table are Groups, vertical columns within the table contain elementswith similar chemical properties resulting from a common number of electrons in theouter shell, and Periods, rows of elements arranged with increasing atomic number,demonstrating an increasing number of outer electrons and a move from metallic tonon-metallic characteristics;

• state that there are periodic variations in the densities, melting points and boiling pointsof the elements across a Period and down a Group.


TOPIC 1. THE PERIODIC TABLE 3

1.1 Prior knowledge

Go onlineTest your prior knowledge

Q1: What does the number of protons in an atom determine?

a) Atomic numberb) Mass numberc) Element named) Atom charge

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Q2: Name the seven diatomic elements.

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Q3: Complete the following table.

Particle Charge Mass Location

Proton

Neutron

Electron

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Q4: What is a row in the Periodic Table is called?

a) Groupb) Periodc) Rowd) Set

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Q5: What is a column in the Periodic Table is called?

a) Columnb) Groupc) Periodd) Set

1.2 Arrangement of elements in the Periodic Table: Introduction

As more and more elements were discovered and their properties investigated, chemists showed anatural desire to simplify the study of the elements by organising them according to similarities intheir chemical behaviour. Eventually this resulted in the modern Periodic Table. If you understandhow the Periodic Table is constructed, you will realise that it contains a huge amount of informationstored in a very compact form. This makes it a vital resource for any chemist.



1.3 History of the Periodic Table

Background Information

In 1800, about thirty-three elements were known but there was no obvious pattern or relationshipbetween them.

By 1830, a further twenty or so elements had been discovered and some similarities in propertieswithin small groups of elements was recognised. The German chemist Johan Wolfgang Dobereinermade a tentative connection between chemical behaviour and the atomic masses of certain groupsof elements, each containing three elements, which he called 'triads'.

Figure 1.1: Triads

Triads Relative atomic massLi Na K 7 23 39S Se Te 32 79 128Cl Br I 35·5 80 127Ca Sr Ba 40 88 137

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'.

Go onlineNewlands' octaves

Figure 1.2 shows how Newlands organised the first fourteen elements. Study the informationcarefully and then answer the questions which follow.

Figure 1.2

Q6: What property did Newlands use to put the elements in order?

a) Name (i.e. alphabetical order)b) Colourc) Atomic massd) Atomic number

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Q7: What property is shared by the elements in the first column (H and F)?

a) Both are gases.b) Both are green.c) Both are reactive metals.d) None - they are completely different.

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Q8: What property is shared by the elements in the second column (Li and Na)?

a) Both are gases.b) Both are purple.c) Both are reactive metals.d) None - they are completely different.

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Q9: Why do you think Newlands referred to these as 'octaves'? (Hint: Think musical scales.)

Now look at the second arrangement in which the next seven elements have been added.

Figure 1.3

Q10: At first sight, the pattern seems to continue. For how many of the further elementsdoes the pattern work?

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Q11: Can you think of a reason why chromium, manganese and iron do not fit with theelements immediately above them?

Key point

When Newlands arranged the elements in order of increasing atomic mass, similar chemicalproperties were repeated with every eighth element, but this only worked for the firstseventeen elements.



Newlands' ideas were subjected to ridicule and it was even suggested unkindly that he would getbetter agreement if he arranged the elements in alphabetical order. However, he was on the rightlines.

The Modern Periodic Table

The major breakthrough, indeed one of the most important advances in all chemistry, was providedby Dmitri Ivanovitch Mendeleev in 1869. Like Newlands, he organised the elements in order ofincreasing atomic mass, but there were other important differences.

Go onlineMendeleev's Periodic Table

Mendeleev organised the first forty five of the sixty two elements known at that time.

Figure 1.4: Constructing Mendeleev's table

To maintain the chemical pattern, Mendeleev left a space (*) for the next element.

To ensure bromine appeared directly below chlorine, Mendeleev now left two spaces.



Iodine has a lower atomic mass than terullium but is chemically similar to bromine.

This is the table that Mendeleev proposed.

Q12: Unlike Newlands, Mendeleev left spaces (shown as *) in his table. Why?

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Q13: Explain why Mendeleev swapped the positions of the elements iodine and tellurium.

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Q14: Name the Group of elements which is found in the modern Periodic Table but is absentfrom both Newlands' and Mendeleev's tables.

Key point

Mendeleev organised the elements in order of increasing atomic mass (mostly) in conjunctionwith similar chemical properties, leaving gaps for elements yet to be discovered.



Go onlineThe modern Periodic Table

The relationship between the modern Periodic Table and that of Mendeleev can be seen.

Figure 1.5

Mendeleev was so confident that his table was correct that he predicted the properties ofsome of the undiscovered elements. His predictions proved to be very accurate when theseelements 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 theone used in the SQA Higher Chemistry Data Booklet. An interactive Periodic Table can befound at http://www.webelements.com/.

A good resource for finding out more about the elements in the Periodic Table can be foundat http://www.periodicvideos.com/. This website is from the University of Nottingham. It hasshort, informative video clips on each of the elements and is updated regularly.

Q15: In the modern Periodic Table, the elements are not arranged in order of increasingatomic mass. What is used instead?


http://www.webelements.com/

http://www.periodicvideos.com/


1.4 Trends and patterns (periodicity)

When the elements are arranged in order of increasing atomic number, many properties vary in aregular way. As you move across a Period from left to right, a pattern emerges. A similar patternappears on crossing the next Period. Properties which behave in this way are said to be periodic.

Throughout this topic, we will concentrate on the properties of the main group elements and, inparticular, 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.

Groups in the Periodic table are vertical columns within the table containing elements with similarchemical properties which result from a common number of electrons in the outer shell. Periods arerows of elements arranged with increasing atomic number, demonstrating an increasing number ofouter electrons and a move from metallic to non-metallic characteristics.

1.4.1 Melting and boiling points

Go onlineMelting and boiling points across a Period

Figure 1.6: Period 2 and 3 (melting point and boiling points)

Melting points and boiling points are also periodic properties and depend on the strength ofthe 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 thesolid state?

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Q17: Which of the Period 3 elements is the easiest to boil?

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Q18: Of all the elements shown in the SQA Higher Chemistry Data Booklet which elementwhich has the weakest forces between its atoms?

Go onlineMelting and boiling points down a Group

A more obvious trend can be seen on descending a Group.

Figure 1.7: Alkali metals (melting point and boiling points)

Q19: Which alkali metal has the strongest forces between its atoms?

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Q20: Which alkali metal has the weakest forces between its atoms?

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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 ofthe first twenty elements is considered later in this topic.



Key point

There are periodic variations in the densities, melting points and boiling points of the elementsacross a Period and down a Group.

1.4.2 Atomic size

The size of an atom is determined by the amount of space taken up by the electrons and so mustbe connected to the electron arrangement of the atom. The electron arrangement is itself a periodicproperty. Covalent radii information can be found on Page 7 of the SQA Higher Chemistry DataBooklet. Use Figure 1.8 in the following activity to explain the trends in covalent radius.

Go onlineAtomic size

Figure 1.8: Trends in nuclear charge and electron arrangement



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Q22: Which of the following provides the best reason for the increase in covalent radius ongoing down a Group?

a) The number of protons increases.b) The number of electrons increases.c) The number of electron shells increases.d) The number of neutrons increases.

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Q23: Which of the following provides the best reason for the decrease in covalent radius ongoing from left to right across a Period?

a) The number of electrons increases.b) The number of electron shells increases.c) The number of neutrons increases.d) The number of protons increases.

Key point

• The covalent radius decreases across a Period because the increase in nuclear chargeattracts the electrons more strongly.

• The covalent radius increases on going down a Group as the number of occupiedelectron shells increases.



1.5 Summary

Summary

• Elements are arranged in the Periodic Table in order of increasing atomic number.

• The Periodic Table allows chemists to make accurate predictions of physical propertiesand chemical behaviour for any element based on its position.

• Features of the table are Groups, vertical columns within the table contain elementswith similar chemical properties resulting from a common number of electrons in theouter shell, and Periods, rows of elements arranged with increasing atomic number,demonstrating an increasing number of outer electrons and a move from metallic tonon-metallic characteristics.

• There are periodic variations in the densities, melting points and boiling points of theelements across a Period and down a Group.

1.6 Resources

Texts

• SQA Higher Chemistry Data Book :https://www.sqa.org.uk/sqa/files_ccc/ChemistryDataBooklet_NewH_AH-Sep2016.pdf

• Higher Chemistry for CfE with Answers: Eric Allan, John Harris, John Anderson, HodderGibson ISBN 1444167529

• How to Pass Higher Chemistry for CfE : John Anderson, Hodder Gibson, ISBN 1471808289

• CfE Higher Chemistry (Bright Red Study Guide): Archie Gibb and David Hawley WilliamBeveridge, Bright Red Publishing ISBN 1906736596

Other resources:

• http://www.webelements.com/A useful website with an interactive Periodic Table.

• http://www.periodicvideos.com/A good resource for finding out more about the elements in the periodic table from theUniversity of Nottingham.


https://www.sqa.org.uk/sqa/files_ccc/ChemistryDataBooklet_NewH_AH-Sep2016.pdf

http://www.webelements.com/

http://www.periodicvideos.com/


1.7 End of topic test

Go onlineEnd of Topic 1 test

Q24: Menedeleev is famous for producing the Periodic Table on which the modern version isbased.

Which of the following statements is true?

a) Mendeleev organised the elements in order of their atomic number.b) Mendeleev left gaps because some elements did not fit the pattern of reactivity.c) Mendeleev left gaps for elements which had not yet been discovered.d) Mendeleev swapped some elements round so that their atomic masses fitted the pattern.

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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 leftto right.

f) There is an increase in electronegativity on going down Group 7.

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Q26: What does the number of protons in an atom determine?

a) Atomic numberb) Mass numberc) Element named) Atom charge

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Q27: Name the seven diatomic elements.

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Q28: What is Group 1 in the Periodic Table called?

a) The halogens.b) The alkali metals.c) The noble gases.d) Transition metals.

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Q31: Elements in the same group in the Periodic Table have the same:

a) number of occupied energy shells.b) density.c) number of outer electrons.d) number of protons.

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Q32: Where in the Periodic Table are non-metals found?

a) The top.b) The bottom.c) The right hand side.d) The left hand side.


17

Topic 2

Bonding and structure

Contents2.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 202.2 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 212.3 Metallic bonds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 25

2.3.1 Metallic structures . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 262.4 Covalent bonds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 27

2.4.1 Covalent molecular structures . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 292.4.2 Covalent network structures . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 33

2.5 Ionic bonds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 372.5.1 Ionic lattice structures . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 392.5.2 Ionic formulae . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 41

2.6 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 482.7 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 482.8 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 49

Prerequisites


• 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 asubstance (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 theirouter electrons (National 4, Unit 1);

• covalent molecular compounds have low melting and boiling points - as a result, theycan 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 twobonded atoms (National 5, Unit 1);

• more than one bond can be formed between atoms leading to double and triple covalentbonds (National 5, Unit 1);

18 TOPIC 2. BONDING AND STRUCTURE

Prerequisites continued

• 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 amolecule and the shape of simple two-element compounds (National 5, Unit 1);

• covalent molecular substances have low melting and boiling points due to only weakforces of attraction between molecules being broken (National 5, Unit 1);

• giant covalent network structures have very high melting and boiling points because thenetwork 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 particleis known as an ion (National 5, Unit 1);

• ionic bonds are the electrostatic attraction between positive and negative ions - ioniccompounds form lattice structures of oppositely charged ions (National 5, Unit 1);

• ionic formulae can be written giving the simplest ratio of each type of ion in thesubstance (National 5, Unit 1);

• ionic bonds are the electrostatic attraction between positive and negative ions (National5, Unit 1);

• 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 mustbe broken in order to break down the lattice - dissolving also breaks down the latticestructure (National 5, Unit 1);

• ionic compounds have high melting and boiling points - as a result, they are found in thesolid state at room temperature (National 4, Unit 1);

• ionic compounds form when metal atoms join to non-metal atoms by transferringelectron(s) from the metal to the non-metal - the resulting charged particles are calledions - 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 thebreakdown of the lattice resulting in the ions being free to move (National 5, Unit 1).


TOPIC 2. BONDING AND STRUCTURE 19

Learning objective


• state that the first twenty elements in the Periodic Table can be categorised accordingto bonding and structure:

◦ metallic Li, Be, Na, Mg, Al, K, Ca;

◦ covalent molecular H2, N2, O2, F2, Cl2, P4, S8 and fullerenes (eg C60);

◦ covalent network B, C, Si (diamond, graphite, the element silicon and thecompound silicon dioxide);

◦ monatomic He, Ne, Ar (noble gases);

• explain each of these types of substance in terms of bonding and structure.



2.1 Prior knowledge


Q1: Covalent bonding involves:

a) a shared pair of electrons.b) transfer of electrons.c) delocalised electrons.d) gaining electrons.

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Q2: Metals can conduct electricity in their solid state because of their:

a) ions.b) positive cores.c) lattice structure.d) delocalised electrons.

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Q3: Ionic bonding involves:

a) a shared pair of electrons.b) transfer of electrons.c) delocalised electrons.d) protons only.

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Q4: One property of covalent networks is that they:

a) have high melting and boiling points.b) have low melting and boiling points.c) are soluble in water.d) conduct electricity when molten or in solution.

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Q5: One property of ionic substances is that they:

a) are all white in colour.b) have low melting and boiling points.c) are insoluble in water.d) conduct electricity when molten or in solution.



2.2 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, Fe2O3). Some arecomplex minerals (e.g. columbite, containing tantalum, (Fe,Mn)(Ta,Nb)2O6).

Haematite (http:/ / bit.ly/ 2aK AD0W by http:/ / commons.wikimedia.org/wiki/ User :Sailko, islicensed under http:/ / creativecommons.org/ l icenses/ by/ 2.0/ deed.en)

Columbite is a black mineral group that is an ore of niobium (http:/ / commons.wikimedia.org/wiki/ F i le:Columbite-75444.jpg, by http:/ / www.irocks.com, is licensed under http:/ / creativeco

mmons.org/ l icenses/ by/ 2.0/ deed.en)


http://bit.ly/2aKAD0W

http://commons.wikimedia.org/wiki/User:Sailko

http://creativecommons.org/licenses/by/2.0/deed.en

http://commons.wikimedia.org/wiki/File:Columbite-75444.jpg

http://commons.wikimedia.org/wiki/File:Columbite-75444.jpg

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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 oftantalum 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 isfor?

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 haveto 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 choicedepends 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 chemicallybonded in various ways to produce a multitude of different molecules.

The different types of chemical bonding determine the structure that elements and compoundsadopt. In turn the structure, together with the size of intermolecular forces mainly determines thephysical and chemical properties possessed by these materials.

Bonding and structure will be studied in this topic. Intermolecular forces (the forces that existbetween molecules) and properties are studied in a later topic.

Examples of some substances with different properties, depending on different bonding andstructures are shown in the following images.

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 allinteract 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 ofincreasing energy levels, around it.

The following activity shows the final electronic structure of a sodium atom.



Go onlineConstruction of a sodium atom

Construction of a sodium atom

In the case of the atoms of noble gases, all shells are completely filled with electrons. Thisarrangement is particularly stable, which makes noble gases unreactive.

Atoms of other elements combine in ways which try to achieve this stable noble gasarrangement of electrons, in other words, to become isoelectronic with a noble gas.

Q8: Look at the Periodic Table on Page 8 of the SQA Higher Chemistry Data Booklet. Whichnoble gas has an electronic structure closest to chlorine?

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Q9: What must happen to a chlorine atom to get the noble gas electronic arrangement?

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Q10: What must happen to the sodium atom for it to have a noble gas electronic structure?

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Q11: What will the sodium atom become?

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Q12: Which noble gas has the same electronic structure as a sodium ion?



2.3 Metallic bonds

Metals occur to the left and centre of the Periodic Table (see Topic 1). When atoms of metals cometogether, the most stable condition is for them to 'pool' their outer electrons and become, in effect, aregular 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 willachieve the stable sodium ion structure (isoelectronic with neon).

Go onlineMetallic bonds

The following diagram shows the difference in structure between gaseous sodium andmetallic sodium.

Such an array of positively charged ions would normally split apart by mutual repulsions, butthe influence of the negative electrons holds the particles in the structure so well that mostmetals 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 normalconditions?



2.3.1 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 structureknown as a lattice.

Part of the structure of the giant lattice for copper is shown in the following activity.

Go onlineCopper lattice

Q14: The electrons in a metallic bond are said to be:

a) loose.b) ionised.c) delocalised.d) hard.

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Q15: What do the ions in a metallic bond form into?

a) Gelb) Gridc) Bard) Lattice

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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 anddelocalised outer electrons.

A metallic structure consists of a giant lattice of positively charged ions and delocalised outerelectrons.



2.4 Covalent bonds

You should already have studied covalent bonds, which are the most common type of bond. Thissection revises your knowledge.

Covalent bonds form when atoms share two electrons, enabling both atoms to complete theirvalency shells.

The following activity illustrates two separate chlorine atoms and the sharing of electrons to form achlorine molecule.

Go onlineCovalent bond formation

This sharing of two electrons (one from each atom) to complete an outer shell of electrons is calleda 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 positivelycharged nuclei will repel each other, as will the negatively charged electrons; but these forces aremore than balanced by the attraction between the nuclei and electrons.

Attractive and repulsive forces in a hydrogen molecule

When atoms require more than one electron to complete their outer shell, they can share two orthree 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 (O2) and ammonia (NH3). Noticethe 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 preceding diagram are the bonding electrons, shared between theatoms. 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?

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Q18: How many lone pairs are there in an ammonia molecule?

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Q19: Which two words complete this sentence?

A covalent bond is formed when two atoms �� a pair of electrons, so that each can achievea noble gas �� configuration.

a) share, protonb) share, electronc) donate, atomicd) donate, ionic

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Q20: The dominant attractive force in a covalent bond is between:

a) the positively charged nuclei.b) the negatively charged shared electrons.c) negatively charged nuclei and positively charged shared electrons.d) positively charged nuclei and negatively charged shared electrons.

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Q21: Draw a diagram showing how bonding occurs in iodine chloride. Look at Page 8 of theSQA Higher Chemistry Data Booklet for the electron structures of the atoms.

Key point

Atoms in a covalent bond are held together by electrostatic forces of attraction betweenpositively charged nuclei and negatively charged shared electrons.

2.4.1 Covalent molecular structures

Most covalent substances (for example all those mentioned in the previous section) exist as discretemolecules. There are strong covalent bonds binding the atoms together in the molecule, butmuch weaker forces between these molecules. You will study these forces in detail later (referto 'Intermolecular Forces' in Topic 5). Consequently, many small covalent molecules are gases (e.g.fluorine, oxygen, nitrogen, carbon dioxide, sulfur dioxide etc). Larger molecules make structureswhich are liquids or low melting point solids (think of candle wax, with molecular mass around 500,but a low melting point around 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.



Go onlineWhite phosphorus

White phosphorus

Phosphorus (in the common form of 'white' phosphorus) forms P4 molecules with eachphosphorus atom at the corner of a tetrahedron.

Tetrahedron of four atoms in white phosphorus



Go onlineRed phosphorus

Red phosphorus

Phosphorus can also form a much less reactive (and less toxic) form, 'red' phosphorus, whichis used in making matches. Its structure consists of chains of these tetrahedra. The structureis as follows.

These P4 tetrahedra are quite stable so that it is usual to write 'P4' in equations wherephosphorus is involved. It is even retained in some compounds - the oxides are molecules offormula P4O6 and P4O10.

Red phosphorus structure



Go onlineSulfur

Sulfur

Molecules of sulfur (in the normal solid state as rhombic sulfur) consist of puckered rings ofeight sulfur atoms, written S8.

Rings of eight sulfur atoms

In this case, using S8 in equations would be rather cumbersome so sulfur in equations iswritten '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 O2; the blue colour of burning sulfur is due to theemission 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 discussedin the next section after the other common forms of carbon.



2.4.2 Covalent network structures

In some elements and compounds the covalent bonds are not limited only to those within themolecules, 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 heldtogether by covalent bonds. Each carbon atom makes covalent bonds to four different carbonatoms, which each bond to three more carbons, and so on, binding the whole structure together asshown in the following activity.

This covalent network structure has the characteristic properties of hardness and a high meltingpoint, both due to the strong bonding throughout the structure.

Go onlineDiamond

Arrangement of carbon atoms in diamond

Q22: How many pairs of electrons are around each carbon atom?

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Q23: What type of bond holds the carbon atoms together in diamond?

a) Single covalent.b) Double covalent.c) Triple covalent.



The following activities allow you explore other covalent network structures.

Go onlineSilicon dioxide

Look at the following model, which compares the structure of silica (quartz, SiO2) withdiamond. Again, this consists of a network of covalent bonds, with each silicon atom bondedto four oxygen atoms which in turn bond to another silicon atom, to form a strongly-boundnetwork. Quartz is hard and has a high melting point, like diamond.

Silica and diamond

Go onlineGraphite structure

Carbon has another interesting structure, graphite. In this case, carbon atoms covalentlybond with only three other atoms to form a sheet of hexagonal rings.

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 bondstrength is:

a) weaker.b) the same.c) stronger.

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Q25: In these sheets, each carbon atom is bonded to only three others. How many bondingelectrons surround each carbon atom in this case?

Graphene

It is only recently that single layers of carbon atoms as found in graphite have been isolatedin sufficient quantities to study their properties. The Nobel Prize for physics for 2010 wasawarded 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 singlelayers have very unusual electrical and physical properties, which have started research intonew electronic applications.

Graphite

The carbon atoms have four electrons in their outer shell. Only three are used to bond withthree other atoms, leaving a fourth electron. These sheets are layered together and thecarbon atoms donate their extra electron to a delocalised pool which holds the sheets in ametallic type bonding.

Layers of carbon atoms in graphite

The distance between these sheets is 335 pm which indicates a weak bonding. Theseweaker forces produce a material which is greasy (as the layers slip over each other) andcan be used as a lubricant, and the delocalised electron pool makes graphite a goodconductor of electricity.



Go onlineFullerenes

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 forminga large sphere, looking like a football. It is called buckminsterfullerene after the architect(Robert Buckminster Fuller) who creates geodesic domes resembling the structure found inthese carbon molecules.

Buckminsterfullerene

These fullerenes have interesting properties, which are currently being researched. Forexample, the 'buckyball', as illustrated, will dissolve in benzene to form a red solution.

Q26: Why do you think that buckminsterfullerene is soluble when diamond and graphite arenot?

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. Whenvery many of these long molecules are combined, the result consists of very large moleculesof carbon atoms bonded to each other with strong bonds, similar to those in sheets ofgraphite. These bonds are much stronger than metallic bonds in steel.

Key point

A covalent molecular structure consists of discrete molecules held together by intermolecularforces (see next topic).

A covalent network structure consists of a giant lattice of covalently bonded atoms.

2.5 Ionic bonds

Polar covalent bonds (you will study 'Polar covalent bonding' in a later topic) are formed when theatoms involved in a bond have different attractions for the bonding electrons. When two atoms havea large difference in their attraction for electrons (e.g. sodium and chlorine), it is most energeticallyfavourable for an electron in the metal to be donated completely to the non-metal. This is theconcept of electronegativity and will be fully explained in the next topic. Both then achieve a noblegas structure as shown in the following activity.

Go onlineIonic bond formation



Q27: How many outer electrons does a sodium atom have?

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Q28: How many outer electrons does a chlorine atom have?

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Q29: When a sodium atom loses its outer electron to a chlorine atom, how many outerelectrons do both ions then have?

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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 becomemore polar, then ionic, as the difference in electronegativity increases.

The extent of covalent or ionic character depends mainly on the difference in electronegativity. Theproperties of the different bonds are summarised in the following table.

Bond Covalent Polar covalent Ionic

Electrons Equally shared Unequally shared Totally transferred

ChargeDistribution

None Partial +ve and -veFully 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 astable noble gas structure. This is reflected in the relatively low first and second ionisationenergies. These elements will require two chlorine atoms to receive these two electrons andmaintain 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-metalions 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 ina later section.



Q31: Ionic bonds are formed between atoms which have a �� difference inelectronegativity.

a) zerob) smallc) larged) 0.2

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Q32: Caesium chloride is an ionic solid consisting of:

a) negative caesium ions and negative chloride ions.b) negative caesium ions and positive chloride ions.c) positive caesium ions and negative chloride ions.d) 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 (M2+ : X3-)in the solid?

a) 1 : 1b) 2 : 3c) 3 : 2d) 3 : 3

2.5.1 Ionic lattice structures

The mutual electrostatic attraction of positive and negative ions is completely non directional (unlikecovalent bonds which are highly directed). When solid ionic compounds form, a positive ion willbe surrounded by several negative ions which, in turn, will attract more positive ions. This processresults in the formation of a lattice of regularly arranged ions, all held together by electrostaticforces. These forces act equally in all directions so that no one ion is attracted particularly to anyother specific ion. For example, in sodium chloride there are no molecules formed. The ionic latticefor sodium chloride is shown in the following activity.



Go onlineSodium chloride ionic lattice

Sodium chloride lattice

This extensively bonded ionic solid is relatively hard, though not as hard as the covalentnetwork solids, but is brittle. (Think of how easily salt (sodium chloride) crystals can becrushed.)

Rock salt cellar

Q34: What are caesium chloride (CsCl) crystals formed from?

a) Covalent molecules.b) A covalent lattice.c) An ionic lattice.d) Ionic molecules.

Now that we have looked at the three types of bonding in elements and compounds, see thefollowing table which roughly compares their strengths. The table also gives a figure for the weakerbonds 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 thinkof 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 holdcovalent molecules together. (These are the intermolecular forces, which are studied in the nexttopic.)

Key point

Ionic bonding is the electrostatic force of attraction between positively and negatively chargedions.

An ionic structure consists of a giant lattice of oppositely charged ions.

2.5.2 Ionic formulae

We can write formulae for ionic substances using the same method as for covalent substances.

Ionic formulae give the simplest ratio of each type of ion in the substance and can show thecharges on each ion, if required.

In ionic formulae, charges must be superscript and numbers of atoms/ions must be subscript.

2.5.2.1 Formulae Involving Group ions

In National 5, we learned that some ions contain more than one type of atom and are referred to asgroup ions.

Some common group ions are shown in the following table.

One positive One negative Two negative Three negativeIon Formula Ion Formula Ion Formula Ion Formula

ammonium NH4+

ethanoatehydrogencarbonate

hydrogensulfatehydrogensulfite

hydroxidenitrate

permanganate

CH3COO-

HCO3-

HSO4-

HSO3-

OH-

NO3-

MnO4-

carbonatechromate

dichromatesulfatesulfite

thiosulfate

CO32-

CrO42-

Cr2O72-

SO42-

SO32-

S2O32-

phosphate PO43-

Examples of group ions can be found on page 21 of the SQA Higher Data Book.



We work out the formula for compounds containing group ions in the same way as before.

The valency of a group ion is the same as the value of its charge.

For example, the nitrate ion (NO3-) has a one negative charge and so its valency is one.

The nitrate ion (PO43-) has a three negative charge and so its valency is three.

Common group ions: Question

Q35: Complete the following table to show the valency of some common group ions.

Group ion name Group ion formula Valency

ammonium NH4+

carbonate CO32-

sulfate SO42-

phosphate PO43-

ethanoate CH3COO-

hydrogencarbonate HCO3-

Examples

1. Calcium nitrate

1. Write out the atomic symbol for calcium and the formula for the nitrate ion.

Ca NO3

Remember that the group ion formulae can be found on page 8 of the data book.

2. Write the valency underneath each substance.

Ca NO3

2 1

Remember, valencies for group ions are the same as the value of their charge.

3. Swap the valencies over.

Ca NO3

1 2

4. Check if you can simplify the valencies by dividing by a common factor.

Ca NO3

1 2

Here, we cannot simplify.

5. Write out the formula.



Ca NO3

1 2

Ca(NO3)2Remember, we do not show the number '1' when we write out formula.

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2. Magnesium sulfate

1. Write out the atomic symbol for magnesium and the formula for the sulfate ion.

Mg SO4



Mg SO4

2 2



Mg SO4

2 2


Mg SO4

1 1

Here, we were able to divide both valencies by 2.


Mg SO4

1 1

MgSO4

Remember, we do not show the number '1' when we write out formula.

Using brackets in formula

When there is more than one of the same group ion in a formula we put brackets around that groupion and use a subscript outside the brackets to show how many of those ions there are present.

Everything inside the bracket is multiplied by the subscript outside the bracket.

For example, the formula for calcium nitrate in the above example is represented as Ca(NO3)2 ratherthan CaNO32 which would have an entirely different meaning!



Example : Ammonium carbonate

1. Write out the formulas for ammonium and the carbonate ion.

NH4 CO3



NH4 CO3

1 2



NH4 CO3

2 1


NH4 CO3

2 1



NH4 CO3

2 1

(NH4)2 (CO3)

Remember, we do not show the number '1' when we write out formula.

Go onlineGroup ion formulae: Questions

What are the formulae for the following compounds containing group ions?

Q36: Copper (II) sulfate

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Q37: Sodium hydroxide

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Q38: Barium hydrogensulfate

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Q39: Beryllium permanganate



2.5.2.2 Writing ionic formulae

We have seen that when forming ionic compounds, metals in Groups 1, 2 and 3 lose their outerelectrons to obtain the stable electron arrangement of a noble gas.

Non-metals in Groups 5, 6 and 7 gain electrons to obtain the stable electron arrangement of a noblegas.

In Group 4, the non-metals carbon and silicon do not form ions based on single atoms.

The charges on the ions of main group elements can be summarised as follows.

Groupnumber

1 2 3 4 5 6 7

Chargeon ion

1+ 2+ 3+ NA 3- 2- 1-

Ionic formulae is a formulae where we show the ionic charges within the formula.

If there is more than one of either of the ions in the ionic formulae then we use brackets as in theexample below.

The positive and negative charges shown on the ions in an ionic formula should always balance.

Examples

1. Calcium chloride

1. Write out the ionic formulae for the calcium and chloride ions.

Ca2+ Cl−

Remember that the charge on a main group ion is linked to its group number in thePeriodic Table.


Ca2+ Cl−

2 1

Remember, valencies for main group elements are linked to their group in the PeriodicTable.


Ca2+ Cl−

1 2


Ca2+ Cl−

1 2





Ca2+ Cl−

1 2(Ca2+

) (Cl−

)2

Remember, if there is more than one of either of the ions in the ionic formulae then weuse brackets.

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2. Copper (II) nitrate

We have already seen how a Roman Numeral can be used to show the valency of a metalatom.

The Roman numeral also gives the size of the charge on the ion that is present.

1. Write out the atomic symbol for copper and the formula for the nitrate ion.

Cu2+ NO3−



Cu2+ NO3−

2 1



Cu2+ NO3−

1 2


Cu2+ NO3−

1 2



Cu2+ NO3−

1 2(Cu2+

) (NO3

−)2

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3. Ammonium carbonate

1. Write out the formula for the ammonium and carbonate ions.

NH4+ CO3

2−





NH4+ CO3

2−

1 2



NH4+ CO3

2−

2 1


NH4+ CO3

2−

2 1



NH4+ CO3

2−

2 1(NH4

+) (

CO32−

)2

Go onlineGroup ion ionic formulae: Questions

What are the ionic formulae for the following compounds containing group ions?

Q40: Nickel (II) sulfate

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Q41: Potassium permanganate

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Q42: Barium hydrogencarbonate

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Q43: Ammonium phosphate

Key point

Ionic formulae

• ions containing more than one type of atom are often referred to as group ions;

• chemical formulae can be written for compounds containing group ions using valencyrules and the data booklet;

• ionic formulae give the simplest ratio of each type of ion in the substance and can showthe charges on each ion, if required;

• in formulae, charges must be superscript and numbers of atoms/ions must be subscript.



2.6 Summary

Summary

• Metallic bonding is the electrostatic force of attraction between positively charged ionsand delocalised outer electrons.

• A metallic structure consists of a giant lattice of positively charged ions and delocalisedouter electrons.

• Atoms in a covalent bond are held together by electrostatic forces of attraction betweenpositively charged nuclei and negatively charged shared electrons.

• A covalent molecular structure consists of discrete molecules held together byintermolecular 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 negativelycharged ions.

• An ionic structure consists of a giant lattice of oppositely charged ions.

• Ionic formulae can be written giving the simplest ratio of each type of ion in thesubstance.

• Ionic bonds are the electrostatic attraction between positive and negative ions.

• Ionic compounds form lattice structures 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

2.7 Resources

Texts










Q44: In which of the following compounds do both ions have the same electron arrangementas argon?

a) Calcium sulfideb) Magnesium oxidec) Sodium sulfided) Calcium bromide

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Q45: Which element is a solid at room temperature and consists of discrete molecules?

a) Sulfurb) Neonc) Silicond) Boron

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Q46: Graphite, a form of carbon, conducts electricity because it has:

a) pure covalent bonding.b) delocalised electrons.c) metallic bonding.d) van der Waals' bonding.

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Q47: Which of the following can be applied to lithium but not to carbon?

a) Covalentb) Metallicc) Made up of discrete molecules.d) Made up of diatomic molecules.e) Gasf) Solid

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Q48: 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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Q49: Identify the element which exists as a covalent network solid.

a) Boronb) Chlorinec) Nitrogend) Phosphoruse) Sodiumf) Sulfur

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Q50: Identify the two elements which exist as discrete covalent molecular solids.

a) Boron

b) Chlorine

c) Nitrogen

d) Phosphorus

e) Sodium

f) Sulfur

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Q51: 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

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Q52: Which of the elements is most likely to have a covalent network structure?

ElementMelting Point(K)

Boiling Point(K)

Density (gcm-3)

Conductivewhen 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

a) Ab) Bc) Cd) D

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Q53: Which type of structure is found in a substance melting at 1069◦C which conductselectricity when molten, but not when solid?

a) Ionicb) Covalent (network structure)c) Metallicd) Covalent (discrete molecules)

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Q54: Identify the covalent network substance.

a) NH4Cl (s)b) CH3OH (l)c) C6H14 (l)d) SO2 (g)e) Na2CO3 (s)f) SiO2 (s)

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Q55: Identify the two substances which are ionic.

a) NH4Cl (s)

b) CH3OH (l)

c) C6H14 (l)

d) SO2 (g)

e) Na2CO3 (s)

f) SiO2 (s)

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Q56: Which of the following structures is never found in compounds?

a) Ionicb) Monatomicc) Covalent Moleculard) Covalent Network



Go onlineSQA style questions

Q57: Diamond and graphite are forms of carbon with very different properties. Graphitecan mark paper, is a lubricant and is a conductor of electricity. Diamond has none of theseproperties.

Draw a diagram to show the structure of diamond.

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Q58: Why is graphite an effective lubricant?

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Q59: Boron nitride can form a similar structure to graphite. The boron and nitrogen atomsalternate throughout the structure as shown.

Why is this substance a non-conductor, while graphite conducts electricity?

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Q60: Suggest why the bonds between the layers in boron nitride are stronger than the bondsbetween the layers in graphite.


53

Topic 3

Periodic Table trends

Contents3.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 55

3.2 Covalent radius . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 56

3.2.1 The trends in covalent radius . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 57

3.3 Ionisation energies . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 59

3.3.1 Explanation of the trends in first ionisation energy . . . . . . . . . . . . . . . . . . . . 62

3.4 Electronegativity . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 64

3.5 Summary of trends in the Periodic Table . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 67

3.6 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 68

3.7 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 68

3.8 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 69

Prerequisites


• atoms contain protons, neutrons and electrons, each with a specific charge, mass andposition within the atom - the number of protons defines an element and is known asthe atomic number (National 4, Unit 1);

• or have knowledge of: sub-atomic particles, including their charge, mass and positionwithin the atom, and the structure of the Periodic Table, including Groups, Periods andatomic number (National 5, Unit 1);

• all matter is made of atoms - when a substance contains only one kind of atom it isknown 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);



54 TOPIC 3. PERIODIC TABLE TRENDS



• more than one bond can be formed between atoms, leading to double and triple covalentbonds (National 5, Unit 1);



• when there is an imbalance in the number of positive protons and electrons, the particleis known as an ion (National 5, Unit 1).

Learning objective


• 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 explainedin terms of the number of occupied shells, and the nuclear charge;

• state that trends in ionisation energies across Periods and down Groups can beexplained in terms of the:

◦ atomic size;

◦ nuclear charge;

◦ 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 hasfor the electrons of the bond;

• state that electronegativity values increase across a Period and decrease down aGroup;

• state that electronegativity trends can be rationalised in terms of nuclear charge,covalent radius and the presence of 'screening' inner electrons.


TOPIC 3. PERIODIC TABLE TRENDS 55

3.1 Prior knowledge


Q1: Complete the following table.


Proton

Neutron

Electron

Q2: Elements in the same Group in the Periodic Table have the same:


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Q3: Elements in the same Period in the Periodic Table have the same:


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Q4: What determines the number of protons in an atom?

a) Atomic number.b) Mass number.c) Name of element.d) Charge of atom.

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a) a shared pair of electrons.b) transfer of electrons.c) delocalised electrons.d) gaining electrons.



3.2 Covalent radius

The size of individual atoms is difficult to measure since the size is determined by the space taken upby the constantly moving electrons. However, the distance between the nuclei of atoms in the solidstate can be measured by a technique called X-ray diffraction. Consequently, the size of an atom isusually described in terms of its covalent radius. This is defined as half the distance between thenuclei of two bonded atoms of the element (see Figure 3.1).

Figure 3.1: Definition of covalent radius

The distance between the nuclei is shown as 2r and so the covalent radius in each case is r.Covalent radii information can be found on Page 7 of the SQA Higher Chemistry Data Booklet.

Go onlineCovalent radius - relative sizes

The covalent radius is another periodic property.

Figure 3.2: The covalent radius



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Q6: What happens to the size of the atoms on crossing a Period from left to right?

a) An increaseb) A decreasec) Nothing

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Q7: What happens to the size of the atoms on descending a group?

a) An increaseb) A decreasec) Nothing

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Q8: Can you suggest why are there no covalent radii quoted for the noble gases (Group 0)?

3.2.1 The trends in covalent radius

The size of an atom is determined by the amount of space taken up by the electrons and so mustbe connected to the electron arrangement of the atom. The electron arrangement is itself a periodicproperty.

Use the following activity to explain the trends in covalent radius.

Go onlineThe trends in covalent radius

Figure 3.3: The trends in covalent radius



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Q9: Which of the following provides the best reason for the increase in covalent radius ongoing down a group?

a) The number of protons increases.b) The number of electrons increases.c) The number of electron shells increases.d) The number of neutrons increases.

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Q10: Which of the following provides the best reason for the decrease in covalent radius ongoing from left to right across a Period?

a) The number of electrons increases.b) The number of electron shells increases.c) The number of neutrons increases.d) The number of protons increases.



Key point

The covalent radius decreases across a Period because the increase in nuclear chargeattracts the electrons more strongly.

The covalent radius increases on going down a group as the number of occupied electronshells increases.

3.3 Ionisation energies

The electrons orbiting the nucleus of an atom are held by the electrostatic attraction between thenegative electrons and the positive nucleus. Some atoms lose electrons relatively easily, whereassome lose electrons only with great difficulty. No atoms simply give away electrons; to remove anelectron always requires energy to overcome the force of attraction between the electron and thenucleus. 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 quotevalues 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 formulamass 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

Ionisation energy

The first ionisation energy (IE) is defined as the energy required to remove one mole of electronsfrom one mole of gaseous atoms (one electron from each atom). The units used are kilojoules permole (kJ mol-1).

Consider the element carbon. The value quoted for the first ionisation energy of carbon is1086 kJ mol-1. In other words, 1086 kJ of energy is required to remove one electron from each atomin one mole (12 g) of gaseous carbon.

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 3.4: First ionisation energy of carbon

The C+ ion is smaller than the C atom because the remaining five electrons in the ion still feel thefull attractive force from six protons and so they are more tightly held. As a result, it is even moredifficult to remove the next electron.

Figure 3.5: Second ionisation energy of carbon

The removal of one mole of electrons from one mole of gaseous C+ ions is known as the secondionisation energy of carbon. This has a value of 2353 kJ mole-1. The following figure showsequations representing the first four ionisation energies of carbon.

Figure 3.6: Successive ionisation energies



Go onlineIonisation energies

Q11: Using Page 11 of the SQA Higher Chemistry Data Booklet, plot a graph of first ionisationenergy against atomic number for the first twenty elements.

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Q12: Is the first ionisation energy a Periodic property?

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Q13: From the shape of the graph, explain your answer to the previous question.

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Q14: Name the group of elements which appears at the peaks in the graph.

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Q15: Name the group of elements which have the lowest ionisation energies.

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Q16: What is the general trend in first ionisation energy, going across a Period from left toright?

a) A steady increase.b) A steady decrease.c) An increase followed by a decrease.d) A decrease followed by an increase.

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Q17: What is the general trend in first ionisation energy, going down a Group?

a) A steady increase.b) A steady decrease.c) An increase followed by a decrease.d) A decrease followed by an increase.

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Q18: Confirm your answer to the previous question by listing the first ionisation energies ofthe alkali metals in order of atomic number.



3.3.1 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 willconcentrate 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 arethe attractive forces between the electron and the nucleus. The strength of this force of attractionwill 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 electronfrom 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 3.7).

Figure 3.7: 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.

Q19: Explain fully why the first ionisation energy decreases on going down Group 1. Youshould mention nuclear charge, covalent radius and screening effect.

(3 marks)



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 thispart.

Go onlineFirst ionisation energy - Going across a Period

Discuss the same points with a partner as before.

• 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 leftto right.

(3 marks)

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Q21: What happens to the nuclear charge on going across Period 2 from left to right?

a) It decreases.b) It increases.c) It stays the same.

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Q22: What effect will this have on the first ionisation energy on going across Period 2 fromleft to right?

a) It increases.b) It decreases.c) It stays the same

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Q23: What happens to the number of inner shell electrons on going across Period 2 from leftto right?


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Q24: What happens to the screening effect on going across Period 2 from left to right?


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Q25: As a result, what happens to the size of the atoms on going across Period 2 from left toright?


• Use page 11 of the SQA Higher Chemistry Data Booklet to calculate the energy requiredto 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 electronfrom 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 onemole of gaseous atoms. The second and subsequent ionisation energies refer to the energiesrequired to remove further moles of electrons.

First ionisation energies increase across a Period and decrease down a Group. This can beexplained in terms of atomic size, nuclear charge and the screening effect due to inner shellelectrons.

3.4 Electronegativity

The first ionisation energy involves the removal of electrons from gaseous atoms and so is ameasure of how strongly an isolated atom holds on to its outermost electrons. In the world aroundus, isolated atoms are very rare and the vast majority of atoms are found bonded to one or moreother atoms. How and why atoms bond together is the basis of chemistry and electrons are thefundamental particles involved in bonding.

In all theories of bonding, different types of bond arise because atoms of different elements havedifferent attractions for electrons. Atoms which tend to attract the electrons within a bond are saidto be electronegative.



An atom with a high electronegativity will tend to attract bonded electrons towards it whereas anatom with a low electronegativity will have a very weak attraction for electrons. There will be a 'tugof war' between the different atoms for the electrons.

Go onlineElectronegativity

In Figure 3.8, the darker shading shows the electrons being pulled towards the moreelectronegative atom.

Figure 3.8: Electronegativity

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Q26: If X is hydrogen, Y could be:

a) Berylliumb) Brominec) Boron

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Q27: The bond X-Y could be:

a) C-Clb) H-Sc) C-S

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Q28: If X is in Group 5, Y could be in:

a) Group 3b) Group 6c) Group 7

There are several different methods for estimating electronegativity values. The one most commonlyused is that devised by double Nobel Prize winner, Linus Pauling. He assigned a value to each ofthe elements most commonly found in bonds. Lithium was assigned a value of 1.0 while fluorinewas assigned the highest value of 4.0. The electronegativity values produced by Pauling are quotedin the SQA data booklet (page 11). Study these values carefully and look for any patterns.

Q29: What is the Group number of the elements which have the lowest electronegativities?

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Q30: What is the Group number of the elements with the highest electronegativities?

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Q31: What is the name of the Group for which no values are quoted?

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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 youmove from left to right, there are more protons and so the nuclear charge is increased. As you movedown a Group, atomic size is increasing so outer electrons are further from the positively chargednucleus. As a result, the electronegativity decreases.



Q33: Which of the following statements describes the trends in electronegativity values in thePeriodic Table?

a) Electronegativity values increase on going from left to right across a Period and increaseon going down a Group.

b) Electronegativity values decrease on going from left to right across a Period and decreaseon going down a Group.

c) Electronegativity values increase on going from left to right across a Period and decreaseon going down a Group.

d) Electronegativity values decrease on going from left to right across a Period and increaseon going down a Group.

Key point

Electronegativity is a measure of the attraction an atom in a bond has for the electrons of thebond. Electronegativity values increase across a Period and decrease down a Group.

3.5 Summary of trends in the Periodic Table

Go onlineSummary: Periodic Table trends

Q34: Complete the table by selecting the appropriate arrow symbol to indicate the increasingsize of each property as you move across a Period and up/down Group 1 and 7.



3.6 Summary

Summary

• The covalent radius decreases across a Period because the increase in nuclear chargeattracts the electrons more strongly.

• The covalent radius increases on going down a Group as the number of occupiedelectron shells increases.

• The first ionisation energy is the energy required to remove one mole of electrons fromone mole of gaseous atoms.

• The second and subsequent ionisation energies refer to the energies required to removefurther 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 effectdue to inner shell electrons.

• Electronegativity is a measure of the attraction an atom in a bond has for the electronsof 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 effectdue to inner shell electrons.

3.7 Resources

Texts










Q35: Menedeleev is famous for producing the Periodic Table on which the modern version isbased.

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.

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Q36: Which of the following statements about the Periodic Table are true?

a) There is a steady decrease in density on going down Group 1.b) There is an increase in electronegativity on going down Group 7.c) There is a decrease in first ionisation energy on going down Group 0.d) There is a decrease and then an increase in boiling point on crossing a Period from left

to right.e) There is a steady increase in melting point across a Period from left to right.f) There is a steady decrease in atomic size across a Period from left to right.

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Q37: Which property of the Group 1 elements could be represented by the following graph?

Li Na K Rb

Arbitrary Scale

a) Electronegativityb) Atomic sizec) Melting pointd) First ionisation energy

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Q38: Some information about four atoms A, B, C and D is as follows.

AB

CD

Covalent Radius

Number of occupied electron shells

A 3

B 4

C 3

D 4

Which atom will have the largest nuclear charge?

a) Ab) Bc) Cd) D

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Q39: Calcium has a larger covalent radius than magnesium because calcium has:

a) a smaller first ionisation energy.b) a larger nucleus.c) a larger nuclear charge.d) more occupied electron shells.

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Q40: The difference in atomic size between sodium and chlorine is mainly due to the numberof:

a) electron shells.b) neutrons.c) protons.d) electrons.

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Q41: Which of the following equations represents the first ionisation energy of magnesium?

a) Mg(s) → Mg+(g) + e-

b) Mg(g) → Mg2+(g) + 2e-

c) Mg(g) → Mg+(g) + 2e-

d) 12Mg(g) → 1

2Mg2+(g) + e−

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Q42: Which of the following equations represents the first ionisation energy of fluorine?

a) F-(g) → F(g) + e-

b) 12F2 (g) → F + (g) + e−

c) F(g) + e- → F-(g)d) F(g) → F+(g) + e-

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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.

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Q44: The change in first ionisation energy from Li to F is mainly due to:

a) increasing number of electron shells.b) increased screening effect.c) increasing number of electrons.d) increasing nuclear charge.

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Q45: The bar graph shows the variation in the first ionisation energy with atomic number forsixteen consecutive elements in the Periodic Table. The element at which the bar graph startsis not specified.

Atomic Number

ZFirst Ionisation

Energy(kJ mol-1)

Element Z is identified in the bar graph. In which Group of the table is element Z?

a) 1b) 3c) 5d) 6

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Q46: Which of the following graphs shows the variation in first ionisation energy on goingfrom left to right across Period 2?

a)

Atomic Numberb)

Atomic Number



c)

Atomic Numberd)

Atomic Number. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q47: Which of the following will not affect the first ionisation energy of an element?

a) Screening effectb) Atomic sizec) Atomic numberd) Atomic mass

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Q48: �� is a measure of the ability of an atom in a bond to attract the bondingelectrons.

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Q49: �� has the highest attraction for bonding electrons.

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Q50: Atoms of different elements have different attractions for the electrons in a bond. Thisproperty shows Periodic variation.

Which of the following shows the correct trends within the Periodic Table?

a)

b)

c)

d)

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Q51: Why are no electronegativity values quoted for the Group 0 elements?

a) They have a value of zero.b) They generally do not form bonds with other elements.c) They are unreactive non-metals.d) Their electronegativities are too high for the scale.


Q52: Explain fully why the first ionisation energy increases on going across Period 2 from leftto right. (3 marks)

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Q53: Explain fully why the first ionisation energy decreases on going down Group 1. Youshould mention nuclear charge, covalent radius and screening effect. (3 marks)


75

Topic 4

Bonding continuum and polar covalentbonding

Contents4.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 774.2 Polar covalent bonds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 784.3 Predicting bonding type using electronegativity . . . . . . . . . . . . . . . . . . . . . . . . . . 814.4 Polar molecules . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 83

4.4.1 Polar molecules and molecular polarity . . . . . . . . . . . . . . . . . . . . . . . . . . 854.5 The bonding continuum . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 884.6 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 904.7 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 904.8 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 91

Prerequisites







• the first twenty elements in the Periodic Table can be categorised according to bondingand structure:

◦ metallic (Li, Be, Na, Mg, Al, K, Ca);

◦ covalent molecular (H2, N2, O2, F2, Cl2, P4, S8 and fullerenes (eg C60));

76 TOPIC 4. 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 electronsof the bond (Higher, Unit 1, Topic 3);

• electronegativity values increase across a Period and decrease down a Group (Higher,Unit 1, Topic 3);

• trends in electronegativity can be explained in terms of atomic size, nuclear charge andthe screening effect due to inner shell electrons (Higher, Unit 1, Topic 3).

Learning objective


• 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 theircommon attraction for the shared pair of electrons;

• explain that polar covalent bonds are formed when the attraction of the atoms for thepair of bonding electrons is different;

• explain that delta positive (δ+) and delta negative (δ-) notation can be used to indicatethe partial charges on atoms, which give rise to a dipole;

• state that pure covalent bonding and ionic bonding can be considered as being atopposite ends of a bonding continuum with polar covalent bonding lying between thesetwo extremes;

• state that the larger the difference in electronegativities between bonded atoms, themore polar the bond will be;

• state that, if the difference is large, then the movement of bonding electrons from theelement 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 alwaysionic;

• physical properties of a compound, such as its state at room temperature, melting point,boiling point, solubility and electrical conductivity, should be used to deduce the type ofbonding and structure in the compound.


TOPIC 4. BONDING CONTINUUM AND POLAR COVALENT BONDING 77

4.1 Prior knowledge



a) transfer of electrons.b) a shared pair of electrons.c) delocalised electrons.d) gaining electrons.

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Q2: Why are no electronegativity values quoted for the Group 0 elements?

a) They have a value of zero.b) They are unreactive non-metals.c) They generally do not form bonds with other elements.d) Their electronegativities are too high for the scale.

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Q3: Which of the following elements has the greatest attraction for the shared pair ofelectrons in a bond?

a) Fluorineb) Carbonc) Hydrogend) Chlorine

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Q4: Which of the following elements has the least attraction for the shared pair of electronsin a bond?

a) Fluorineb) Carbonc) Hydrogend) Chlorine



4.2 Polar covalent bonds

Previously, you saw how the electronegativity of atoms is a measure of their attraction for electronsin a bond. Only when the two atoms bonded by a covalent bond are the same (e.g. Cl 2 in chlorinegas, C-C in diamond) will the electrons be exactly shared equally. In other cases there will bean unequal distribution of the electrons and the atom with the higher electronegativity will have agreater share of the electrons.

Since electrons carry a negative charge, an unequal distribution will result in the bond having apartial negative charge (called 'delta negative' and drawn δ-) where there is an excess of electronsaround the most electronegative atom, and a partial positive charge (δ+) where there is adeficiency. This is not to be confused with a full charge as found on ions.

Figure 4.1: Polar bond

In the case of HCl (Figure 4.1), chlorine, with electronegativity 3·0, becomes negative whilehydrogen, with electronegativity 2·2, becomes positive. The δ partial charge is about 0·17 of afull charge.

This type of bond is called a polar covalent bond (sometimes abbreviated to a polar bond). Thegreater the difference between the electronegativities of the atoms, the greater the distortion of theelectrons in the bond, and the greater the charge distribution. This effect is shown in the followingactivity.



Go onlineElectronegativity

In Figure 4.2, the darker shading shows the electrons being pulled towards the moreelectronegative atom.

Figure 4.2: Electronegativity and charge

Q5: How much ionic character could this bond have?

a) 0%b) 50%c) 100%

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Q6: This bond has almost zero ionic character. It could be:



a) Li-Fb) Mg-Oc) C-S

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Q7: This bond could be:

a) X is Na and Y is Ob) X is S and Y is Kc) X is N and Y is Cl

Use the table of electronegativities, where necessary, (Page 11 of the SQA HigherChemistry Data Booklet) to answer the following questions.

Q8: Which molecule contains polar bonds?

a) Hydrogen (H2)b) Chlorine (Cl2)c) Hydrogen chloride (HCl)d) Nitrogen (N2)

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Q9: Which molecule contains the most polar bonds?

a) Hydrogen fluoride (HF)b) Hydrogen chloride (HCl)c) Hydrogen bromide (HBr)d) Hydrogen iodide (HI)

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Q10: Which molecule has the bonds that are most polar?

a) Hydrogen sulfide (H2S)b) Phosphine (PH3)c) Methane (CH4)d) Ammonia (NH3)

The symbol can be used to show the direction of the dipole, the arrow pointing to thenegative side.



Go onlineDirection of charge

Use the electronegativity values (Page 11 of the SQA Higher Chemistry Data Booklet) to helpshow 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 electronsunequally causing the atoms to have partial positive and negative charges.

The polarity of a covalent bond depends on the difference in electronegativity between thebonded atoms, the most electronegative becoming more negative.

4.3 Predicting bonding type using electronegativity

One of the main factors determining the type of bond formed between two elements in a compoundis the difference in electronegativity. As previously discussed, electronegativity is a periodicproperty, with increasing electronegativity as you move across a Period from left to right, anddecreasing electronegativity as you move down a Group.

The relationship between electronegativity and bond type is summarised as follows.



Figure 4.3: 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 (SiO2) and silicon carbide (SiC,carborundum), which has a similar structure to diamond. Both of these are used as abrasiveson account of their hardness.

Use the electronegativity values on page 11 of the SQA Higher Chemistry Data Booklet to help youanswer the following questions.

Q12: Hydrogen, bromine and potassium have electronegativities of 2·2, 2·8 and 0·8respectively. What type of bonding would you expect:

• between two bromine atoms in a Br2 molecule?

• between hydrogen and bromine in HBr?

• between potassium and bromine in KBr?

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Q13: What type of bonding would you expect between the Group 1 metal caesium and theGroup 6 element sulfur?

a) Pure covalentb) Polar covalentc) Ionicd) Metallic

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Q14: What type of bonding would you expect between phosphorus and hydrogen inphosphine (PH3) ?

a) Pure covalentb) Polar covalentc) Ionicd) Metallic

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Q15: Which of the following elements will bond with oxygen, to produce a polar bond with apartial positive charge on the oxygen atom?

a) Carbonb) Fluorinec) Hydrogend) Lithium

Key point

The type of bonding between two atoms depends mainly on the difference inelectronegativity between the atoms.

When the difference is zero, the bond will be covalent. With a small difference, a polarcovalent bond is likely. When the difference is large, fully charged ions are produced andionic bonding will be predicted.

4.4 Polar molecules

When two atoms of different electronegativity are bonded together by sharing electrons, one atomattracts the electrons more than the other and a polar bond results. In simple molecules likehydrogen chloride and iodine chloride this polar bond has a permanent dipole.

Polar molecules are attracted to one another by forces called permanent dipole-permanent dipoleinteractions (see Figure ??) as well as the London dispersion attractions caused by themovement of electrons observed in the last section.



Go onlinePermanent dipole interactions

Figure 4.4: Permanent dipoles in iodine chloride

How do permanent dipole-permanent dipole interactions compare to London dispersionforces of attraction in terms of strength?

A comparison of strengths can be made by looking at the following table.

Figure 4.5: Boiling points of halogen containing compounds

A B C

Cl - Cl I - Cl Br - Br

-34◦C 97◦C 59◦C

Any comparison of boiling points has to be between molecules of similar size and shape sothat the London dispersion forces are similar.

Q16: Which molecule has permanent dipole-permanent dipole interactions?

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Q17: Which other molecule would have almost the same London dispersion forces?

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Q18: Which of the molecules has the strongest intermolecular forces?



Key point

Permanent dipole-permanent dipole interactions act in addition to London dispersionelectrostatic attractions between polar molecules and are stronger than these attractions formolecules of equivalent size.

4.4.1 Polar molecules and molecular polarity

Not all covalent molecules with polar bonds result in polar molecules. Those molecules which arehighly symmetrical prove to be non-polar (see Figure 4.6).

Figure 4.6: Non-polar molecules

In each of the two molecules shown, although there is a difference in electronegativity in the polarbonds, the charge is distributed around the central carbon atoms with the positive and negativecharges balancing out. The molecules have no overall dipole.

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 4.7), the dipoles cannotcancel each other out. If the molecule has a permanent slight positive charge at one side and anegative charge at the other, then it is a polar molecule. Bond polarity can thus be predicted fromelectronegativity differences and molecular polarity can be predicted from electronegativitydifferences if we also take into account the shape of the molecule.

The symbol can be used to show the direction of the dipole, the arrow pointing to thenegative side. The molecules shown in the following figure are both polar molecules.

Figure 4.7: Polarity of molecules



Go onlinePredicting molecular polarity

Electronegativity values are used to predict the polarity of bonds, and the shapes ofmolecules are used to predict molecular polarity. Page 11 in the SQA Higher Chemistry DataBooklet may be helpful.

Q19: Complete the table using the following molecules, bond arrangements and diploes.

Go onlineDetecting polar molecules

An experiment is performed to show the attractions between polar molecules andelectrostatically charged rods.

Figure 4.8

Water Paraffin

ChargedRods



Q20: Which liquid is affected by the charged rod?

a) Waterb) Paraffin

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Q21: Which liquid is polar?

a) Waterb) Paraffin

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Q22: Which of these statements is correct when a jet of polar molecules passes a chargedrod?

a) It is not affected by the rod.b) It is attracted by the rod.c) It is repelled by the rod.d) It is sped up by the rod.

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Q23: Which of these statements is correct when a jet of non-polar molecules passes acharged rod?

a) It is not affected by the rod.b) It is attracted by the rod.c) It is repelled by the rod.d) It is sped up by the rod.

Key point

An electrostatically charged rod can be used to detect the presence of polar molecules in aliquid. Polar molecules are attracted to both a negative and positive rod.



4.5 The bonding continuum

Covalent bonding usually involves elements close to each other in the Periodic Table. Ionic bondinggenerally 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 andcovalent bonding as being at opposite ends of a bonding continuum with varying degrees of polarcovalent bonding lying between the extremes (see Figure 4.9).

Figure 4.9: Bonding spectrum

In most cases, the bigger the difference in electronegativity between the atoms, the more polarthe bond and the greater the ionic character. However, other factors make a contribution and careneeds to be taken before jumping to conclusions. Identical atoms share the electrons in a covalentbond equally. In the hydrogen molecule there is no charge distribution. Non-identical atoms attractthe bonding electrons unequally, e.g. when hydrogen and chlorine bond, chlorine attracts electronsmore strongly than hydrogen. A polar bond results (see Figure 4.10).

Figure 4.10: Polar bond

In some cases, the bond polarity results in a polar molecule. In certain symmetrical molecules thepolar nature of the bonds tends to cancel out (see Figure 4.11).



Figure 4.11: Polar and non-polar molecules

In water, the molecule is asymmetrical so the δ+ on the hydrogen atoms will add and the moleculewill 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 bondswill cancel out and the molecule will be non-polar, despite having polar bonds.

Q24: Will trichloromethane (CHCl3, chloroform) be a polar or non-polar molecule?

a) Polarb) Non-polar



4.6 Summary

Summary

• Polar covalent bonds occur when the atoms of the bond attract the bonding electronsunequally, causing the atoms to have partial positive and negative charges.

• The polarity of a covalent bond depends on the difference in electronegativity betweenthe 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 inelectronegativity 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 willbe predicted.

• Permanent dipole-permanent dipole interactions act in addition to London dispersionelectrostatic attractions between polar molecules and are stronger than theseattractions 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 moleculesin 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 basedon electronegativity.

4.7 Resources

Texts










Q25: In which of the following compounds do both ions have the same electron arrangementas argon?

a) Calcium bromideb) Magnesium oxidec) Sodium sulfided) Calcium sulfide

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Q26: Which element is a solid at room temperature and consists of discrete molecules?

a) Sulfurb) Carbonc) Silicond) Boron

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Q27: Graphite, a form of carbon, conducts electricity because it has:

a) metallic bonding.b) pure covalent bonding.c) delocalised electrons.d) London dispersion forces.

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Q28: Which of the following is true of lithium but not of carbon?

a) Covalentb) Metallicc) Made up of discrete molecules.d) Made up of diatomic molecules.e) Gasf) Solid

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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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a) Boronb) Chlorinec) Nitrogend) Phosphoruse) Sodiumf) Sulfur

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Q31: Identify the two elements which exist as discrete covalent molecular solids.

a) Boronb) Chlorine

c) Nitrogend) Phosphorus

e) Sodiumf) Sulfur

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a) Boron

b) Chlorinec) Nitrogen

d) Phosphoruse) Sodium

f) Sulfur

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Q33: In which molecule will the chlorine atom carry a partial positive charge (δ+)?

a) Cl-Fb) Cl-Brc) Cl-Cld) Cl-I

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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) Conductionwhen 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

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Q35: Which of the following chlorides is likely to have least ionic character?

a) LiClb) CsClc) BeCl2d) CaCl2

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Q36: Which type of structure is found in a substance melting at 1263◦C which conductselectricity when molten, but not when solid?

a) Covalent (network structure)b) Covalent (discrete molecules)c) Ionicd) Metallic

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Q38: Identify the two substances which are ionic.

a) NH4Cl (s)

b) CH3OH (l)

c) C6H14 (l)

d) SO2 (g)

e) Na2CO3 (s)

f) SiO2 (s)

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95

Topic 5

Intermolecular forces

Contents5.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 985.2 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 985.3 London dispersion forces . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1005.4 Hydrogen bonding . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1045.5 Relating properties to bonding . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 107

5.5.1 Boiling point . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1075.5.2 Density . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 108

5.6 Viscosity . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1105.7 Predicting solubilities from solute and solvent polarities . . . . . . . . . . . . . . . . . . . . . 1115.8 The solubility of flavour molecules . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1155.9 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1165.10 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1185.11 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 118

Prerequisites


• the first twenty elements in the Periodic Table can be categorised according to bondingand structure:

◦ metallic Li, Be, Na, Mg, Al, K, Ca;◦ covalent molecular H2, N2, O2, F2, Cl2, P4, S8 and fullerenes (eg C60);◦ 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 electronsof the bond (Higher, Unit 1, Topic 3);

• electronegativity values increase across a Period and decrease down a Group (Higher,Unit 1, Topic 3);

• this can be explained in terms of atomic size, nuclear charge and the screening effectdue to inner shell electrons (Higher, Unit 1, Topic 3);

96 TOPIC 5. INTERMOLECULAR FORCES


• in a covalent bond, atoms share pairs of electrons (National 4, Unit 1);

• the covalent bond is a result of two positive nuclei being held together by their commonattraction 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 bondingelectrons is different (Higher, Unit 1, Topic 4);

• delta positive and delta negative notation can be used to indicate the partial charges onatoms, which give rise to a dipole (Higher, Unit 1, Topic 4);

• pure covalent bonding and ionic bonding can be considered as being at opposite endsof a bonding continuum with polar covalent bonding lying between these two extremes(Higher, Unit 1, Topic 4);

• the larger the difference in electronegativities between bonded atoms, the more polarthe bond will be (Higher, Unit 1, Topic 4).

Learning objective


• state that all molecular elements and compounds and monatomic elements condenseand freeze at sufficiently low temperatures - for this to occur, some attractive forcesmust exist between the molecules or discrete atoms;

• state that any 'intermolecular' forces acting between molecules are known as van derWaals' forces;

• state that there are several different types of van der Waals' forces such as:

◦ London dispersion forces;

◦ permanent dipole-permanent dipole interactions;

◦ hydrogen bonding (another example of permanent dipole-permanent dipoleinteractions);

• state that London dispersion forces are forces of attraction that can operate between allatoms 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 attractionbetween temporary dipoles and induced dipoles, caused by movement of electrons inatoms and molecules;

• state that the strength of London dispersion forces is related to the number of electronswithin an atom or molecule;

• state that a molecule is described as polar if it has a permanent dipole - the spatialarrangement of polar covalent bonds can result in a molecule being polar;


TOPIC 5. INTERMOLECULAR FORCES 97

Learning objective continued

• state that permanent dipole-permanent dipole interactions are additional electrostaticforces of attraction between polar molecules;

• state that permanent dipole-permanent dipole interactions are stronger than Londondispersion forces for molecules with similar numbers of electrons;

• state that bonds consisting of a hydrogen atom bonded to an atom of a stronglyelectronegative element such fluorine, oxygen or nitrogen are highly polar;

• state that hydrogen bonds are electrostatic forces of attraction between molecules whichcontain these highly polar bonds;

• state that a hydrogen bond is stronger than other forms of permanent dipole-permanentdipole interaction, but weaker than a covalent bond;

• explain that melting points, boiling points and viscosity can all be rationalised in termsof 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 meltingand boiling points of non-polar substances with similar numbers of electrons;

• state that the anomalous boiling points of ammonia, water and hydrogen fluoride are aresult of hydrogen bonding;

• explain that boiling points, melting points, viscosity and solubility/miscibility in water areproperties of substances which are affected by hydrogen bonding;

• explain that hydrogen bonding between molecules in ice results in an expandedstructure which causes the density of ice to be less than that of water at lowtemperatures;

• predict the solubility of a compound by considering the following key features:

◦ presence in molecules of O-H or N-H bonds, which implies hydrogen bonding;

◦ spatial arrangement of polar covalent bonds, which could result in a moleculepossessing a permanent dipole.



5.1 Prior knowledge


Q1: Which of the following is a measure of the ability of an atom in a bond to attract thebonding electrons?

a) Ionisation Energyb) Periodicityc) Electronegativityd) Polarity

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Q2: Which of the following elements has the greatest attraction for the shared pair ofelectrons in a bond?

a) Fluorineb) Nitrogenc) Phosphorusd) Lithium

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Q3: As you move across the Periodic Table, electronegativity:

a) stays the same.b) decreases.c) increases.d) decreases then increases.

5.2 Introduction

Simple molecules like nitrogen N2, methane CH4 and water H2O, have their atoms held together bycovalent 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 avery low temperature. Water is a liquid at room temperature and other covalent molecules likecandle wax are solid. Forces of attraction must exist between all these covalent molecules or itwould never be possible to have covalent molecular solids at all. The molecules would fly apart andall 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).



Figure 5.1 shows both intermolecular and intramolecular forces in a sample of water.

Figure 5.1

The way that the charge is distributed in a water molecule is not symmetrical. There is morenegative charge on one side than the other. It is said to have a dipole. The Greek letter delta, δ, isused to show a 'small amount'.

Van der Waals' forces

Johannes Diderik van der Waals recognised that relatively weak forces were responsible for thechange 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 followingsections.

• London dispersion forces.

• Permanent dipole-permanent dipole interactions.

• Hydrogen bonding.

In this topic, the forces of attraction caused by dipoles will be explored.



5.3 London dispersion forces

The strength of the intermolecular forces between covalent substances can be estimated by lookingat the boiling points. The lower the boiling point, the weaker the forces must be, since it is theintermolecular forces which have to be overcome to change the liquid into gas.

Water

Steam

Heat energy addedLiquid to gas

The covalent diatomic elements like hydrogen, H2 gas, and bromine, Br2, which is a liquid at roomtemperature, 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 couldhave a temporary dipole at a particular instant in time and that if there were other atoms ormolecules nearby; this temporary dipole might affect them and induce a dipole in the nearby atomor molecule. The resultant electrostatic attraction between the temporary dipole and the induceddipole might, although very weak, be able to hold the substance together.

These forces of attraction are known as London dispersion forces.



Go onlineInduced dipoles

Figure 5.2 shows how the temporary dipole in a helium atom can exist at a single briefinstant in time. This can induce a dipole in its neighbour. Hydrogen molecules are shownforming the same temporary dipole-induced dipole pair. The attractions between the dipolepairs form the same type of bonds; London dispersion forces.

Figure 5.2: Temporary and induced dipoles

Q4: When the electron movement is paused, as summarised in Figure 5.2, the electrondistribution is:

a) evenly spread.b) unevenly spread.

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Q5: This distribution of electrons causes:

a) a permanent dipole.b) a temporary dipole.

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Q6: What effect does this have on the neighbouring particle?

a) The permanent dipole causes an induced dipole.b) The temporary dipole causes an induced dipole.c) The permanent dipole induces a temporary dipole.d) The temporary dipole induces a permanent dipole.

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Q7: The London dispersion forces of attraction formed can be said to be:

a) intramolecular.b) intermolecular.



London dispersion forces of attraction can operate between all atoms and molecules. They aremuch weaker than other types of bonding, at around 2 kJ mol-1 compared to covalent bonds ofbetween 150-500 kJ mol-1. In many substances, London dispersion forces are not large enoughto 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 Londondispersion 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.

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Q9: Which noble gas shown has the highest boiling point?

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Q10: Which noble gas shown must have the strongest London dispersion forces holding theatoms together?

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Q11: Look at the relationship between the London dispersion forces and size. Which of thesestatements is true?

a) The size of the atom doesn't affect the London dispersion forces.b) The larger the atom the weaker the London dispersion forces.c) The larger the atom the stronger the London dispersion forces.d) 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?

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 dispersionforces. The following table shows the boiling points of the first four alkane molecules.



Name Formula Boiling point (◦C)

Methane CH4 -162

Ethane C2H6 -89

Propane C3H8 -42

Butane C4H10 -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 explainyour answer fully, work through the next four questions before attempting the complete answerin the repeated question. If you feel confident, check your answer.)

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Q14: Name the alkane shown which has the largest molecules.

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Q15: Name the alkane shown which has the highest boiling point.

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Q16: Name the alkane shown which must have the strongest London dispersion forces.

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Q17: Which relationship between molecule size and London dispersion forces is true?

a) The size of the molecule doesn't affect the London dispersion forces.b) The larger the molecule the weaker the London dispersion forces.c) The larger the molecule the stronger the London dispersion forces.d) The smaller the molecule the stronger the London dispersion forces.

Key point

London dispersion forces of attraction operate between all atoms and molecules and areweaker than all other types of bonding. The strength of London dispersion forces is related tothe size of the atoms or molecules.



5.4 Hydrogen bonding

Hydrogen bonding is the name given to an intermolecular force which is actually a type ofpermanent dipole-permanent dipole attraction (Topic 4). These hydrogen bonds are separatedfrom other examples of this type of interaction because they are unusually strong (about 15-25 kJmol-1) as shown in the following diagram.

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 theseconditions.

Hydrogen bonding elements

Notice that the bonds called hydrogen bonds are the forces of electrostatic attraction between themolecules which contain these highly polar bonds, i.e. the hydrogen bonds are the intermolecularforces.

Key point

Bonds consisting of a hydrogen atom bonded to an atom of a strongly electronegativeelement, such as fluorine, oxygen or nitrogen, are highly polar and the electrostatic attractionsbetween molecules which contain these highly polar bonds are called hydrogen bonds.



Go onlineStrength of hydrogen bonds

The first diagram shows the formation of hydrogen bonds in a sample of water, and thefollowing diagram shows the boiling points and relative sizes of the first four hydrides of theGroup 6 elements.

Strength of hydrogen bonds

Boiling points of the Group 6 hydrides

A comparison of the strength of hydrogen bonds with other permanent dipole-permanentdipole interactions and London dispersion forces can be made by considering the data in thediagram. Remember that London dispersion forces of attraction will operate between allatoms and molecules, and that some of the hydride molecules will have permanentdipole-permanent dipole attractions since they are polar molecules.

Q18: Use the boiling point data to compare the intermolecular forces present in water withthose present in the other Group 6 hydrides and explain fully how hydrogen bonds comparein 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 explainyour answer fully, work through the next four questions before attempting the complete answerin the repeated question. If you feel confident, check your answer.)

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Q19: Name the type of weak intermolecular force of attraction which increases as sizeincreases.

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Q20: Name the substance which shows boiling point evidence which goes against this trend.

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Q21: Name the substance which has the strongest intermolecular force of attraction holdingthe molecules together.

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Q22: How do the hydrogen bonds in water compare in strength to the London dispersionforces present and the other permanent dipole-permanent dipole interactions present in someof the other molecules?

a) The hydrogen bonds are weaker than London dispersion forces and the other permanentdipole-permanent dipole interactions.

b) The hydrogen bonds are stronger than London dispersion forces and the other permanentdipole-permanent dipole interactions.

c) The hydrogen bonds are stronger than London dispersion forces but weaker than theother permanent dipole-permanent dipole interactions.

A comparison of the strength of hydrogen bonds with intramolecular bonds like the covalentbond between hydrogen and oxygen can be made by remembering that the hydrogen bondhas a value of about 15-25 kJ mol-1 and the strength of the covalent O-H bond can be foundon Page 10 of the SQA Higher Chemistry Data Booklet. The table of 'Mean Bond Enthalpies'gives the value.

Q23: Which statement is true about hydrogen bonds in comparison to covalent bonds?

a) Covalent bonds are much stronger than hydrogen bonds.b) Hydrogen bonds are much stronger than covalent bonds.c) Covalent bonds are the same strength as hydrogen bonds.d) Hydrogen bonds are twice as strong as covalent bonds.

Key point

A hydrogen bond is stronger than other forms of permanent dipole-dipole interaction butweaker than a covalent bond.



5.5 Relating properties to bonding

Intermolecular bonds are the relatively weak bonds which attract molecules to each other. Theyare not as strong as the bonds, for example covalent bonds, that bind the atoms together into amolecule.

Q24: Which of the following properties of a substance are a reflection of the intermolecularforces 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 involveslooking for water or evidence that water was present at some time. Water is the commonest liquidon Earth and also one of the most unusual! The next sections explore some properties of waterfrom the point of view of hydrogen bonding.

5.5.1 Boiling point

The boiling points of the hydrides along the Periods for Groups 4 to 7 elements are:

Boiling points of several hydrides



Go onlineHydrides of non-metals

Q25: The trend in boiling points in Group 4 from CH4 to SnH4 shows a regular increase withformula mass. Is there any evidence of polar attractions between molecules?

a) Yesb) No

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Q26: Three of the second period hydrides have higher boiling points than expected bycomparison with the trends seen for other elements. (Remember that higher formula massusually means higher boiling point.)

Which are the three hydrides?

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Q27: What is the main force between molecules in group 4 hydrides, from CH4 to SnH4?

a) Polar-polar attractionsb) Van der Waals attractionsc) Covalent bondsd) Ionice) Hydrogen bonds

Key point

The higher-than-expected boiling points of ammonia, water and hydrogen fluoride resultfrom additional intermolecular forces, in this case, hydrogen bonds.

5.5.2 Density

Most substances contract when they are cooled and solids are normally denser than their ownliquids. This causes the majority of solids to sink when placed in their own liquid. The structure ofice, 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 astrong 'open' structure with large spaces in it (see the following diagram). This makes ice lessdense and able to float on water.



The arrangement of water molecules in ice maximises the hydrogen bonding between them andleads to an open structure. As it melts the structure collapses into the spaces and the liquid becomesless dense.

The fact that ice floats on water is good news for fish living in ponds. The ice forms on the surfaceand the fish and other aquatic life can exist below this layer. The layer of ice thermally insulates thewater beneath.

A disadvantage of water expanding as it forms ice is that water trapped in pipes over a cold spellcauses 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 causesthe density of ice to be less than that of water at low temperatures.



5.6 Viscosity

The stronger the intermolecular forces are between molecules in a liquid, the more viscous theliquid. The viscosity is a measure of how easily it flows (how thick or syrupy a liquid is). Themolecules have to be able to move past each other to flow.

Go onlineTesting viscosity

Figure 5.3: Testing the viscosity of three different liquids

Q28: In which liquid does the marble fall fastest?

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Q29: Which chemical structure has the least hydrogen bonding?

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Q30: Which chemical structure has the most hydrogen bonding?

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Q31: Given that the marble in water took two seconds to reach the bottom, calculate its rateof 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) 0·8 m s-1

b) 0·6 m s-1

c) 0·4 m s-1

d) 0·1 m s-1

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Q33: Assuming your answer to the last question is correct, how long would it take a marbleto travel through the ethoxyethane?

Viscosity is a term which applies to fluids (both liquids and gases), and since these are covalentsubstances, 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.

5.7 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 willdissolve 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 itdissolve.

A polar solvent is one whose molecules exhibit strong permanent dipoles. Because of this, polarsolvents like water can generally dissolve polar substances, like sucrose, and ionic solids, forexample 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 aremore 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. Oppositesdo not dissolve.

Polar and ionic substances tend to be insoluble in non-polar solvents. Non-polar substances tendto 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. Thepolar water molecules can sometimes pull the ions into solution and surround them.

Dissolving solid sodium chloride

The way the water molecules direct themselves depends on the charge the ion carries. The ionsare said to be hydrated.

Dissolving of sodium chloride can be shown in an equation.



Dissolving sodium chloride

A polar molecule, like hydrogen chloride, dissolves in water as the water 'interacts' with themolecules.

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 (C12H22O11), containing several -OH groups,can bond with polar water molecules.



Q34: Ammonia (NH3) is a gas which is very soluble in water, whereas nitrogen (N2) is almostinsoluble. Can you explain why?

Key point

Ionic compounds and polar molecular compounds tend to be soluble in polar solvents such aswater and insoluble in non-polar solvents. Non-polar molecular substances tend to be solublein non-polar solvents and insoluble in polar solvents.

Miscibility

When two liquids mix thoroughly with no visible boundary between them, they are said to bemiscible, e.g. water and methanol. Both of these molecules have hydrogen bonds shown in Figure5.4 (a).

Miscibility arises when the intermolecular attractions between two types of substances are fairlysimilar. They are then able to mix easily.

If one substance has different intermolecular attractions from the other, its molecules stay groupedaround each other and form a separate layer, e.g. water (which is hydrogen bonded), andtetrachloromethane (which is non-polar) shown in Figure 5.4 (b).

Figure 5.4: Miscible and immiscible liquids

Whether two liquids are miscible or immiscible can only be determined by experiment, but thegeneral rule of like dissolves like can be useful in trying to make a rough prediction.



5.8 The solubility of flavour molecules

Our sense of smell can detect an enormous number of different types of molecules. Two examples:toluene, an aromatic hydrocarbon named because of its pleasant odour, and methanol, both ofwhich have distinct odours. The former is almost insoluble in water; the latter will mix with water inall proportions (because they can form hydrogen bonds with each other easily).

Chemical analysis of the volatile compounds in asparagus showed a majority of these were alcoholsand aldehydes. Hexanal, hexenal and oct-1-en-3-ol predominated.

These three compounds are all fairly soluble in water so what do you imagine will happen to theflavour of asparagus if it is cooked by boiling it in water? The cooked asparagus will lose its essentialtaste which will be thrown out with the cooking water.

How do you get round this problem? Substances that are soluble in water are not usually solublein oils so heat the asparagus in cooking oil or butter and the flavour will remain with the vegetable.On the other hand, the flavour components of broccoli are not very soluble in water so boiling orsteaming are appropriate cooking methods.

Asparagus Broccoli

Go onlineWhich cooking method?

For the following compounds, decide whether they will be soluble in water or oil and whetherthey will be volatile or not. Both of these properties depend on the intermolecular forcesbetween molecules so focus on that aspect of each compound.

Q35: Aspartame

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Q36: Limonene

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Q37: Vanillin

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Q38: Capsaicin

You might like to discuss some foods that you particularly like or dislike with members of yourclass, trying to give some reasons for your choice.

5.9 Summary

Summary

• All molecular elements, compounds and monatomic elements condense and freezeat sufficiently low temperatures. For this to occur, some attractive forces must existbetween 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 dispersionforces and permanent dipole-permanent dipole interactions, which include hydrogen



Summary continued

bonding.

• London dispersion forces are forces of attraction that can operate between all atomsand molecules.

• These forces are much weaker than all other types of bonding.

• They are formed as a result of electrostatic attraction between temporary dipoles andinduced dipoles caused by movement of electrons in atoms and molecules.

• The strength of London dispersion forces is related to the number of electrons within anatom or molecule.

• Bonds consisting of a hydrogen atom bonded to an atom of a strongly electronegativeelement, such fluorine, oxygen or nitrogen, are highly polar.

• Hydrogen bonds are electrostatic forces of attraction between molecules which containthese highly polar bonds.

• A hydrogen bond is stronger than other forms of permanent dipole-permanent dipoleinteraction, but weaker than a covalent bond.

• Melting points, boiling points and viscosity can all be rationalised in terms of the natureand strength of the intermolecular forces which exist between molecules.

• By considering the polarity and number of electrons present in molecules, it is possibleto make qualitative predictions of the strength of the intermolecular forces.

• The melting and boiling points of polar substances are higher than the melting andboiling points of non-polar substances with similar numbers of electrons.

• The anomalous boiling points of ammonia, water and hydrogen fluoride are a result ofhydrogen bonding.

• Boiling points, melting points, viscosity and solubility/miscibility in water are propertiesof substances which are affected by hydrogen bonding.

• Hydrogen bonding between molecules in ice results in an expanded structure whichcauses 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 solventssuch as water and insoluble in non-polar solvents.

• Non-polar molecular substances tend to be soluble in non-polar solvents and insolublein polar solvents.

• The solubility of a compound can be predicted by considering the following key features:

◦ presence in molecules of O-H or N-H bonds, which implies hydrogen bonding;

◦ spatial arrangement of polar covalent bonds, which could result in a moleculepossessing a permanent dipole.



5.10 Resources

Texts







Q39: London dispersion forces of attraction are a result of:

a) an induced dipole causing an temporary dipole.b) a permanent dipole causing an permanent dipole.c) an induced dipole causing an permanent dipole.d) a temporary dipole causing an induced dipole.

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Q40: Which of these is the strongest?

a) Hydrogen bondsb) Covalent bondsc) London dispersion forcesd) Permanent dipole-permanent dipole interactions

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Q41: What is the connection between the size of molecules and the London dispersion forcesof attraction?

a) The size makes no difference.b) The smaller the molecule the weaker the force.c) The smaller the molecule the stronger the force.d) The weakest force is between the largest molecules.

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Q42: Which of these molecules has the largest permanent dipole?

a) HFb) HClc) HBr




d) HI

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Q43: 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.

Look at the following table which shows the boiling points of some substances.

Substance Boiling Point ◦C Bonds (or forces) broken at the boilingpoint

Sodium 883 Metallic

Neon -246 A

Water 100 B

Q44: Which word(s) should be inserted at A?

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Q45: Which word(s) should be inserted at B?

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Q46: Identify the substance which has neither covalent or metallic bonding.

a) NH3 (l)b) CCl4 (l)c) He (g)d) Hg (l)e) H2 (g)f) HBr (l)

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Q47: Identify the substance which has hydrogen bonding between the molecules.


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Q48: Identify the substances which are polar molecules.


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Q49: Identify the substance which has polar bonds but is a non-polar molecule.


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Q50: The following diagram shows the structure of kevlar, which is strong because of theintermolecular bonding between neighbouring molecules.

Name the type of intermolecular bonding involved.

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Q51: Which area of the diagram shows the position of one such bond?


121

Topic 6

Oxidising or reducing agents

Contents6.1 Prior knowledge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 124

6.2 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 125

6.3 Elements as oxidising and reducing agents . . . . . . . . . . . . . . . . . . . . . . . . . . . . 126

6.3.1 Displacement reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 129

6.4 Molecules and group ions as oxidising and reducing agents . . . . . . . . . . . . . . . . . . . 130

6.4.1 Oxygen . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 131

6.5 Uses for strong oxidising agents . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 133

6.6 Ion-electron half equations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 136

6.7 Combining ion-electron equations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 139

6.8 Complex ion-electron equations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 141

6.9 Summary exercise . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 145

6.10 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 147

6.11 Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 148

6.12 End of topic test . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 149

Prerequisites


• balanced ionic equations can be written to show the reaction of metals with water,oxygen and acids (National 5, Unit 3);

• ion-electron equations can be written for electrochemical cells, including those involvingnon-metals (National 5, Unit 3);

• combinations of these reactions form redox equations (National 5, Unit 3);

• spectator ions can be described and identified from equations (National 5, Unit 3);

• students should be able to explain the principles of and carry out an acid / base titration(National 5, Unit 3).

122 TOPIC 6. OXIDISING OR REDUCING AGENTS

Learning objective

By the end of this topic, you should know that:

Elements as oxidising or reducing agents

• A redox reaction is a reaction in which reduction and oxidation occur together, reductionbeing the gain of electrons by a reactant and oxidation being the loss of electrons by areactant in a reaction.

• An oxidising agent is a substance which accepts electrons.

• A reducing agent is a substance which donates electrons.

• Oxidising and reducing agents can be identified in redox reactions.

• Elements with low electronegativities (metals) tend to form ions by losing electrons(oxidation) and so can act as reducing agents.

• Elements with high electronegativities (non-metals) tend to form ions by gainingelectrons (reduction) and so can act as oxidising agents.

• The strongest reducing agents are found in Group 1.

• The strongest oxidising agents are found in Group 7.

• The electrochemical series indicates the effectiveness of oxidising and reducing agents.

Compounds as oxidising or reducing agents

• Compounds can also act as oxidising or reducing agents.

• Electrochemical series contain a number of ions and molecules.

• The dichromate and permanganate ions are strong oxidising agents in acidic solutions,while hydrogen peroxide is an example of a molecule which is a strong oxidising agent.

• Carbon monoxide is an example of a gas that can be used as a reducing agent.

• Oxidising and reducing agents can be selected using an electrochemical series fromPage 12 of the SQA Higher Chemistry Data Booklet or can be identified in the equationshowing a redox reaction.

Use of oxidising agents

• Oxidising agents are widely employed because of the effectiveness with which they cankill fungi and bacteria, and can inactivate viruses.

• The oxidation process is also an effective means of breaking down colouredcompounds, making oxidising agents ideal for use as 'bleach' for clothes and hair.

Ion-electron equations

• Oxidation and reduction reactions can be represented by ion-electron equations.


TOPIC 6. OXIDISING OR REDUCING AGENTS 123

Learning objective continued

• The electrochemical series represents a series of reduction reactions.

• The strongest oxidising agents are at the bottom of the left-hand column of theelectrochemical series.

• The strongest reducing agents are at the top of the right-hand column of theelectrochemical series.

• When molecules or group ions are involved, if the reactant and product species areknown, a balanced ion-electron equation can be written by adding appropriate numbersof water molecules, hydrogen ions and electrons.

• Ion-electron equations can be combined to produce redox equations.

Practical applications

• Displacement reactions are examples of redox reactions; oxidising and reducing agentscan be identified in these and other redox reactions.



6.1 Prior knowledge


Q1: Reduction can be defined as the:

a) gain of electrons.b) loss of electrons.c) loss in mass of a substance.d) gain in mass of a substance.

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Q2: Potassium iodide will react with lead nitrate to produce a precipitate of lead iodide.Which of the following best represents this?

a) 2Pb+(aq) + 2NO3-(aq) + 2K+(aq) + 2l-(aq) → 2Pb+(l-)2(s) + 2K+(aq) + 2NO3

-(aq)b) Pb+(aq) + 2NO3

-(aq) + 2K+(aq) + 2l-(aq) → Pb2+(l-)2(aq) + 2K+(aq) + 2NO3-(aq)

c) Pb2+(aq) + 2NO3-(aq) + 2K+(aq) + 2l-(aq) → Pb2+(l-)2(s) + 2K+(aq) + 2NO3

-(aq)d) Pb2+(NO3

-)2(aq) + 2K+l-(aq) → Pb2+(l-)2(s) + 2K+NO3-(aq)

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Q3: Combine the following reactions to give the overall redox reaction.

• Oxidation H2O2(l) → O2(g) + 2H+(aq) + 2e-

• Reduction 5e- + MnO4-(aq) + 8H+(aq) → Mn2+(aq) + 4H2O(l)

a) H2O2(l) + 5e- + MnO4-(aq) + 8H+(aq) → Mn2+(aq) + 4H2O(l) + O2(g) + 2H+(aq) + 2e-

b) H2O2(l) + MnO4-(aq) + 8H+(aq) → Mn2+(aq) + 4H2O(l) + O2(g) + 2H+(aq)

c) 5H2O2(l) + 10e- + 2MnO4-(aq) + 16H+(aq) → 2Mn2+(aq) + 8H2O(l) + 5O2(g) + 10H+(aq)

+ 10e-

d) 5H2O2(l) + 2MnO4-(aq) + 16H+(aq) → 2Mn2+(aq) + 8H2O(l) + 5O2(g) + 10H+(aq)

e) 5H2O2(l) + 2MnO4-(aq) + 6H+(aq) → 2Mn2+(aq) + 8H2O(l) + 5O2(g)



6.2 Introduction

The term 'redox' is an abbreviation of the chemical reactions commonly known as reduction andoxidation. These processes occur together and include many common reactions such as theburning of fossil fuels, the corrosion of metals and photosynthesis.

Rusting is a common redox reaction in whichiron is converted into iron oxides.

Photosynthesis is another common redoxreaction, involving both reduction and oxidation

stages.

An early definition of reduction describes the removal of oxygen from metal oxides. The metaloxide is said to be reduced to the metal. An example is iron(III) oxide being reduced to iron.

2Fe2O3 + 3C → 4Fe + 3CO2

Iron oxide reduction

An early definition of oxidation describes the addition of oxygen to a metal forming a metal oxide. Afamiliar example is the rusting of iron to give iron(III) oxide. In this case the iron metal is beingoxidised.

4Fe + 3O2 → 2Fe2O3

Iron oxidation

Iron(III) oxide consists of Fe3+ and O2- ions. Close analysis of the above reaction shows that theiron atoms have lost electrons; these electrons are transferred to the oxygen.

The iron has been oxidised and the oxygen has been reduced. The changes can be shown clearlyby writing ion-electron half equations for each reaction.

4Fe → 4Fe3+ + 12e-

3O2 + 12e- → 6O2-

Iron oxide half equations

Notice that the equations have been balanced by ensuring that the number of atoms on each sideof the arrow is the same and that the number of electrons exchanged between the two halfequations is the same.



A useful memory aid for the definitions of oxidation and reduction is provided by the first letters inthe phrase 'OIL RIG'.

Oxidation Is Loss of electrons.Reduction Is Gain of electrons.

6.3 Elements as oxidising and reducing agents

The corrosion of iron results in the formation of iron(III) oxide.

4Fe + 3O2 → 2Fe2O3

Iron oxidation (2)

As already seen, this reaction can be represented by two ion-electron equations as follows.

4Fe → 4Fe3+ + 12e-

3O2 + 12e- → 6O2-

Iron oxide half equations (2)

The iron metal is being oxidised by the oxygen. Another way of expressing this is to say that theoxygen is acting as an oxidising agent. The oxygen is receiving electrons from the iron. Anoxidising agent is therefore an electron acceptor.

Elements with high electronegativities tend to form ions by gaining electrons and so act as oxidisingagents.

In a similar way, the oxygen molecules are being reduced by the iron. The iron is acting as areducing agent. The iron is donating electrons to the oxygen. A reducing agent is therefore anelectron donor.

Elements with low electronegativities tend to form ions by losing electrons and so act as reducingagents. In the Periodic Table, the strongest reducing agents are in Group 1, and the strongestoxidising agents are in Group 7.



Sometimes there are ions present which do not react in a redox reaction. These ions are present inthe reactants and are still present in the same amount when the products have formed. Such ionsare known as spectator ions and are not normally shown in the balanced redox equation.

The electrochemical series (also known as the table of standard reduction potentials) can befoundon Page 12 of the SQA Higher Chemistry Data Booklet and can give an indication of thestrength of oxidising and reducing agents. The easiest reductions occur at the foot of the series, leftto right. The best oxidising agents are therefore found at the bottom left side. Similarly, the easiestoxidations occur at the top of the series, right to left . This is summarised as follows.

Reducing and oxidising agents

Example exercise

When bromine is added to potassium iodide solution, iodine is displaced and potassium bromidesolution is formed.

2KI + Br2 → 2KBr + I2

Potassium bromide formation

The two potassium ions are unchanged in this reaction. These are spectator ions; ignoring theseions results in the ion electron equations as follows.

2I- → I2 + 2e-

Br2 + 2e- → 2Br-

Iodide and bromine half equations



Q4: Is the first ion-electron equation (see preceding iodide and bromine half equations) anoxidation or a reduction?

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Q5: Which of these is acting as an oxidising agent?

a) Iodideb) Iodinec) Bromined) Bromide

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Q6: Name the chemical species acting as a reducing agent.

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Q7: Where in the electrochemical series in the SQA Higher Chemistry Data Booklet will youfind the reaction which is most likely to go in the opposite direction?

a) Top of the table.b) Bottom of the table.

Use the following equation showing a redox change in the following three questions.

Mg + S → Mg2+S2-

Q8: Name the reactant being reduced.

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Q9: Name the reactant acting as an electron donor.

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Q10: Name the oxidising agent.

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Q11: Which of the following is the best oxidising agent?

a) Na (s)b) Br2 (l)c) Sn2+ (aq)d) I- (aq)e) Zn (s)f) Cu2+ (aq)

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Q12: Which of the following is the best reducing agent?

a) Na (s)b) Br2 (l)



c) Sn2+ (aq)d) I- (aq)e) Zn (s)f) Cu2+ (aq)

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Q13: Which substance could be used as an oxidising agent or a reducing agent?

a) Na (s)b) Br2 (l)c) Sn2+ (aq)d) I- (aq)e) Zn (s)f) Cu2+ (aq)

6.3.1 Displacement reactions

Go onlineDisplacement reactions

A strip of zinc metal is placed into copper(II) sulfate solution. The zinc ionises and passesinto solution, and copper ions plate onto the surface of the strip as copper atoms, replacingthe zinc. The solution loses its blue colour since the resulting zinc sulfate solution iscolourless.

Cu2+

Cu2+

Cu2+

Cu2+

Cu2+

CuZn2+

Zn

ZnZn

ZnZn

SO2+4

SO2+4

SO2+4

SO2+4

SO2+4

SO2+4

Cu2+

Cu2+

Cu2+

CuZn2+Zn

Cu Zn

Zn Cu

SO2+4

SO2+4

SO2+4

SO2+4

SO2+4

SO2+4

Zn2+

Zn2+

Copper Sulfate solutionZinc metal

Zinc Sulfate solution

Zinc metal

Copper metal



Q14: Name the spectator ion in this reaction.

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Q15: Complete the ion-electron half equation for the oxidation taking place:

Zn →. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q16: Complete the ion-electron half equation for the reduction taking place:

→ Cu

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Q17: Which of these is acting as an oxidising agent?

a) Zinc atomsb) Zinc ionsc) Copper atomsd) Copper ions

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Q18: Name the reactant acting as an electron donor.

a) Zinc atomsb) Zinc ionsc) Copper atomsd) Copper ions

Key point

An oxidising agent is a substance which accepts electrons; a reducing agent is a substancewhich donates electrons. Oxidising and reducing agents can be identified in redox reactions.

6.4 Molecules and group ions as oxidising and reducing agents

You might have noticed that as well as elements being listed in the table of standard reductionpotentials, there are some compounds. Hydrogen peroxide (H2O2) is a molecule that is an oxidisingagent found in bleach. Both dichromate (Cr2O7

2-) and permanganate (MnO4-) ions are group ions

that are strong oxidising agents in acidic solutions. Carbon monoxide (CO) is a gas that can be usedas a reducing agent in a blast furnace (National 5). You will study the reactions of permanganateand dichromate ions in more detail in a later section.



Q19: Can you think of a practical reason why sodium dichromate or potassiumpermanganate are frequently used as oxidising agents in chemistry as opposed to theelements studied so far?

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Q20: A traditional method for producing chlorine is the reaction between concentratedhydrochloric acid and potassium permanganate.

The formulae for the reactants and products are:

KMnO4 + HCl → KCl + MnCl2 + H2O + Cl2

It is quite difficult, but can you balance this equation?

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Q21: Remembering that KMnO4 consists of K+ and MnO4- ions, and HCl is also ionised in

solution, what is the oxidising agent in this reaction and what are the reducing agents?

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Q22: Look carefully at the reactants and products again. What are the products from theseoxidising and reducing agents?

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Q23: Finally, are there any spectator ions in this reaction?

6.4.1 Oxygen

Almost all life on Earth depends on oxygen to provide energy. When the Earth was formed 4·5 billion(4·5 x 109) years ago* it was a very different place.

A volcano

As the hot fireball cooled, it formed a solid crust. Volcanoes belched out gases: steam, carbondioxide and ammonia. These formed the oceans and an atmosphere with no oxygen.



Planet Earth

We now have an atmosphere consisting of about 21% oxygen. The rest is mainly nitrogen.

[* Evidence for the age of the Earth comes mainly from studying the ages of rocks using a variety ofradioactive decay data.]

Cyanobacteria (blue-green algae)

The change was brought about somewhere around 2·7 to 2·2 billion years ago by vast numbers ofcyanobacteria (blue-green algae) in the oceans.

These organisms used the energy of sunlight in photosynthesis to convert the abundant suppliesof carbon dioxide and water into carbohydrates and oxygen. This dramatic change in atmosphericcomposition fuelled the abundant numbers of plant and animal species which now use oxygen toprovide energy for life.



Q24: A simple equation for photosynthesis, making glucose, C6H12O6, is:

6CO2+ 6H2O + light energy → C6H12O6 + 6O2

Can you see in this case what is oxidised and what is reduced?

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q25: When we humans use glucose to provide energy for life, the equation is simplyreversed:

C6H12O6 + 6O2 → 6CO2+ 6H2O + energy

The atmospheric concentration of carbon dioxide measured at Mauna Loa, Hawaii, has risenfrom 315 ppmv in 1960 to 390 ppmv in 2010; ppmv is parts per million by volume; μ� CO2

per � atmosphere. Assuming the 'effective volume' of Earth's atmosphere is 4 x 1020 �, whatvolume of additional carbon dioxide has the atmosphere contained in going from 1960 to2010?

6.5 Uses for strong oxidising agents

The main domestic and industrial uses of strong oxidising agents are as biocides and bleaches.

Sodium chlorate(l) solution is apowerful germicide and bleach.

This weedkiller containssodium chlorate(lll)

The original washingdetergents contained

perborates, percarbonates andsilicates, from which Persil

derived its name.



Biocides

Biocides are used for killing fungi and bacteria, and inactivating viruses. The main oxidising agentsused are chlorine, hypochlorite, chlorine dioxide, iodine and permanganate.

Disinfection during a foot-and-mouth outbreakusing hypochlorite

Disinfection with iodine during a bone marrowdonation

Bleaches

These are used for bleaching textiles, paper and hair. Common oxidising agents include sulfite,chlorine and peroxides.

Perborate is used for adding extra whiteness to laundry and dibenzoyl peroxide for treating spotsand acne.



The original washing detergents containedperborates, percarbonates and silicates, from

which Persil derived its name.Gel for treating acne

Water purification

Water from reservoirs is purified considerably before it is pumped into buildings for consumption.As well as removing particulate matter, there are possible pathogens to be removed, comprisingviruses, bacteria, including Escherichia coli, Campylobacter and Shigella, and protozoa, includingGiardia lamblia and other cryptosporidia.

The most common disinfection method involves some form of chlorine or its compounds suchas chloramine or chlorine dioxide. Chlorine dioxide is a faster-acting disinfectant than elementalchlorine; however, it is relatively rarely used because, in some circumstances, it may createexcessive amounts of chlorite. The use of chloramine is becoming more common as a disinfectant.Although chloramine is not as strong an oxidant, it does provide a longer-lasting residual than freechlorine and it won't form undesirable by-products. Ozone (O3) is an unstable molecule which readilygives up one atom of oxygen, providing a powerful oxidizing agent which is toxic to most waterborneorganisms. It is a very strong, broad spectrum disinfectant that is widely used in Europe.

For disinfection of small quantities of water, when camping for example, tablets which producehypochlorite (chlorate(I)) in situ can be used.



Propellants

A further use for strong oxidants is as propellants for space rockets. These include hydrogenperoxide and liquid oxygen.

The Apollo-Saturn V first stage used kerosene-liquid oxygen, the upper stages used liquidhydrogen-liquid oxygen.

6.6 Ion-electron half equations

When sodium burns in chlorine the product, sodium chloride, is an ionic compound.

2Na + Cl2 → 2Na+Cl-

The reaction is a redox reaction and can be described by writing two ion-electron equations. Thefirst shows the sodium.

2Na → 2Na+

Sodium ionising

This ion-electron equation can have its charge balanced by showing the two electrons lost from thesodium atom on the right hand side of the equation.

2Na → 2Na+ + 2e- (oxidation)

Sodium ionising (2)

Similarly, the chlorine ion-electron half equation can be balanced by showing the gain of twoelectrons by the chlorine on the left of the equation.



Cl2 → 2Cl-

Cl2 + 2e- → 2Cl- (reduction)

Chlorine half equations

Practise writing ion-electron equations in the following activity.

Go onlineWriting ion-electron equations

Complete the following oxidation and reduction ion electron half equations and then checkyour answers.

Q26: Oxidation of potassium, forming a K+ ion:

K →. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q27: Oxidation of magnesium:

Mg →. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q28: Oxidation of aluminium:

Al →. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q29: Reduction of a fluorine atom to a fluoride ion:

→ F-

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q30: Reduction of a nitrogen atom to nitride ion:

→ N3-

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q31: Reduction of an oxygen atom to oxide ion:

→ O2-

Many ion-electron equations do not have to be worked out as they can be found in a table knownas the 'electrochemical series: standard reduction potentials'. This is available in most chemistrytexts or in the SQA Higher Chemistry Data Booklet.

The oxidation ion-electron equation which applies can be obtained by reversing the reductionequation given in the table.



Simple rules

In a redox reaction, a simple rule applies when determining which equation will be the reductionand which the oxidation.

The ion-electron equation appearing lowest in the electrochemical series will go as written (i.e. asa reduction) and this reaction will be able to force any ion-electron equation higher up in the tableto go in reverse (i.e. as an oxidation). This can be described as the 'anticlockwise rule' as follows.

Anticlockwise rule diagram

Use the diagram and the electrochemical series on Page 12 of the SQA Higher Chemistry DataBooklet to answer the following questions.

Q32: In the electrochemical series shown, where will the easiest reduction be found?

a) At the top of the table.b) At the bottom of the table.

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q33: In the electrochemical series shown, where will the easiest oxidation be found?

a) At the top of the table, going 'as written'.b) At the bottom of the table, going 'as written'.c) At the top of the table, in reverse.d) At the bottom of the table, in reverse.

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q34: Which of the following reductions is the easiest?

a) Na+(aq) + e- → Na(s)b) F2(g) + 2e- → 2F-(aq)c) Br2(g) + 2e- → 2Br-(aq)d) Cu2+(aq) + 2e- → Cu(s)

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .



Q35: Which of these reactions would be reversed by the following reduction?

2H+(aq) + 2e- → H2(g)

a) Na+(aq) + e- → Na(s)b) F2(g) + 2e- → 2F-(aq)c) Br2(g) + 2e- → 2Br-(aq)d) Cu2+(aq) + 2e- → Cu(s)

Key point

Ion-electron equations can be written for oxidation and reduction reactions. Many ion-electronequations can be obtained from the electrochemical series available on Page 12 of the SQAHigher Chemistry Data Booklet.

6.7 Combining ion-electron equations

The balanced redox equation for a chemical change can be worked out from the ion-electronequations which make up the oxidation and reduction stages. Often these can be found in the tableof standard reduction potentials: the electrochemical series. This is available in most chemistrytexts or on Page 12 of the SQA Higher Chemistry Data Booklet.

The oxidation ion-electron equation which applies can be obtained by reversing the reductionequation given in the table.

Writing the balanced redox equation can best be illustrated by use of an example.

Example : Sodium reacting with chlorine

When sodium is placed into chlorine gas, the product, sodium chloride, is an ionic compound.Write a balanced redox equation from the two ion-electron equations.

Consider the two reactants (sodium and chlorine). The ion-electron equation appearinglowest in the electrochemical series will go as written (i.e. as a reduction). This is describedin the 'anticlockwise rule' as mentioned previously.

Cl2 + 2e- → 2Cl- (reduction)

Chlorine to chloride

In this reaction, the chlorine is the oxidising agent and will be able to force any of the ion-electron equations higher up in the table to go in reverse (i.e. as an oxidation). In this casesodium will be oxidised.

Na → Na+ + e- (oxidation)

Sodium oxidation



Combining the individual ion-electron equations into the redox equations requires a numberof 'balancing' stages.

Go onlineBalancing redox equations

The following diagram shows the stages required.

Balancing redox equations

Go onlineTutorial - simple ion-electron equations

Use the electrochemical series to try these examples. In each case:

• write the ion-electron equation for the reduction;

• write the ion-electron equation for the oxidation;

• combine these to form the redox equation.

Q36: When magnesium is placed into chlorine gas, the product, magnesium chloride, is anionic compound.

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .



Q37: When chlorine reacts with a solution containing iodide ions, iodine is one of theproducts.

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q38: Magnesium reacts with sulfuric acid to give hydrogen gas and magnesium sulfatesolution. The spectator ions can be ignored when writing the redox equation.

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q39: Aluminium displaces silver ions from a solution of silver(I) nitrate, giving aluminiumnitrate solution and silver. The spectator ions can be ignored when writing the redox equation.

Key point

Ion-electron equations can be combined to produce redox equations.

6.8 Complex ion-electron equations

Many of the more complex ion-electron equations involve hydrogen ions and/or water moleculesin the chemical change. This is why reagents like 'acidified' potassium permanganate or 'acidified'potassium dichromate (as follows) require an acid solution in which to act in redox reactions.

MnO4- + 8H+ + 5e- → Mn2+ + 4H2O

Cr2O72- + 14H+ + 6e- → 2Cr3+ + 7H2O

Complex ion-electron equations

Not all of these more complex ion-electron equations are listed on Page 12 of the SQA HigherChemistry Data Booklet but, given the reactant and product species, ion-electron equationsinvolving H+(aq) and H2O(l) can be written.

Example

Acidified dichromate ions, Cr2O72-, can be used to oxidise Fe2+ ions. The Fe2+ ions are

changed to Fe3+ ions and the dichromate to chromium(III), Cr3+.

Write ion-electron equations for each reaction and combine these to give a redox equationfor the change.



Go onlineComplex equations

Work out the balanced ion-electron equation for this next example using the steps as followsand then answer the questions.

Complex equations

The oxidation step involves the Fe2+ ions being changed into Fe3+ ions.

Fe2+ → Fe3+ + e- (oxidation)

The other steps in the diagram are:

1. The reduction given in the question can be written and the number of atoms in thegiven ions balanced.

2. Balance the number of oxygens by adding the correct number of water molecules.

3. Balance the number of hydrogen atoms by adding the correct number of hydrogenions.

4. Balance the charge on each side of the equation by adding electrons to the mostpositive side.

Writing the full redox equation requires the oxidation and reduction ion-electron equations tobe combined.



Balance the number of electrons being transferred between the ion-electron equations andadd them to give the redox equation.

Example

Fe2+ ions can be changed into Fe3+ ions in reaction with hypochlorite ions. The hypochloriteion is converted to chloride (Figure 6.1).

Figure 6.1: hypochlorite reduction

ClO- → Cl-

Write a balanced ion-electron equation for this change (Figure 6.1) using the steps shownpreviously and then answer the questions.

Q40: How many water molecules have to be added to the product side?

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q41: How many hydrogen ions have to be added to the reactant side?

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q42: How many electrons have to be added to balance the charge?

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q43: On which side do the electrons get added?

a) Reactantb) Product

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q44: Is the completed ion-electron equation an oxidation or a reduction?

a) Oxidationb) Reduction

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .



Q45: Write a complete balanced redox equation for the reaction between hypochlorite andFe2+.

Go onlineTutorial - complex ion-electron equations

In the first example, write a balanced ion-electron equation for the given change and thenanswer the questions which take the process step by step.

Example

V3+ ions can be changed into VO3- ions as part of a redox reaction. Write a balancedion-electron equation for the change and then answer the questions which take the processstep by step.

V3+ → VO3-

Q46: How many water molecules have to be added to the reactant side?

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q47: How many hydrogen ions have to be added to the product side?

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .


. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .



. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q50: Is the completed ion-electron equation an oxidation or a reduction?

a) Oxidationb) Reduction

In the following example, write a balanced ion-electron equation for the given change andthen check your answer.

Q51: ClO3- ions can be changed into Cl2 molecules as part of a redox reaction. Write a

balanced ion-electron equation for the change. (Hint: balance the number of chlorine atomsfirst).

ClO3- → Cl2



Use the electrochemical series on Page 12 of the SQA Higher Chemistry Data Booklet tohelp with the questions. In each case:

• write the ion-electron equation for the reduction;

• write the ion-electron equation for the oxidation;

• combine these to form the redox equation.

Q52: When zinc is placed into sulfuric acid, the products are hydrogen gas and zinc sulfatesolution.

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q53: In acid solution, potassium permanganate oxidises iron(II) to iron(III).

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q54: In dilute nitric acid solution, copper metal is oxidised to copper(II) ions. The nitrate ionis reduced to nitrogen monoxide gas.

Key point

Given reactant and product species, ion-electron equations which include H+(aq) and H2O(l)can be combined to produce redox equations.

6.9 Summary exercise

Go onlineSummary exercise

Q55: The thermite reaction is used to generate molten iron, for example to weld railway linesin situ.

The reactants are iron(iii) oxide and aluminium powder. The equation for the reaction is:

a) Fe2O3 + Al → Fe + Al2O3

b) Fe2O3 + 2Al → 2Fe + Al2O3

c) 3FeO + 2Al →3 Fe + Al2O3

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q56: What word can be used to describe the oxide ion (O2-)?

a) Commonb) Watchingc) Spectator

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .



Q57: The ionic equation can be written as:

a) Fe + Al → Fe + Alb) Fe+ + Al → Fe + Al+

c) Fe3+ + Al → Fe + Al3+

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q58: Oxidation is defined as �� of electrons.

a) changeb) lossc) gain

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q59: What is oxidised in the thermite reaction ?

a) Alb) Fec) Al3+

d) Fe3+

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q60: What is acting as the oxidising agent?

a) Alb) Fec) Al3+

d) Fe3+

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q61: Using the electrochemical series, which metal could not be used to reduce Fe3+ ion toFe?

a) Copperb) Magnesiumc) Manganese

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q62: You are required to make a mixture of aluminium (Al) and iron(III) oxide (Fe2O3) to usefor this reaction. Use the SQA Higher Chemistry Data Booklet to calculate how much Fe2O3

is required to add to 1 kg of Al to make a reaction mixture.

a) 1 kgb) 2 kgc) 3 kg



6.10 Summary

Summary

Elements as oxidising or reducing agents

• A redox reaction is a reaction in which reduction and oxidation occur together, reductionbeing the gain of electrons by a reactant and oxidation being the loss of electrons by areactant in a reaction.

• An oxidising agent is a substance which accepts electrons.

• A reducing agent is a substance which donates electrons.

• Oxidising and reducing agents can be identified in redox reactions.

• Elements with low electronegativities (metals) tend to form ions by losing electrons(oxidation) and so can act as reducing agents.

• Elements with high electronegativities (non-metals) tend to form ions by gainingelectrons (reduction) and so can act as oxidising agents.

• The strongest reducing agents are found in Group 1.

• The strongest oxidising agents come from Group 7.

• The electrochemical series indicates the effectiveness of oxidising and reducing agents.

Compounds as oxidising or reducing agents

• Compounds can also act as oxidising or reducing agents.

• The electrochemical series represents a series of reduction reactions.

• The strongest oxidising agents are at the bottom of the left-hand column of theelectrochemical series.

• The strongest reducing agents are at the top of the right-hand column of theelectrochemical series.

• Electrochemical series contain a number of ions and molecules.

• The dichromate and permanganate ions are strong oxidising agents in acidic solutions,while hydrogen peroxide is an example of a molecule which is a strong oxidising agent.

• Carbon monoxide is an example of a gas that can be used as a reducing agent.

• Oxidising and reducing agents can be selected using an electrochemical series fromPage 12 of the SQA Higher Chemistry Data Booklet or can be identified in the equationshowing a redox reaction.

Use of oxidising agents

• Oxidising agents are widely employed because of the effectiveness with which they cankill fungi and bacteria, and can inactivate viruses.



Summary continued

• The oxidation process is also an effective means of breaking down coloured compoundsmaking oxidising agents ideal for use as 'bleach' for clothes and hair.


• Oxidation and reduction reactions can be represented by ion-electron equations.

• When molecules or group ions are involved, if the reactant and product species areknown, a balanced ion-electron equation can be written by adding appropriate numbersof water molecules, hydrogen ions and electrons.

• Ion-electron equations can be combined to produce redox equations.

Practical applications

• Displacement reactions are example of redox reactions and oxidising and reducingagents can be identified in these and other redox reactions.

6.11 Resources

Texts










This end of topic test is available online. If you do not have access to the internet, here is apaper version.

Q63: Which of the following reductions is the easiest?

a) Cl2 + 2e- → 2Cl-

b) 2H+ + 2e- → H2

c) Fe3+ + e- → Fe2+

d) Na+ + e- → Na

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q64: An oxidising agent is a substance which

a) is oxidised in a reaction.b) accepts electrons.c) donates electrons.d) causes reduction.

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q65: Which of these is the equation for oxidation of bromide ions?

a) 2Br- + 2e- → Br2

b) 2Br- → Br2 + 2e-

c) Br2 + 2e- → 2Br-

d) Br2 → 2Br- + 2e-

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q66: In which of these reactions is the reactant a reducing agent?

a) Fe3+ + 3e- → Feb) Fe2+ → Fe3+ + e-

c) H+ + HO- → H2Od) 2H2O + 2e- → H2 + 2OH-

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q67: Which of these is the equation for the reduction of copper ion?

a) Cu2+ + 2e- → Cub) Cu2+ → Cu + 2e-

c) Cu + 2e- → Cu2+

d) Cu → Cu2+ + 2e-



Q68: Consider the following reactions.

Cl2 + 2e- → 2Cl-

I2 + 2e- → 2I-

Which of these reactions will take place in solution?

a) Chlorine will react with iodine.b) Iodide ions will react with chloride ions.c) Iodide ions will react with chlorine.d) Chloride ions will react with iodine.

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q69: The following reduction fits into the electrochemical series: standard reductionpotentials at a value of Eo = +0·68 V.

O2 + 2H+ + 2e- � H2O2

Which of these chemicals would convert H2O2 into O2?

a) Sn4+

b) Fe2+

c) Sn2+

d) Fe3+

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q70: How many electrons must be added to balance this incomplete ion-electron equation?

Hg2Cl2 → 2Hg + 2Cl-

a) Add one electron to the left hand side.b) Add two electrons to the right hand side.c) Add two electrons to the left hand side.d) Add one electron to the right hand side.

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q71: Consider the following oxidations.

Fe2+ → Fe3+ + e-

Sn2+ → Sn4+ + 2e-

Which of these reactions is likely to take place in solution?

a) Tin(IV) ions will react with iron(II) ions.b) Tin(IV) ions will react with iron(III) ions.c) Iron(II) ions will react with tin(II) ions.d) Iron(III) ions will react with tin(II) ions.



. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q72: In the following equation, highlight the chemical acting as a reducing agent.

CrO72- + 14H+ + 6I- → 2Cr3+ + 7H2O + 3I2

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q73: In the following equation, highlight the chemicals which act as an oxidising agent.

CrO72- + 14H+ + 6I- → 2Cr3+ + 7H2O + 3I2

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q74: Which of the following contains the best oxidising agent?

a) Cs+(aq)b) Br-(aq)c) Ca (s)d) Cl2(g)e) Fe2+(aq)f) Sn (s)

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q75: Which of the following contains the best reducing agent?


. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q76: Which of the following contains a substance which could be used as an oxidising agentor a reducing agent?




Look at the following incomplete ion-electron equation.

XeO3 → Xe

Q77: How many water molecules have to be added to the product side?

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q78: How many hydrogen ions have to be added to the reactant side?

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .


. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .



. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q81: The completed ion-electron equation is

a) an oxidation.b) a reduction.


153

Topic 7

End of unit test

154 TOPIC 7. END OF UNIT TEST

Go onlineEnd of Unit 1 test

Q1: On going down a group in the Periodic Table, the first ionisation energy:

a) increases.b) decreases.

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q2: What ion is oxidised in the following redox reaction?

HgCl2(aq) + SnCl2(aq) → Hg(l) + SnCl4(aq)

a) Cl-(aq)b) Sn2+(aq)c) Sn4+(aq)d) Hg2+(aq)

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q3: Ionic bonding is most likely when the electronegativity difference between the elementsis:

a) exothermic.b) endothermic.c) large.d) small.

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q4: Which of the following exist as diatomic molecules?

a) Ethaneb) Potassium bromidec) Carbon monoxided) Neon

. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Q5: When an atom X of an element in Group 2 reacts to become X2+:

a) the mass number of X increases.b) the charge on the nucleus increases.c) the atomic number of X decreases.d) the number of occupied energy levels decreases.

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Q6: Which of the following reactions can be classified as reduction?

a) CH3CH2COCH3 → CH3CH2CH(OH)CH3

b) CH3CH(OH)CH3 → CH3COCH3

c) CH3CH2OH → CH3COOHd) CH3CH2CHO → CH3CH2COOH


TOPIC 7. END OF UNIT TEST 155

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Q7: Which of the following is a redox reaction?

a) ZnO + 2HCl → ZnCl2 + H2Ob) ZnCO3 + 2HCl → ZnCl2 + H2O + CO2

c) Zn(OH)2 + 2HCl → ZnCl2 + 2H2Od) Zn + 2HCl → ZnCl2 + H2

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Q8: In which of the following reactions is a positive ion reduced?

a) Sulfate → sulfiteb) Gold(II) → gold(III)c) Iron(III) → iron(II)d) Bromide → bromine

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Q9: Which of the following molecules may be described as polar?

a) C1 Be C1b) H C1c)

C1 C1d) C1

CC1 C1

C1

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Q10: The melting points of Group 7 elements increase on descending the group because the�� increase.

a) mean bond energiesb) nuclear chargesc) covalent bond lengthsd) London dispersion forces

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Q11: The difference between the covalent radii of lithium and carbon is mainly due to thedifference in the:

a) number of electrons.b) number of protons.c) mass of each atom.d) number of neurons.

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Q12: In general, covalent substances have lower melting points than ionic substancesbecause:

a) covalent bonds have no electrostatic attractive forces.b) bonds between molecules are weaker than bonds between ions.c) covalent compounds are composed of non-metals which have low melting points.d) ionic bonds are stronger than covalent bonds.

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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.

The apparatus shown can be used to prepare iron(III) chloride.

Q14: Which substance in the diagram contains metallic bonding?

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Q15: By considering the method of production, predict the type of bonding in iron(III) chloride.

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Q16: Which of the following describes the bonding in the chlorides across Period 3 from NaClto SCl2?

a) Ionic ⇒ pure covalentb) Ionic ⇒ polar covalent ⇒ pure covalentc) Polar covalent ⇒ ionicd) Ionic ⇒ pure covalent ⇒ polar covalent



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Q17: Atoms of different elements have different attractions for bonded electrons. What termis used as a measure of these differing attractions?

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Q18: In which of the following reactions is the hydrogen ion acting as an oxidising agent?

a) Fe + 2HCl → FeCl2 + H2

b) KOH + HNO3 → KNO3 + H2Oc) MgCO3 + H2SO4 → MgSO4 + H2O + CO2

d) CH3COONa + HCl → NaCl + CH3COOH

The water in swimming pools is often disinfected by adding 'chlorinated lime'. Thisproduces chlorate(I) ions (OCl-) in solution, which effectively kill pathogens as long as theirconcentration is sufficiently high.

One method of estimating the OCl- concentration is to take a known volume of pool waterand add an excess of potassium iodide solution. The reactions given in the following twoquestions produce iodine, which can then be measured.

Q19: Fill in the gap to complete the ion-electron equation.

OCl- + � H+ + 2e- → Cl- + H2O

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Q20: Fill in the gap to complete the ion-electron equation.

2I- → I2 + �e-

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Q21: Which ionic equation is correct for the whole reaction?

a) OCl- + 2H+ + 2I- → Cl- + H2O + I2b) OCl- + H+ + 2I- → Cl- + H2O + I2c) OCl- + 2H+ + I- → Cl- + H2O + I2d) 2OCl- + 2H+ + I- → 2Cl- + H2O + I2

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Q22: If 100 cm3 of the pool water produced 0·002 moles of iodine, what was theconcentration of OCl- in mol �-1?

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Q23: In order to be effective, the concentration of chlorate(I) should be between 0·4 and 1·5mg �-1. What is the concentration of OCl- in g �-1?

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Q24: Using the table of standard reduction potentials, which of the following redox reactionswill take place?

a) Cl2(g) + 2Br-(aq) → 2Cl-(aq) + Br2(aq)

b) Br2(l) + 2Fe2+(aq) → 2Br-(aq) + 2Fe3+(aq)

c) Ca2+(aq) + Cu(s) → Ca(s) + Cu2+(aq)

d) Zn2+(aq) + Cu(s) → Zn(s) + Cu2+(aq)

e) Mg2+(aq) + H2(g) → Mg(s) + 2H+(aq)

f) Cu(s) + Cu2+(aq) → 2Cu+(aq)

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Q25: The iodate ion, IO3-, can be converted to iodine. Which is the correct ion-electron

equation for the reaction?

a) 2IO3-(aq) + 12H+(aq) + 10e- → I2(aq) + 6H2O(l)

b) 2IO3-(aq) + 12H+(aq) → 2I-(aq) + 6H2O(l)

c) 2IO3-(aq) + 12H+(aq) → 2I-(aq) + 6H2O(l)

d) IO3-(aq) + 6H+(aq) → I-(aq) + H2O(l)

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Q26: A major source of iodine is caliche, mined in Chile.

The mass of iodine in a 10 g sample of caliche can be determined by dissolving the samplein water and adding hydrogen peroxide solution to oxidise the iodide to iodine molecules. Theion-electron equation for the reduction reaction is shown as follows.

H2O2(aq) + 2H+ + 2e- → 2H2O(l)

Write a balanced redox equation for the reaction of hydrogen peroxide with iodide ion.

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Q27: Using starch solution as indicator, the iodine solution is then titrated with sodiumthiosulfate solution to determine the mass of iodine in the sample. The balanced equationfor the reaction is shown as follows.

2Na2S2O3(aq) + I2(aq) → 2NaI(aq) + Na2S4O6(aq)

In an analysis of a sample, 13·25 cm3 of 0·0100 mol �-1 sodium thiosulfate solution wasrequired to reach the end-point.

Calculate the mass of iodine present in the sample of caliche.

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Q28: Hydrogen peroxide is used in gels to whiten teeth. The ion-electron equation for theoxidation of hydrogen peroxide is as follows:

H2O2 → O2 + 2H+ + 2e-

Using your knowledge of chemistry, comment on possible methods for measuring andcomparing the concentration of hydrogen peroxide present in two different gels.

(3 marks)

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Q29: Concentrated solutions of hydrogen peroxide are used in the propulsion systems oftorpedoes. Hydrogen peroxide decomposes naturally to form water and oxygen:

2H2O2(aq) → 2H2O(�) + O2(g) ΔH = -196·4 kJ mol-1

Transition metal oxides act as catalysts in the decomposition of the hydrogen peroxide.Unfortunately, there are hazards associated with the use of hydrogen peroxide as a fuel intorpedoes. It is possible that a leak of hydrogen peroxide solution from a rusty torpedo maytrigger an explosion.

Using your knowledge of chemistry, comment on why this could happen.

(3 marks)

Q30: Match the following chemical terms to the correct definitions.

DefinitionChemical term(randomised)

Regular arrangement of positively charged ions surrounded bydelocalised electrons.

Covalent bonding

The electrostatic force of attraction of a shared pair of electronsfor two positive nuclei.

Metallic bonding

The electrostatic attraction between negative and positive ions. Intramolecular forces

Weak bonds between molecules. Dipole

A slight negative or positive charge caused by the unevendistribution of electrons.


Bonds holding atoms together within molecules. Ionic bonding

(6 marks)

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Q31: Match the following chemical terms to the correct definitions.

DefinitionChemical term(randomised)

Half the distance between the nuclei of two covalently bonded atomsof an element.

Ionisation energy

The strength of the attraction of an element for the electrons of itsbonding electrons. Covalent radius

The energy required to remove an electron from a gaseous atom toform 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

(5 marks)

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Q32: 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

(4 marks)

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Q33: 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 across a period, electronegativity . . . increases / decreases.

• Going down a group, electronegativity . . . increases / decreases.

• Going down Group 7, boiling point . . . increases / decreases.

(5 marks)




Q34: The elements in the third row of the Periodic Table are as follows.

Na Mg Al Si P S Cl Ar

Why does the atomic size decrease crossing the period from sodium to argon?

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Q35: Use a table of these values from Page 11 of the SQA Higher Chemistry Data Bookletto explain why nitrogen chloride contains pure covalent bonds.

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Q36: Atoms of different elements have different ionisation energies.

Explain clearly why the first ionisation energy of sodium is less than the first ionisation energyof lithium.

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Q37: 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 anelement combines with one mole of electrons to form gaseous negative ions.

Write the equation, showing state symbols, that represents the electron affinity of chlorine.

A student writes the following two statements. Both are incorrect. In each case, explain themistake in the student's reasoning.

Q38: Statement 1

All ionic compounds are solids at room temperature. Many covalent compounds are gases atroom temperature. This proves that ionic bonds are stronger than covalent bonds.

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Q39: Statement 2

The formula for magnesium chloride is MgCl2 because, in solid magnesium chloride, eachmagnesium ion is bonded to chloride ions.

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Q40: Although propane and ethanol have similar molecular masses, the alkane is a gasat room temperature while the alcohol is a liquid. Explain why propane is a gas at roomtemperature whereas ethanol is a liquid.

(3 marks)


162 GLOSSARY

Glossary

Allotropes

one of two or more existing forms of an element, e.g. graphite and diamond are allotropes ofcarbon

Bonding electrons

shared pairs of electrons from both atoms forming a 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

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 andare shared amongst the entire structure

Diatomic

molecules with only two atoms, e.g. oxygen, O2, and carbon monoxide, CO

Dipole

an atom or molecule in which a concentration of positive charges is separated from aconcentration of negative charge

Electronegativity

a measure of the attraction that an atom involved in a bond has for the electrons of the bond

Fullerenes

molecules of pure carbon constructed from 5- and 6-membered rings combined into hollowstructures; the most stable contains 60 carbon atoms in a shape resembling a football

Group ions

ions which contain more than one type of atom

Hydrogen bonds

electrostatic forces of attraction between molecules containing a hydrogen atom bonded to anatom of a strongly electronegative element, such as fluorine, oxygen or nitrogen, and a highlyelectronegative atom on a neighbouring molecule


forces of attraction which exist between molecules; they are weaker than chemical bonds


GLOSSARY 163

Intramolecular forces

forces of attraction which exist within a molecule

Ion

formed when atoms lose or gain electrons to obtain the stable electron arrangement of a noblegas; In general, metal atoms lose electrons forming positive ions and non-metal atoms gainelectrons forming negative ions


a half-equation, either an oxidation or a reduction, which in combination of the opposite typecan be part of a complete redox equation

Ionic formulae

give the simplest ratio of each type of ion in the substance and can show the charges on eachion, if required

Ionisation energy

the energy required to remove one mole of electrons from one mole of atoms in the gaseousstate

Isoelectronic

molecules which have the same arrangement of electrons, e.g. the noble gas neon, a sodiumion (Na+) and a magnesium ion (Mg2+) 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 dipolesand 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

Oxidation

the loss of electrons by a reactant in any reaction

Oxidising agent

a substance which accepts electrons

Periodicity

the regular recurrence of similar properties when the elements are arranged in order ofincreasing atomic number

Polar covalent bond

a covalent bond between atoms of different electronegativity, which results in an unevendistribution of electrons and a partial charge along the bond


164 GLOSSARY

Properties

physical and chemical characteristics of a substance; these are often a reflection of thechemical bonding and structure of the material

Reducing agent

a substance which donates electrons

Reduction

the gain of electrons by a reactant in any reaction

Superscript

ionic charges must be shown as superscripts, using a small number to the top right after theatomic symbol, e.g. Al3+

Viscosity

the resistance to flow that is exhibited by all liquids


ANSWERS: UNIT 1 TOPIC 1 165

Answers to questions and activities

Topic 1: The Periodic Table

Test your prior knowledge (page 3)

Q1: a) Atomic number

Q2:

• Bromine• Chlorine• Fluorine• Hydrogen• Iodine• Nitrogen• Oxygen

Q3:


Proton +1 1 Nucleus

Neutron 0 1 Nucleus

Electron -1 0Orbiting nucleus or inshells

Q4: b) Period

Q5: b) Group

Newlands' octaves (page 4)

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 eighthelement. If the elements are numbered in order, element 1 has properties in common with element8, 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 noteafter G is A and so on. A to G is said to be one octave. The same applies if the notes are describedas 'doh-ray-me-fah-soh- etc'.

Q10: 3

Q11: Chromium, manganese and iron are metals. The elements above them in Newlands' tableare non-metals.


166 ANSWERS: UNIT 1 TOPIC 1

Mendeleev's Periodic Table (page 6)

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 whichcontains bromine.

Q14: Noble gases

The modern Periodic Table (page 8)

Q15: Atomic number

Melting and boiling points across a Period (page 9)

Q16: Carbon

Q17: Argon

Q18: Helium

Melting and boiling points down a Group (page 10)

Q19: Lithium

Q20: Caesium

Q21: Melting point and boiling point both decrease and so the forces must be getting weaker.

Atomic size (page 11)

Q22: c) The number of electron shells increases.

Q23: d) The number of protons increases.



End of Topic 1 test (page 14)

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 andd) 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.



Topic 2: Bonding and structure


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 21.

Q6:

You have probably thought of gold, nuggets of which are found in rocks and as gold particles in someriver beds. Platinum and other platinum group metals (rhodium, iridium, ruthenium and osmium) arealso found native.

Sulfur is also found as yellow crystals in rocks and around volcanos. Copper is also found inelemental form, as well as in other ores.

Gold Sulfur Native copper


Q7: Niobium is used in steel superalloys which can resist high temperatures and are used in jetengine components and heat shields.



Construction of a sodium atom (page 24)

Q8: Argon

Q9: Gain an outer electron.

Q10: Lose the outer electron.

Q11: A positive sodium ion (Na+)

Q12: Neon

Metallic bonds (page 25)

Q13: The Group 1 metals sodium and potassium are easily cut with a knife (which is made fromhard metals - iron and chromium). Mercury is liquid (melting point -39◦C) and gallium melts in thehand (at 30◦C).

Copper lattice (page 26)

Q14: c) delocalised.

Q15: d) Lattice

Q16: Each aluminium atom will ionise to produce three electrons and a Al3+ ion. Sodium willproduce only one electron (to carry the electric current) from each atom.


Q17: 4

Q18: 1

Q19: b) share, electron

Q20: d) positively charged nuclei and negatively charged shared electrons.

Q21:



Diamond (page 33)

Q22: 4

Q23: a) Single covalent.

Graphite structure (page 34)

Q24: c) stronger.

Q25: 6

Fullerenes (page 36)

Q26: Diamond and graphite are network covalently bonded so that it would require a very largeinput of energy to separate the atoms in a solution. 'Buckyballs' are covalent molecular bonded withstrong bonds binding the C60 atoms into a molecule, but much weaker bonds between the moleculesin the solid. These can be easily broken to form a solution.

Ionic bond formation (page 37)

Q27: 1

Q28: 7

Q29: 8

Q30: Both ions have a stable noble gas electron configuration.


Q31: c) large

Q32: c) positive caesium ions and negative chloride ions.

Q33: c) 3 : 2

Sodium chloride ionic lattice (page 40)

Q34: c) An ionic lattice.



Common group ions: Question (page 42)

Q35:

Group ion name Group ion formula Valency

ammonium NH4+ 1

carbonate CO32- 2

sulfate SO42- 2

phosphate PO43- 3

ethanoate CH3COO- 1

hydrogencarbonate HCO3- 1

Group ion formulae: Questions (page 44)

Q36: CuSO4

Q37: NaOH

Q38: Ba(HSO4)2

Q39: Be(MnO4)2

Group ion ionic formulae: Questions (page 47)

Q40: (Ni2+)SO42-

Q41: K+(MnO4-)

Q42: Ba2+(HCO32-)

Q43: (NH4+)3(PO4

3-)


Q44: a) Calcium sulfide

Q45: a) Sulfur

Q46: b) delocalised electrons.

Q47: b) Metallic

Q48: a) Covalent andc) Made up of discrete molecules.

Q49: a) Boron



Q50: d) Phosphorus andf) Sulfur

Q51: b) Chlorine ande) Sodium

Q52: c) C

Q53: a) Ionic

Q54: f) SiO2 (s)

Q55: a) NH4Cl (s) ande) Na2CO3 (s)

Q56: b) Monatomic

SQA style questions (page 52)

Q57:

Q58: The bonds between the layers of carbon in graphite are comparatively weak, which meansthat the layers can easily slide over each other providing lubrication.

Q59: The carbon atoms in graphite have electrons not involved on the covalent bonds betweenatoms 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 covalentbonds, so there is no delocalised pool of electrons to conduct electricity.

Q60: The nitrogen atoms in the layers have two free electrons which can form strong covalent bondsbetween adjacent layers in boron nitride.



Topic 3: Periodic Table trends


Q1:


Proton +1 1 Nucleus

Neutron 0 1 Nucleus

Electron -1 0Orbiting 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 56)

Q6: b) A decrease

Q7: a) An increase

Q8: The covalent radius is defined as half the distance between the nuclei of bonded atoms. Noblegases do not form bonds because they are so unreactive.

The trends in covalent radius (page 57)

Q9: c) The number of electron shells increases.

Q10: d) The number of protons increases.



Ionisation energies (page 61)

Q11:

Q12: Yes

Q13: The shape of the first part of the graph (from atomic number 3-10) is repeated for elements11-18.

Q14: Noble gases

Q15: Alkali metals

Q16: a) A steady increase.

Q17: b) A steady decrease.

Q18:

MetalFirst ionisation energy / kJ

mol-1

Lithium 520

Sodium 496

Potassium 419

Rubidium 403

Caesium 376



First ionisation energy - Going down a Group (page 62)

Q19:

The nuclear charge increases on going down a group which should make the outer electron moredifficult 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 shouldmake the outer electron easier to remove. (1 mark)

On going down the group, there are more and more inner electrons which prevent the outer electronfrom experiencing the full effect of the nuclear charge. These inner electrons increasingly screenthe outer electron from the nucleus. Consequently, the outer electron becomes easier to remove asthe 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 63)

Q20:

On going from left to right across Period 2, the nuclear charge increases. The electrons are heldmore 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 screeningeffect. Consequently, the increased nuclear charge pulls in the electrons and the covalent radiusdecreases. (1 mark)

Because the outer electrons are closer to the nucleus, they are more strongly held and so the firstionisation 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 65)

Q26: b) Bromine

Q27: c) C-S

Q28: a) Group 3


Q29: 1

Q30: 7



Q31: Noble gases

Q32: The noble gases do not generally form bonds with other elements.


Q33: c) Electronegativity values increase on going from left to right across a Period and decreaseon going down a Group.

Summary: Periodic Table trends (page 67)

Q34:


Q35: c) Mendeleev left gaps for elements which had not yet been discovered.

Q36: c. There is a decrease in first ionisation energy on going down Group 0, andf. 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) + 2e-

Q42: d) F(g) → F+(g) + e-

Q43: 4. The first ionisation energy decreases from Li to Cs, and5. Electronegativity decreases from Li to Cs.

Q44: d) increasing nuclear charge.

Q45: d) 6



Q46: b)

Atomic Number

Q47: d) Atomic mass

Q48: Electronegativity

Q49: Fluorine

Q50: d)

Q51: b) They generally do not form bonds with other elements.


Q52: On going from left to right across Period 2, the nuclear charge increases. The electrons areheld 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 screeningeffect. Consequently, the increased nuclear charge pulls in the electrons and the covalent radiusdecreases. (1 mark)

The outer electrons are more strongly held because they are closer to the nucleus, so the firstionisation increases across the Period. (1 mark)

Q53: The nuclear charge increases on going down a group, which should make the outer electronmore 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 shouldmake the outer electron easier to remove. (1 mark)

On going down the group, there are more and more inner electrons which prevent the outer electronfrom experiencing the full effect of the nuclear charge. These inner electrons increasingly screenthe outer electron from the nucleus. Consequently, the outer electron becomes easier to remove asthe atom gets bigger, i.e. the first ionisation energy decreases on going down a group. (1 mark)



Topic 4: Bonding continuum and polar covalent bonding


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 79)

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 (NH3)

Direction of charge (page 81)

Q11:




Q12:

• Pure covalent - electronegativity difference 0; Br - Br

• Polar covalent - electronegativity difference 0·6; δ+H - Brδ-

• Ionic - electronegativity difference 2·0; K+ ...... Br-

Q13: c) Ionic

Q14: a) Pure covalent

Q15: b) Fluorine

Permanent dipole interactions (page 84)

Q16: B

Q17: C

Q18: B

Predicting molecular polarity (page 86)

Q19:

Detecting polar molecules (page 86)

Q20: a) Water

Q21: a) Water

Q22: b) It is attracted by the rod.

Q23: a) It is not affected by the rod.


Q24: a) Polar




Q25: d) Calcium sulfide

Q26: a) Sulfur

Q27: c) delocalised electrons.

Q28: b) Metallic

Q29: a) Covalent, andb) Made up of discrete molecules.

Q30: a) Boron

Q31: d) Phosphorus, andf) Sulfur

Q32: b) Chlorine, ore) Sodium

Q33: a) Cl-F

Q34: C

Q35: c) BeCl2

Q36: c) Ionic

Q37: f) SiO2 (s)

Q38: a) NH4Cl (s), ande) Na2CO3 (s)

Q39: b) CH3OH (l)



Topic 5: Intermolecular forces


Q1: c) Electronegativity

Q2: a) Fluorine

Q3: c) increases.

Induced dipoles (page 101)

Q4: b) unevenly spread.

Q5: b) a temporary dipole.

Q6: b) The temporary dipole causes an induced dipole.

Q7: b) intermolecular.


Q8: Krypton

Q9: Krypton

Q10: Krypton

Q11: c) The larger the atom the stronger the London dispersion forces.


Q12: The energy required to break forces which are only effective at such low temperatures is verysmall, and in fact, London dispersion forces of attraction are weaker than all other types of bonding.


Q13: As the alkane molecules in the family get bigger, from methane to butane, the boiling pointincreases (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 havestronger London dispersion forces than methane. (1 mark)

Therefore the larger the molecule the stronger the London dispersion forces. (1 mark)

Q14: Butane

Q15: Butane

Q16: Butane

Q17: c) The larger the molecule the stronger the London dispersion forces.



Strength of hydrogen bonds (page 105)

Q18: As size increases, London dispersion forces increase, so one might expect water (thesmallest) to have the lowest boiling point. However, the boiling point of water goes against the trendand is much higher than might be expected. This is because water molecules exhibit hydrogenbonding between the molecules, due to the highly polar character of the water molecule. Althoughthe other hydrides have some permanent dipole-permanent dipole interactions, the hydrogenbonded permanent dipole-permanent dipole interactions in water are stronger than other forms ofpermanent dipole-permanent dipole interactions and London dispersion forces, as evidenced by thehigh boiling point.

Q19: London dispersion forces

Q20: Water

Q21: Water

Q22: b) The hydrogen bonds are stronger than London dispersion forces and the other permanentdipole-permanent dipole interactions.

Q23: a) Covalent bonds are much stronger than hydrogen bonds.


Q24: All of the above.

Hydrides of non-metals (page 108)

Q25: b) No

Q26: NH3, H2O and HF

Q27: b) Van der Waals attractions

Testing viscosity (page 110)

Q28: Methanol

Q29: Methanol

Q30: Glycerol

Q31: 0.5 m s-1

Q32: a) 0·8 m s-1

Q33: 1.25 seconds


Q34: Ammonia is a polar substance which can form hydrogen bonds with water; nitrogen is non-polar and unable to bond to water.



Which cooking method? (page 115)

Q35: Not volatile; soluble in water.

Aspartame has many -OH, -C=O and -NH2 groups which can form relatively strong hydrogen bondsbetween aspartame molecules in the solid (making it non-volatile, with a high boiling point).

These will also bond easily with water molecules making it water soluble.

Q36: Volatile; soluble in oil.

This is a fairly low molecular mass hydrocarbon. It will only have weak London dispersion forcesbetween molecules so will be volatile and soluble in oil.

Q37: Volatile; soluble in water.

This has a low molecular mass so is volatile. It has several polar groups: -OH, -CHO (and - OCH3)so is likely to form hydrogen bonds with water and be fairly water soluble.

Q38: Not volatile; soluble in oil.

This molecule is more difficult. It has some polar groups (-OH, -NH and -C=O), but a benzene ringand a long hydrocarbon chain. It is likely to be non-volatile and soluble in oil.


Q39: d) a temporary dipole causing an induced dipole.

Q40: b) Covalent bonds

Q41: b) The smaller the molecule the weaker the force.

Q42: a) HF

Q43: b) The London dispersion forces become stronger.e) The relative atomic mass increases.

Q44: London dispersion forces.

Q45: Hydrogen bonding (a form of permanent-permanent dipole interaction) between themolecules.

Q46: c) He (g)

Q47: a) NH3 (l)

Q48: a) NH3 (l)f) HBr (l)

Q49: b) CCl4 (l)

Q50: Hydrogen bonding



Q51:



Topic 6: Oxidising or reducing agents


Q1: a) gain of electrons.

Q2: c) Pb2+(aq) + 2NO3-(aq) + 2K+(aq) + 2l-(aq) → Pb2+(l-)2(s) + 2K+(aq) + 2NO3

-(aq)

Q3: e) 5H2O2(l) + 2MnO4-(aq) + 6H+(aq) → 2Mn2+(aq) + 8H2O(l) + 5O2(g)


Q4: Oxidation

Q5: c) Bromine

Q6: Iodide

Q7: a) Top of the table.


Q8: Sulfur

Q9: Magnesium

Q10: Sulfur

Q11: b) Br2 (l)

Q12: a) Na (s)

Q13: c) Sn2+ (aq)

Displacement reactions (page 129)

Q14: Sulfate

Q15: Zn → Zn2+ + 2e-

Q16: Cu2+ + 2e- → Cu

Q17: d) Copper ions

Q18: a) Zinc atoms


Q19: These salts are soluble in water so can be used in reactions in solution - the commonest onesin chemistry.

Q20: 2KMnO4 + 16HCl → 2KCl + 2MnCl2 + 8H2O + 5Cl2



Q21: The permanganate ion (MnO4-) is the oxidising agent and the reducing agents are chloride

ion (Cl-) and hydrogen ion (H+).

Q22: MnO4- is reduced to manganese(II) ion (Mn2+). Chloride ion is oxidised to chlorine and

hydrogen ion oxidised to water.

Q23: K+ is a spectator ion.


Q24: Looking at the reverse reaction, it is much easier to see that glucose is oxidised by oxygen tocarbon dioxide and water. In the case of photosynthesis, the carbon dioxide is reduced to glucose,and the water is oxidised to molecular oxygen.

Q25: Increase in ppmv 390 - 315 = 75 ppmv CO2

Volume of CO2 is: Total volume of atmosphere x ppmv x 10-6

4 x 1020 x 75 x 10-6 = 3 x 1016 litres of CO2

Writing ion-electron equations (page 137)

Q26: K → K+ + 1e-

Q27: Mg → Mg2+ + 2e-

Q28: Al → Al3+ + 3e-

Q29: F + 1e- → F-

Q30: N + 3e- → N3-

Q31: O + 2e- → O2-


Q32: b) At the bottom of the table.

Q33: c) At the top of the table, in reverse.

Q34: b) F2(g) + 2e- → 2F-(aq)

Q35: a) Na+(aq) + e- → Na(s)



Tutorial - simple ion-electron equations (page 140)

Q36:

Q37:

Q38:

Q39:


Q40: 1

Q41: 2

Q42: 2

Q43: a) Reactant

Q44: b) Reduction



Q45: The oxidation equation for Fe2+ has to be multiplied by 2 and then added to the reductionequation:

Tutorial - complex ion-electron equations (page 144)

Q46: 3

Q47: 6

Q48: 2

Q49: b) Product

Q50: a) Oxidation

Q51: 2ClO3- + 12H+ + 10e- → Cl2 + 6H2O

Q52:

Q53:



Q54: The reduction ion-electron equation can be developed first:

NO3- + 4H+ + 3e- → NO + 2H2O

Combining this with the oxidation:

Summary exercise (page 145)

Q55: b) Fe2O3 + 2Al → 2Fe + Al2O3

Q56: c) Spectator

Q57: c) Fe3+ + Al → Fe + Al3+

Q58: b) loss

Q59: a) Al

Q60: d) Fe3+

Q61: a) Copper

Q62: c) 3 kg


Q63: a) Cl2 + 2e- → 2Cl-

Q64: b) accepts electrons.

Q65: b) 2Br- → Br2 + 2e-

Q66: b) Fe2+ → Fe3+ + e-

Q67: b) Cu2+ → Cu + 2e-

Q68: c) Iodide ions will react with chlorine.

Q69: d) Fe3+

Q70: c) Add two electrons to the left hand side.

Q71: d) Iron(III) ions will react with tin(II) ions.

Q72: CrO72- + 14H+ + 6I- → 2Cr3+ + 7H2O + 3I2



Q73: CrO72- + 14H+ + 6I- → 2Cr3+ + 7H2O + 3I2

Q74: d) Cl2(g)

Q75: c) Ca (s)

Q76: e) Fe2+(aq)

Q77: 3

Q78: 6

Q79: 6

Q80: a) Reactant

Q81: b) a reduction.



Topic 7: End of unit test

End of Unit 1 test (page 154)

Q1: b) decreases.

Q2: b) Sn2+(aq)

Q3: c) large.

Q4: c) Carbon monoxide

Q5: d) the number of occupied energy levels decreases.

Q6: a) CH3CH2COCH3 → CH3CH2CH(OH)CH3

Q7: d) Zn + 2HCl → ZnCl2 + H2

Q8: c) Iron(III) → iron(II)

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)

Q15: Polar covalent

Q16: b) Ionic ⇒ polar covalent ⇒ pure covalent

Q17: Electronegativity

Q18: a) Fe + 2HCl → FeCl2 + H2

Q19: OCl- + 2H+ + 2e- → Cl- + H2O

Q20: 2I- → I2 + 2e-

Q21: a) OCl- + 2H+ + 2I- → Cl- + H2O + I2

Q22: 0·02 mol �-1

Q23: 1·03 g �-1

Q24: a) Cl2(g) + 2Br-(aq) → 2Cl-(aq) + Br2(aq), andb) Br2(l) + 2Fe2+(aq) → 2Br-(aq) + 2Fe3+(aq)

Q25: a) 2IO3-(aq) + 12H+(aq) + 10e- → I2(aq) + 6H2O(l)

Q26: H2O2(aq) + 2H+(aq) + 2I-(aq) → 2H2O(l) + I2(aq)



Q27: 0·0168 g

Q28: As a general rule, award yourself a mark (up to a maximum of three marks in total) for eachpoint you made. (Note: this is a general rule only, remember there are no half marks awarded atHigher.)

• Measure the rate of reaction of the decomposition of hydrogen peroxide with manganeseperoxide and, therefore, the faster the reaction, the higher the concentration of hydrogenperoxide.

• You could also measure the conductivity as electrons are produced. The greater the numberof electrons released, the more hydrogen peroxide has decomposed. Therefore, the higher theconductivity, the more hydrogen peroxide present in the gel.

• You could measure the pH of the products as the lower the pH (higher H+ concentration), themore hydrogen peroxide is present in the gel.

Q29: As a general rule, award yourself a mark (up to a maximum of three marks in total) for eachpoint you made. (Note: this is a general rule only, remember there are no half marks awarded atHigher.)

• If hydrogen peroxide was to leak from a rusty torpedo, it would decompose to form water andoxygen.

• This reaction is exothermic and so releases energy in the form of heat. The oxygen gas whichis released would combust with any fuel it came into contact with and this could cause anexplosion.

• The decomposition of hydrogen peroxide is slow at room temperature and pressure, but thepresence of the iron (in the rust) would act as a catalyst and greatly increase the rate ofdecomposition.

Q30:

Definition Chemical term

Regular arrangement of positively charged ions surrounded bydelocalised electrons.

Metallic bonding

The electrostatic force of attraction of a shared pair of electrons fortwo positive nuclei.

Covalent bonding

The electrostatic attraction between negative and positive ions. Ionic bonding

Weak bonds between molecules. Intermolecular forces

A slight negative or positive charge caused by the unevendistribution of electrons.

Dipole

Bonds holding atoms together within molecules. Intramolecular forces



Q31:

Definition Chemical term

Half the distance between the nuclei of two covalently bonded atoms of anelement.

Covalent radius

The strength of the attraction of an element for the electrons of its bondingelectrons.

Electronegativity

The energy required to remove an electron from a gaseous atom to form anion 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

Q32:

1. Covalent bonding

2. Hydrogen bonding

3. Permanent dipole - permanent dipole interactions

4. London dispersion forces

Q33:

• Going across a period, atomic size decreases.

• Going down a group, atomic size increases.

• Going across a period, electronegativity increases.

• Going down a group, electronegativity decreases.

• Going down Group 7, boiling point increases.


Q34: The number of protons increases OR greater nuclear charge OR greater nuclear attraction.

Q35: Same electronegativity values.

Q36: A sodium atom is bigger than a lithium atom so its electrons are further from the nucleus; theinner electrons shield (screen) the outer electron from the attraction of the nucleus.

Q37: Cl(g) + e- → Cl-(g)

Q38: Covalent bonds are not being broken OR It is intermolecular bonds that are breaking.

(Any alternative wording that recognises that covalent bonds are not broken when covalentsubstances melt or boil will be accepted.)

Q39: Formula refers to the ratio of Mg2+:Cl- ions (in lattice) OR alternative wording, i.e. in the latticethere are twice as many chloride ions as magnesium ions OR Mg2+ ions surrounded by > 2 Cl- ionsOR Cl- surrounded by >1 Mg2+. Any reference to 'chlorine ions' is not acceptable.



Q40: As a general rule, award yourself a half mark, up to a maximum of three marks in total, foreach point you have successfully made (NB this is a general rule only, there are no half marksawarded at Higher):

• Propane molecules are held together by weak intermolecular forces / ethanol molecules areheld 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 atomscreating 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 inthe 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 toovercome the stronger hydrogen bonds than the weak London Dispersion forces.


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SCHOLAR Study Guide Higher Chemistry Unit 1: Chemical ... › 2020 › 04 › scholar-unit-1 … · 5, Unit 1); • all matter is made of atoms - when a substance contains only one