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© 2009 Aristo Educational Press Ltd.14/F Lok's Industrial Building,204 Tsat Tsz Mui Road,North Point,Hong Kong.Tel.: 2811 2908Fax: 2565 6626Website: http://www.aristo.com.hk
All rights reserved. No part of this publication may be reproduced, stored in a retrieval system, ortransmitted in any form or by any means,electronic, mechanical, photo-copying, recording or otherwise, without the prior permission of Aristo Educational Press Ltd.
First published July, 2009
A teacher’s book is available for use by teachers.
Chapter 6 The Periodic Table 26
6.1 Elements with similar chemical properties 26
6.2 The Periodic Table 27
6.3 Patterns in the Periodic Table 30
6.4 Groups — similarities and trends 32
6.5 Predicting chemical properties of an unfamiliar element 35
Key terms 35
Summary 36
Part II Microscopic World I
Chapter 5 Atomic structure 1
5.1 What is an element? 1
5.2 Classification of elements based on physical states 1
5.3 Classification of elements into metals and non-metals 1
5.4 Chemical symbols for elements 5
5.5 Atoms 7
5.6 Structure of atoms 8
5.7 Atomic number and mass number 10
5.8 Isotopes 13
5.9 Relative masses of atoms 15
5.10 Arrangement of electrons 18
5.11 Stability of noble gases related to their electronic arrangements 21
Key terms 23
Summary 24
Chapter 7 Chemical bonding: ionic bonding 37
7.1 Formation of ions from atoms 37
7.2 Colours and migration of ions 38
7.3 Formulae of ions 41
7.4 Elements and ions 45
7.5 Chemical bonds 47
7.6 Ionic bond and ionic substances 47
7.7 Structures of solid ionic compounds 49
7.8 Formulae and names of ionic compounds 50
Key terms 55
Summary 56
Chapter 8 Chemical bonding: covalent bonding 58
8.1 Covalent bonding and covalent substances 58
8.2 Prediction of formulae for covalent compounds 67
8.3 Particles that make up matter — a summary 67
8.4 Relative molecular mass and formula mass 69
Key terms 72
Summary 73
Chapter 9 Structures and properties of substances 75
9.1 Structure of substances 75
9.2 Simple molecular structures 77
9.3 Macromolecules 79
9.4 Giant ionic structures 80
9.5 Giant covalent structures 82
9.6 Giant metallic structures 86
9.7 Comparison of structures and properties of substances 88
9.8 Predicting structure from physical properties 90
9.9 Predicting physical properties from bonding and structure 91
9.10 Applications of substances according to their structures 93
Key terms 94
Summary 94
1
Chapter 5 Atomic structure
In Chapter 1, we have defined that an element is a pure
substance which cannot be broken down into anything simpler
by chemical methods.
5.1 What is an element? 5.1
The simplest way of classifying elements is based on physical
states. At room temperature and pressure, 11 elements are gases, 2
are liquids and the rest are solids. 1 1 2
5.2 Classification of elements based onphysical states
5.2
Metals and non-metals
Another important way of classifying elements is to group
them into metals and non-metals.
• If the element is a gas, it must be a non-metal.
• If the element is a liquid, we have to look at its colour:
♦ Silvery colour indicates the metal mercury (mercury is
the only liquid metal).
♦ Dark red colour indicates the non-metal bromine
(bromine is the only liquid non-metal).
• If the element is a solid, we have to test its electrical
conductivity. (This will be further discussed on p.5.)
♦ Good conductivity indicates a metal in general.
♦ Nil (or poor) conductivity indicates a non-metal
(except graphite).
•
•
♦
( ) (
)
♦
( ) (
)
•
( 5 )
♦
♦ ( )
( )
5.3 Classification of elements into metalsand non-metals
5.3
2
Part II Microscopic World I I
Metals are usually shiny when freshly cut. They are silvery
white in colour, with only a few exceptions (such as copper and
gold).
Solid non-metals usually have a dull appearance. Unlike
metals, they show a variety of colours (e.g. sulphur is yellow,
phosphorus is red or yellow, while carbon is black).
Metals and non-metals differ not only in appearance and
electrical conductivity. They also differ in other ways. See Table
5.1.
(
)
(
)
5.1
Table 5.1 Some typical differences in physical properties of metals and non-metals.
Property
State at room temperature and
pressure
solids (except mercury — a liquid)
usually high
shiny; mostly silvery white in
colour (except copper and gold)
malleable and ductile
hard and strong
usually high
good conductors of heat and
electricity
Melting point and boiling point
Appearance
Hardness and strength
Malleability and ductility
Density
Thermal conductivity and
electrical conductivity
Metals
gases or solids (except bromine —
a liquid)
usually low
usually dull in appearance (when
solid); in various colours
brittle i.e. easily broken when hit
(when solid); not malleable and
not ductile
not uniform in hardness and
strength
low
poor conductors of heat;
non-conductors of electricity
Non-metals
3
Chapter 5 Atomic structure
Note that there are exceptions. Sodium is so soft that it can
be easily cut with a knife; so low-melting that it melts below
100°C; so light that it floats on water. Another example is the
non-metal carbon (in the form of graphite). It is a good
electrical conductor, shiny, and has a very high melting point
(3730°C).
100°C
( )
(3730°C)
Example 5.1Metal or non-metal?
A reddish brown solid element X conducts electricity well.Is X a metal or non-metal? Give reasons.
Solution
X is a metal. All metals conduct electricity while non-metalscannot. (An exception to this rule is the non-metal graphite(a form of carbon), but its colour is black, not reddishbrown.)
5.1
XX
X(
( ) )
5.1
What characteristics do the two elements mercury and bromine
have in common?
Class practice 5.1
✘ All elements can be classified as metals or non-metals.
✔ Many, but not all elements, can be classified as metals or non-metals. A few elements have properties in between those ofmetals and non-metals — they are classified as semi-metals.
Check your concept
✘
✔
4
Part II Microscopic World I I
The in-between elements — the semi-metals
A few elements, called semi-metals (or metalloids), have
properties in between those of metals and non-metals.
Examples of semi-metals include boron and silicon.
Most semi-metals have important uses in industry. An
example is silicon, a semi-conductor. It is widely used in
making transistors and silicon chips.
(
) ( )
Example 5.2Statements about graphite
This question consists of two separate statements. Decidewhether each of the two statements is true or false; if bothare true, then decide whether or not the second statement isa correct explanation of the first statement.
‘Graphite can be considered as a semi-metal.’
‘Graphite conducts electricity as metals do.’
Solution
The first statement is false. Graphite (one form of carbon) isa non-metal. The second statement is true.
5.2
()
1.
(a)
(b)
(c)
2.
(
)
(a)
(b)
(c)
(d) ( )
5.2
1. Would you classify the following elements/compounds as a
metal or non-metal? Why?
(a) Water
(b) Graphite
(c) Mercury
2. Decide which is the odd one in each of the following
groups of elements. Give reason(s) for your choice in each
case.
(a) Iron, copper, mercury, silver
(b) Magnesium, sulphur, lead, tin
(c) Iodine, oxygen, nitrogen, argon
(d) Phosphorus, bromine, helium, carbon (in the form of
graphite)
Class practice 5.2
5
Chapter 5 Atomic structure
Finding whether an element is a metal or non-metal
To find whether an element is a metal or non-metal, a simple
but effective way is to test whether it conducts electricity. We
can use the set-up shown in Figure 5.1. If the element under
test is a metal, the bulb will light up. When non-metals are
tested, the bulb will not light up. All non-metals (except
graphite) are non-conductors of electricity.
5 . 1
(
)
crocodile clip
6 V battery6 V
solid pieceunder test
light bulb
(a)
carbon(graphite) rods
( )
crucible
solid powder (orliquid) under test
( )
(b)
Figure 5.1 Testing electrical conductivity of substances(a) in form of solid piece and (b) in form of solid powder or liquid.
(a)
(b)
5.4 Chemical symbols for elements 5.4
It is useful to give each element a chemical symbol. Chemical
symbols of some common metals, non-metals and semi-metals
are given in Table 5.2.5.2
6
Part II Microscopic World I I
Metal
Table 5.2 Chemical symbols of some common elements (classified into metals, semi-metals and non-metals).
( )
Aluminium ( ) Al
Barium ( ) Ba
Beryllium ( ) Be
Calcium ( ) Ca
Chromium ( ) Cr
Cobalt ( ) Co
Copper ( ) [Cuprum] Cu
Gold ( ) [Aurum] Au
Iron ( ) [Ferrum] Fe
Lead ( ) [Plumbum] Pb
Lithium ( ) Li
Magnesium ( ) Mg
Manganese ( ) Mn
Mercury ( ) [Hydrargyrum] Hg
Nickel ( ) Ni
Platinum ( ) Pt
Potassium ( ) [Kalium] K
Silver ( ) [Argentum] Ag
Sodium ( ) [Natrium] Na
Tin ( ) [Stannum] Sn
Zinc ( ) Zn
Element [Latin name] ChemicalSymbol
Boron ( ) B
Silicon ( ) ( ) Si
Semi-metal
Element ChemicalSymbol
Argon ( ) Ar
Bromine ( ) Br
Carbon ( ) C
Chlorine ( ) Cl
Fluorine ( ) F
Helium ( ) He
Hydrogen ( ) H
Iodine ( ) I
Neon ( ) Ne
Nitrogen ( ) N
Oxygen ( ) O
Phosphorus ( ) P
Sulphur ( ) S
Non-metal
Element ChemicalSymbol
Each chemical symbol shown in the table consists of one or
two letters. The first (or the only) letter is a capital letter; the
second one (if any) is a small letter.
5.2
(a) (i) (ii) (iii)
(b) (i) (ii) (iii)
(c) (i) F (ii) Br (iii) Hg
5.3
Referring to Table 5.2,
(a) Give the chemical symbols for (i) magnesium, (ii) silver and
(iii) sodium.
(b) Give the chemical symbols for the noble gases (i) argon, (ii)
helium and (iii) neon.
(c) Write the names of (i) F, (ii) Br and (iii) Hg.
Class practice 5.3
7
Chapter 5 Atomic structure
5.5 Atoms 5.5
What are atoms?
Everything consists of a basic type of particles called atoms.
An atom is the smallest part of an element which has the
chemical properties of that element.
Size and mass of an atom
If atoms are taken to be spherical, they have diameters of about
10–8 cm. They have masses of around 10–23 g.
Elements and atoms
An element is a substance that is made up of only one kind
of atoms.
Different elements have different properties because they
consist of different kinds of atoms. Until January 2008, 118
kinds of atoms have been discovered or reported,
corresponding to the 118 different elements.
Symbols for atoms
You have learnt chemical symbols of some elements on p.6 —
these are also the atomic symbols for their atoms. Thus the
letter C is the chemical symbol for the element carbon; it is also
the atomic symbol for a carbon atom.
1 0 – 8
cm 10–23 g
2008 1 118
118
6
C
(a)
(b)
(c)
(d) Cu
5.4
(a) What is the total number of atomic symbols at present?
(b) What is the chemical symbol for the element bromine?
(c) What is the atomic symbol for a nitrogen atom?
(d) What does Cu stand for?
Class practice 5.4
8
Part II Microscopic World I I
5.6 Structure of atoms 5.6
Experiments have shown that atoms are in fact made up of
even smaller and simpler particles.
What are atoms made up of?
Atoms are made up of three fundamental sub-atomic particles
— protons, neutrons and electrons.
Atoms are made up of protons, neutrons and electrons. The
protons (positively charged) and neutrons (neutral) are
concentrated in the very tiny nucleus. The electrons
(negatively charged) move around the nucleus.
( ) ( )
(
)
More about protons, neutrons and electrons
Table 5.3 summarizes some data of the three fundamental sub-
atomic particles.5 . 3
Sub-atomicparticle
pProton
Position withinthe atom
Electric charge (relativeto that on a proton)Relative
massMass (in g)(g)
Symbol
1.6725 � 10–24 1 +1 inside nucleus
nNeutron 1.6748 � 10–24 1 0 inside nucleus
e–Electron 9.109 � 10–28
negligible
( )1
1837
–1space outside
nucleus
Table 5.3 Data on the three fundamental sub-atomic particles.
Building up different atoms from protons, neutrons andelectrons
Different atoms have different numbers of protons, neutrons
and electrons.
9
Chapter 5 Atomic structure
The hydrogen atom is the simplest of all atoms. The
commonest type of hydrogen atoms consists of 1 proton and 1
electron (with no neutron). The next simplest one is helium
atom, with 2 protons, 2 neutrons and 2 electrons (Figure 5.2).
1 1
( )
2 2 2 ( 5.2)
neutron
electron
nucleus
proton
}
hydrogen atom helium atom
Table 5.4 gives the number of protons, neutrons and
electrons in the 20 simplest atoms.5.4 20
Atomelectrons neutrons protons
Number of Symbol
Hydrogen ( ) H 1 0 1
Lithium ( ) Li 3 4 3
Boron ( ) B 5 6 5
Nitrogen ( ) N 7 7 7
Fluorine ( ) F 9 10 9
Sodium ( ) Na 11 12 11
Aluminium ( ) Al 13 14 13
Phosphorus ( ) P 15 16 15
Chlorine ( ) Cl 17 18 17
Potassium ( ) K 19 20 19
Helium ( ) He 2 2 2
Beryllium ( ) Be 4 5 4
Carbon ( ) C 6 6 6
Oxygen ( ) O 8 8 8
Neon ( ) Ne 10 10 10
Magnesium ( ) Mg 12 12 12
Silicon ( ) Si 14 14 14
Sulphur ( ) S 16 16 16
Argon ( ) Ar 18 22 18
Calcium ( ) Ca 20 20 20
Figure 5.2 Diagrammaticrepresentations of a hydrogenatom and a helium atom.
Table 5.4 Number of protons,neutrons and electrons in the 20simplest atoms.
20
10
Part II Microscopic World I I
Atoms are electrically neutral
Although an atom contains electrically charged particles, the
atom itself has no overall charge. That is, an atom is electrically
neutral. This is because in an atom, the number of protons is equal
to the number of electrons.
On the other hand, the number of neutrons may not be
equal to the number of protons (look at Table 5.4 again). 5.4
(a)
(b) 91
(c) 8 8
10
5.5
(a) All atoms (except one) are made up of protons, electrons
and neutrons. Which atom does not contain any neutron at
all?
(b) A certain atom contains 91 protons. How many electrons
and neutrons does it have?
(c) A certain particle has 8 protons, 8 neutrons and 10
electrons. Is it an atom? Why?
Class practice 5.5
5 .4
12 17
5.6
Refer to Table 5.4. What would happen if an atom with 12
protons were changed to one with 17 protons?
Class practice 5.6
5.7 Atomic number and mass number 5.7
Atomic number
The atomic number of an atom is the number of protons in
the atom.
For example, a silver atom contains 47 protons. The atomic
number of silver is therefore 47.47
47
11
Chapter 5 Atomic structure
Mass number
The mass number of an atom is the sum of the number of
protons and neutrons in the atom.
For example, a sodium atom (with 11 protons and 12
neutrons) has a mass number of 11 + 12 = 23.( 11
12 ) 11 + 12 = 23
The electrons in an atom have almost no mass. So the mass of anatom is nearly all due to protons and neutrons. For this reason, thenumber of protons plus the number of neutrons in an atom is calledthe mass number.
Learning tip
The atomic number (Z) and mass number (A) of an atom
are usually shown in a full atomic symbol as follows:(Z) (A)
EXAMPLE
mass number
atomic number
42 He
mass number= number of protons + number of neutrons = +
Atomicsymbol
atomic number = number of protons= number of electrons of a neutral atom
= =
A
Z X
Example 5.3Working out the number of protons, electrons andneutrons in an atom
Consider Cl. Work out the number of protons, electrons
and neutrons in the atom.
Solution
Atomic number (Z) = 17, so number of protons = 17 (bydefinition)As an atom is electrically neutral, number of electrons = number of protons = 17Mass number (A) = 35, ∴ number of neutrons = mass number – number of protons
= 35 – 17 = 18
3517
5.3
C l
(Z) = 17 = 17( )
= = 17(A) = 35
∴ = – = 35 – 17 = 18
3517
12
Part II Microscopic World I I
Example 5.4Statements about oxygen atom
This question consists of two separate statements. Decidewhether each of the two statements is true or false; if bothare true, then decide whether or not the second statement isa correct explanation of the first statement.
‘The atomic number of O is 8.’
‘An O atom contains 8 neutrons.’
Solution
Both statements are true, but the second statement does notexplain the first one. A correct explanation would be: ‘An O atom contains 8 protons.’
168
168
168
5.4
O 8
O 8
O 8168
168
168
1.
(a)
47
(b)
47 47
(c)
2. ( = 13)
27
(a)
(b)
(c) (i) (ii)
(iii)
5.7
1. Fill in the following blanks:
A (a) atom has 47 protons. This is what
makes it different from atoms of all other elements. Only
(b) atoms have 47 protons, and any atom
with 47 protons must be a (c) atom.
2. A particular atom of an element (atomic number = 13) has a
mass number of 27.
(a) Name the element.
(b) Write the full atomic symbol for the atom, showing the
mass number and atomic number.
(c) Give the number of (i) protons (ii) electrons (iii)
neutrons in the atom.
Class practice 5.7
13
Chapter 5 Atomic structure
5.8 Isotopes 5.8
What are isotopes?
Isotopes are different atoms of the same element, with the
same number of protons (and electrons) but different numbers of
neutrons.
Let us take hydrogen as an example. Not all of the atoms of
hydrogen are identical. Actually, there are three types of
hydrogen atoms, as shown in Figure 5.3 and Table 5.5. They all
have the same number of protons (same atomic number) and
electrons but different numbers of neutrons. Therefore, hydrogen
has 3 isotopes: H, H and H.31
21
11
( )
5 .3 5.5
( )
H H
H31
21
11
Figure 5.3 The three isotopes of hydrogen. 1H1
2H1
3H1
electron
proton
neutron
Table 5.5 Number of protons, electronsand neutrons in the three isotopes ofhydrogen.
IsotopeNumber of
p e– n
1H1 1 1 0
2H1 1 1 1
3H1 1 1 2
Relative abundance of isotopes
Most elements consist of more than one isotope. In most cases,
one of the isotopes is present in a much higher percentage than
the others in Nature (see Table 5.6). ( 5.6)
14
Part II Microscopic World I I
Element% abundance of
isotopes in NatureMassnumber
AtomicnumberIsotopes
Hydrogen
Carbon
1 H1
2 H1
3 H1
1
1
1
1
2
3
12 C6
13 C6
14 C6
6
6
6
12
13
14
99.984
0.016
very small percentage
98.892
1.108
very small percentage
Oxygen
16 O8
17 O8
18 O8
8
8
8
16
17
18
99.76
0.04
0.20
Sodium 23 Na11 11 23 100
Chlorine 35 Cl17
37 Cl17
17
17
35
37
75.4
24.6
Table 5.6 Isotopes of some elementsin Nature.
5.6
(a)
(b)
5.8
Refer to Table 5.6.
(a) How many natural isotopes does oxygen have?
(b) Which is the most abundant isotope of oxygen?
Class practice 5.8
Comparing properties of different isotopes
Isotopes of the same element have the same number of protons
and electrons in their atoms. They therefore have the same
chemical properties. However, since they have different numbers
of neutrons, they have different masses and hence slightly
different physical properties.
15
Chapter 5 Atomic structure
5.9 Relative masses of atoms 5.9
Relative isotopic mass
The carbon-12 scale
Scientists choose a carbon-12 isotope, which has 6 protons and
6 neutrons, to be the standard atom. Then they fixed it as exactly
12.000 units (atomic mass unit, a.m.u.). Masses of all other
atoms are compared with this reference standard to give their
relative masses.
On the 12C = 12.000 00 scale, the relative masses of a proton
and a neutron are both very close to 1; the relative mass of an
electron is nearly 0. Thus the relative isotopic mass of an
isotope is roughly equal to its mass number.
-12
-12 ( 6
6 )
12.000 (
a.m.u.)
12C = 12.000 00
1
0
‘Relative isotopic mass’ and ‘relative atomic mass’ are both relativevalues; they carry no units.
Learning tip
Relative isotopic mass ≈ mass number ≈
(a) Cl (b) Cl (c) He
(d) U (e) K19238
43517
3717
5.9
What is the relative isotopic mass of
(a) Cl (b) Cl (c) He (d) U (e) K?19238435
173717
Class practice 5.9
16
Part II Microscopic World I I
The relative atomic mass of an element is the weighted
average of the relative isotopic masses of its natural isotopes
on the 12C = 12.000 00 scale.
For example, for an element consisting of three isotopes A,
B and C:
Relative atomic mass = a% � MA + b% � MB + c% � MC
where a%, b%, c% = percentage abundance of isotopes A, B and
C respectively
MA, MB, MC = relative isotopic masses of isotopes A, B
and C respectively
A B C
a% b% c% = A B C
MA MB MC = A B C
(12C = 12.000 00 )
= a% � MA + b% � MB
+ c% � MC
Relative atomic mass
In general, if an element consists of n isotopes, there would be
n different relative isotopic masses, one for each of the isotopes.
However, for the element as a whole, there is only one relative
atomic mass. Hence the relative atomic mass of an element is
determined by: (1) the relative isotopic masses and (2) the relative
abundance of the natural isotopes present in the element. (1)
(2)
17
Chapter 5 Atomic structure
Example 5.5Calculating relative atomic mass and percentageabundance of isotopes
(a) Chlorine consists of two natural isotopes, 35Cl and 37Cl,with percentage abundance of 75.4% and 24.6%respectively. Calculate the relative atomic mass ofchlorine.
(b) Naturally occurring bromine (relative atomic mass =79.9) consists of a mixture of two isotopes: 79Br and 81Br.Calculate the percentage abundance of each of the twoisotopes in natural bromine.
Solution
(a) By approximation,
relative isotopic mass of 35Cl isotope = its mass no. = 35
relative isotopic mass of 37Cl isotope = its mass no. = 37
Relative atomic mass of chlorine
= average mass of 1 chlorine atom on the 12C = 12.000 00 scale
= weighted average of the relative isotopic masses
= � 35 + � 37 = 35.5
(Note: The relative atomic mass of 35.5 is not the relativemass of any one chlorine atom, but the weightedaverage of all the chlorine atoms present.)
(b) Let the percentage abundance of 79Br and 81Br be y%and (100 – y)% respectively.
Relative atomic mass of bromine = weighted average ofthe relative isotopicmasses
79.9 =
7990 = 79y + 8100 – 81y
∴ y = 55
Thus the percentage abundance of 79Br is 55% and thatof 81Br is 45%.
79y + 81(100 – y)100
24.6100
75.4100
5.5
(a) 3 5C l3 7 C l 7 5 . 4 %2 4 . 6 %
(b) 7 9B r81Br
(= 79.9 )
(a)35Cl =
35Cl = 3537Cl =
37Cl = 37
= 1 12C = 12.000 00
=
= � 35 + � 37 = 35.5
(35.5
35.5 )
(b) 79Br 81Bry% (100 – y) %
=
79.9 =
7990 = 79y + 8100 – 81y
∴ y = 55
79Br 81Br55% 45%
79y + 81(100 – y)100
24.6
100
75.4
100
Shell number,
n
Maximum number of electrons (= 2n2)
1234...
28
1832...
18
Part II Microscopic World I I
1. Na
2. 9 0 % N e
10% Ne2210
2010
2311
5.10
1. There is only one kind of sodium atoms in nature, i.e. Na.
What is the relative atomic mass of sodium?
2. Neon in air contains 90% of Ne and 10% of Ne.
Calculate the relative atomic mass of neon.
2210
2010
2311
Class practice 5.10
The accurate relative atomic masses of elements are very
seldom whole numbers (why?). ( )
✘ The relative atomic mass of chlorine is 35.5 g.
✔ The relative atomic mass is a relative value. It carries no unit.The relative atomic mass of chlorine should be 35.5.
Check your concept
✘ 35.5 g
✔
35.5
5.10 Arrangement of electrons 5.10
Electronic arrangement
Scientists think that electrons in an atom exist in a number of
regions (called electron shells) surrounding the central
nucleus.
Each electron shell is given a number 1, 2, 3, 4 and so on,
starting from the one closest to the nucleus (i.e. the innermost
shell). Each shell can hold up to a certain maximum number of
electrons (Table 5.7).
The arrangement of electrons in a sodium atom can be
shown by Figure 5.4.
(
)
1 2 3 4
( )
( 5.7)
5.4
Table 5.7 Maximum number of electronsthe first four shells can hold.
19
Chapter 5 Atomic structure
Figure 5.4 Arrangement of electrons ina sodium atom.
nucleus
electron
3rd shell (outermost shell in sodium atom)
( )
2nd shell
1st shell (innermost shell)( )
Rules for finding electronic arrangement
To find the electronic arrangement of an atom, we use the
following rules:
1. The atomic number of the element is first found. This is
equal to the number of protons, and hence the number of
electrons present in an atom of the element.
2. Electrons go into the shells one by one, starting from the
innermost shell. When a certain shell is ‘full’ (refer to Table
5.7 again), any remaining electrons would go into the next
outer shell and so on, until all are placed.
Ways of representing electronic arrangement
Electronic arrangement by numbering
Electronic arrangement may be shown by numbering. The
number of electrons in each shell is listed, starting from the
first shell (innermost shell); the numbers are separated by
commas. For example, the electronic arrangement of a sodium
atom is 2, 8, 1 (Figure 5.5).
Electrons in an atom are arranged in shells. The distribution
of electrons in the various shells is called electronic
arrangement (or electronic configuration).
1.
2.
( 5 .7)
( )
2, 8, 1
( 5.5)
( )
20
Part II Microscopic World I I
Figure 5.5 Showing the electronicarrangement of a sodium atom bynumbering.
Electronic arrangement of sodium atom:
2, 8, 1Number of 1st 2nd 3rdelectrons in: shell shell shell
Electronic arrangement by diagram
Besides numbering, electronic arrangement can also be
represented by an electron diagram. In such diagrams, the
nucleus is often represented by the symbol of the atom.
Electron shells are shown by circles around the nucleus.
Electrons are shown by dots or crosses. Figure 5.6 is the
electron diagram of a sodium atom.
5.6
Na
Figure 5.6 The electron diagram of a sodium atom.
(a) (b)
(c) (d)
5.11
Draw electron diagrams for the following atoms:
(a) Helium (b) Oxygen (c) Silicon (d) Calcium
Class practice 5.11
Electronic arrangements of the first 20 elements
Following the above rules, we can find the electronic
arrangements of the elements with atomic numbers 1 – 20
(Table 5.8).
1 –
20 ( 5.8)
21
Chapter 5 Atomic structure
1
2
Electronicarrangement
2, 1
2, 2
2, 3
2, 4
2, 5
2, 6
2, 7
2, 8
2, 8, 1
2, 8, 2
2, 8, 3
2, 8, 4
2, 8, 5
2, 8, 6
2, 8, 7
2, 8, 8
Element Symbol Atomicnumber
No. of electrons in electron shells
1st 2nd 3rd 4th
Hydrogen
Helium
H
He
1
2
1
2
1
2
Lithium
Beryllium
Boron
Carbon
Nitrogen
Oxygen
Fluorine
Neon
Li
Be
B
C
N
O
F
Ne
3
4
5
6
7
8
9
10
3
4
5
6
7
8
9
10
2
2
2
2
2
2
2
2
1
2
3
4
5
6
7
8
Na
Mg
Al
Si
P
S
Cl
Ar
11
12
13
14
15
16
17
18
11
12
13
14
15
16
17
18
2
2
2
2
2
2
2
2
8
8
8
8
8
8
8
8
1
2
3
4
5
6
7
8
Potassium
Calcium
K
Ca
19
20
19
20
2
2
8
8
8
8
1
2
2, 8, 8, 1
2, 8, 8, 2
Table 5.8 The electronic arrangements (by numbering) of the elements with atomic numbers 1 – 20.1 – 20 ( )
Sodium
Magnesium
Aluminium
Silicon ( )
Phosphorus
Sulphur
Chlorine
Argon
Number ofelectrons
(a) (
5.8 )
(b) (i) (ii)
5.12
(a) What is the atomic number of chlorine? (See Table 5.8)
(b) Show the electronic arrangement of a chlorine atom by
(i) numbering (ii) an electron diagram.
Class practice 5.12
5.11 Stability of noble gases related to theirelectronic arrangements
5.11
The term ‘noble gases’ is a collective name for Group 0
elements, which are very unreactive.0
22
Part II Microscopic World I I
The exceptional stability of noble gases is related to their
electronic arrangements:
Helium (He) 2
Neon (Ne) 2, 8
Argon (Ar) 2,8, 8
Krypton (Kr) 2,8,18, 8
Xenon (Xe) 2,8,18,18, 8
Radon (Rn) 2,8,18,32,18, 8
All noble gases (except helium) have 8 outermost shell
electrons in their atoms. Helium atom has 2 electrons in the
only one occupied shell. This suggests that a particle has great
stability if it has
• an octet of electrons (i.e. 8 electrons in the outermost shell)
or
• a duplet of electrons (i.e. 2 electrons in the only one
occupied shell).
Atoms of elements other than noble gases are usually not
stable. They will become stable if they attain an octet or a
duplet.
(He) 2
(Ne) 2, 8
(Ar) 2, 8, 8
(Kr) 2, 8, 18, 8
(Xe) 2, 8, 18, 18, 8
(Rn) 2, 8, 18, 32, 18, 8
( ) 8
2
• ( 8
)
• (
2
)
23
Chapter 5 Atomic structure
KK ee yy tt ee rr mm ss
1. atom 7
Page
3. chemical symbol 5
4. duplet 22
7. electronic arrangement 19
8. electronic configuration 19
9. element 1
10. isotope 13
11. mass number 11
12. metal 1
13. non-metal 1
14. octet 22
15. relative abundance 13
16. relative atomic mass 16
17. relative isotopic mass 15
18. semi-metal/metalloid / 4
5. electron diagram 20
2. atomic number 10
6. electron shell ( ) 18
24
Part II Microscopic World I I
SS uu mm mm aa rr yy5.1 What is an element?
1. An is a pure substance which cannot be broken down into anything simpler bychemical methods.
5.2 Classification of elements based on physical states
2. Elements can be classified based on states, that is, whether the elements aresolids, liquids or gases. (a silvery metal) and (a dark red non-metal) are the only two liquid elements.
5.3 Classification of elements into metals and non-metals
3. Elements can be classified into , and .
4. All metals conduct . All non-metals (except carbon in the form of graphite) donot conduct . To tell whether an element is a metal or non-metal, a simple buteffective way is to test whether it conducts .
(Refer to Table 5.1 on p.2 for some typical differences in physical properties between metals andnon-metals.)
5.4 Chemical symbols for elements
5. Chemists use chemical to represent elements. Chemical symbols of mostelements come from their English names.
(Refer to Table 5.2 on p.6 for chemical symbols of some common metals, non-metals and semi-metals.)
5.5 Atoms
6. An is the smallest part of an element which has the chemical properties of thatelement.
7. An is a substance that is made up of only one kind of atoms. Different elementshave different properties because they consist of different kinds of atoms.
5.6 Structure of atoms
8. (a) An atom consists of three types of sub-atomic particles — ,and .
(b)
–1
Relative chargeSub-atomic particle Relative mass
Proton (p) 1 +1
Neutron (n) 1 0
Electron (e–) negligible ( )11837
25
Chapter 5 Atomic structure
(c) An atom has an extremely small centre called . The protons and neutronsare in the nucleus.
(d) Electrons move around the nucleus in .
(e) An atom is electrically .
5.7 Atomic number and mass number
9. of an atom = number of protons in the atom
of an element = number of protons in an atom of the element
10. of an atom = number of protons + number of neutrons in theatom
11. Full atomic symbol
5.8 Isotopes
12. are different atoms of the same element, with the same number of protons (andelectrons) but different numbers of neutrons. Different isotopes of the same element have thesame chemical properties but slightly different physical properties.
5.9 Relative masses of atoms
13. ≈ mass number
14. of an element = weighted average of the relative isotopic massesof its natural isotopes on the 12C = 12.000 00 scale.
5.10 Arrangement of electrons
15. The of an atom is the distribution of electrons in the variousshells of the atom. (Refer to Table 5.8 on p.21.)
5.11 Stability of noble gases related to their electronic arrangements
16. Noble gases have great stability because their atoms have either an of electrons(8 electrons in the outermost shell), or a of electrons (2 electrons in the only oneoccupied shell) as in helium.
EXAMPLE
mass number
atomic number
42 He
mass number= number of protons + number of neutrons
Atomicsymbol
atomic number = number of protons= number of electrons of a neutral atom
A
Z X
26
Part II Microscopic World I I
Grouping elements
There are 92 naturally occurring elements. If we can find a way
to group these elements, we can study them more easily and
systematically.
A. Action of water on potassium, sodium and iron
Both potassium and sodium react vigorously with water.
Iron has no immediate reaction with water. Thus
potassium and sodium behave similarly.
B. Action of dilute hydrochloric acid on calcium,magnesium and copper
Both calcium and magnesium react with dilute
hydrochloric acid to give a colourless gas. Copper has no
reaction with the acid. Thus calcium and magnesium
behave similarly.
C. Action of sodium sulphite solution on aqueouschlorine solution, aqueous bromine solution,aqueous iodine solution and sulphur
On adding sodium sulphite solution, aqueous solutions of
chlorine, bromine and iodine all turn colourless; sulphur
has no reaction. Thus chlorine, bromine and iodine behave
similarly.
92
A.
B.
C.
6.1 Elements with similar chemicalproperties
6.1
27
Chapter 6 The Periodic Table
6.2 The Periodic Table 6.2
Development of the Periodic Table
In 1869, the Russian chemist Mendeleev arranged the 63
elements known at that time in a table form. He put elements
with similar chemical properties in the same vertical column of
the table. He called his table the Periodic Table of Elements.
This table has been much modified over the years, to become
the modern Periodic Table.
The modern Periodic Table
In the modern Periodic Table (Table 6.1), elements are arranged
in ascending order of atomic number. For example, hydrogen
(atomic number 1) comes first. Helium (atomic number 2) comes
second and so on.
1869
63
( 6.1)
Table 6.1 Part of the modern Periodic Table. (A complete Periodic Table is shown on the inside front cover.)
( )
GROUPS
Transition elements
main groups
atomic number relative atomic mass
electronic arrangement
Keys:
metal semi-metal non-metal gas liquid solid
Hal
ogen
s
Nob
le g
ases
Alka
li m
etal
s
Alka
line
eart
h m
etal
s
PER
IOD
S
28
Part II Microscopic World I I
Period number = number of occupied electron shells
Periods
A horizontal row of elements is called a period. Each period
has a number: from 1 to 7. Period 1 contains only two elements.
Period 2 and Period 3 each contains eight elements. Other
periods are longer.
We should note that Period 1 elements have one occupied
electron shell, Period 2 elements have two occupied electron
shells, and so on.
Group number = number of electrons in the outermost shell
Groups
A vertical column of elements is called a group. There are
altogether eight main groups. Each group has a number (I, II,
III, IV, V, VI, VII or 0).
We should note that Group I elements have one outermost
shell electron, Group VII elements have seven outermost shell
electrons, and so on.
There are exceptions to this rule: (1) Hydrogen does not belong toany group. (2) For Group 0 elements, helium has two electrons inthe outermost shell, while all the others have eight.
Learning tip
(1) (2) 0 2
8
=
I II III IV V VI VII 0
I
V I I
=
The elements are arranged in periods and groups of the
Periodic Table.
29
Chapter 6 The Periodic Table
Figure 6.1 illustrates the above two rules for period
number and group number.
Some of the groups have special names:
Group I : Alkali metals
Group II : Alkaline earth metals
Group VII : Halogens
Group 0 : Noble gases
6.1
I
II
VII
0
no. of occupiedelectron shells = 3 =period no.
= 3=
no. of electrons in theoutermost shell = 7 =group no. (VII)
= 7 = (VII)
Electronic arrangement of a chlorine atom:
Figure 6.1 The relation among electronicarrangement, period number and groupnumber.
2, 8, 7
Example 6.1Identifying an unknown element based on its atomicnumber
Element X has an atomic number of 15.
(a) Deduce the electronic arrangement of an atom of X.
(b) In which (i) group (ii) period of the Periodic Tableshould X be placed?
(c) Is X a metal or a non-metal?
Solution
(a) 2,8,5 (b) (i) Group V (ii) Period 3(c) Non-metal
6.1
X 15
(a) X
(b) X (i) (ii)
(c) X
(a) 2,8,5(b) (i) V (ii) (c)
30
Part II Microscopic World I I
W 2, 8, 18, 32, 18,
8, 2
(a) W
W
(b) W
(c) W
6.1
Element W has the electronic arrangement of 2, 8, 18, 32, 18, 8,
2.
(a) To which period and group of the Periodic Table does W
belong? What is the special name of the group?
(b) By referring to the Periodic Table, name element W.
(c) Predict whether W can conduct electricity. Give your
reason.
Class practice 6.1
The elements in between Group II and Group III are called
the transition elements (or transition metals). Many common
metals such as iron (Fe) and copper (Cu) are transition
elements.
II III
( )
(Fe) (Cu)
6.3 Patterns in the Periodic Table 6.3
Changing from metals to non-metals across a period
Across a period, the elements change from metals through
semi-metals to non-metals. For example, across Period 2, there
is a gradual change from a reactive metal (lithium), through a
less reactive metal (beryllium), a semi-metal (boron), less
reactive non-metals (carbon, nitrogen), to reactive non-metals
(oxygen, fluorine), and finally to a noble gas (neon). See Figure
6.2.
( )
( ) ( )
( ) ( )
( ) 6.2
Group
Period I
Li
Na
Be
Mg
B
Al
C
Si
N
P
O
S
F
Cl
Ne
Ar
II III IV V VI VII 0
2
3
more metallic
more non-metallic
reactive metals
less reactive metals
less reactive non-metals
reactive non-metals
noblegases
semi-metals
Figure 6.2 Elements change from metalsto non-metals across Period 2 and Period3 of the Periodic Table.
31
Chapter 6 The Periodic Table
X II
Y 0
Z IV
X Y Z
6.2
Element X belongs to Group II of the Periodic Table.
Element Y is a Group 0 element.
Element Z is a Group IV element.
Try to classify X, Y and Z as a metal or a non-metal.
Class practice 6.2
Electronic arrangement and chemical properties
Electronic arrangements of some elements in the Periodic Table
are given below:
Group I I Group VII VII Group 0 0
Period 2 Period 3 Period 4 Period 5 Period 6
Li 2, 1Na 2,8, 1K 2,8,8, 1Rb 2,8,18,8, 1Cs 2,8,18,18,8, 1
F 2, 7Cl 2,8, 7Br 2,8,18, 7I 2,8,18,18, 7At 2,8,18,32,18, 7
Ne 2, 8Ar 2,8, 8Kr 2,8,18, 8Xe 2,8,18,18, 8Rn 2,8,18,32,18, 8
Let us take Group I as an example. All Group I elements
have one outermost shell electron. They have similar chemical
properties. This suggests the following relationship:
Chemical properties of an element depend mainly on the
number of outermost shell electrons.
I I
✘ Elements in the same group have the same chemicalproperties.
✔ Elements in the same group have similar chemical properties.
Check your concept
✘
✔
32
Part II Microscopic World I I
Example 6.2Deciding which elements show similar chemicalproperties
Which of the following pairs of atoms would have similarchemical properties? Explain your answer.
A. X and Y B. X and Y
C. X and Y D. X and Y
Solution
The subscripts stand for atomic numbers. Electronicarrangements of the atoms:
A. X (2, 4) and Y (2, 8, 5)
B. X (2, 2) and Y (2, 8, 8, 2)
C. X (2, 7) and Y (2, 8, 6)
D. X (2, 5) and Y (2, 8, 7)
In B, X and Y have the same number of outermost shell
electrons, so they should have similar chemical properties.204
177
169
204
156
177169
204156
6.2
A. X Y B. X Y
C. X Y D. X Y
A. X (2, 4) Y (2, 8, 5)
B. X (2, 2) Y (2, 8, 8, 2)
C. X (2, 7) Y (2, 8, 6)
D. X (2, 5) Y (2, 8, 7)
X Y204
177
169
204
156
177169
204156
P 20
(a) P
(b) P
(c) P
(i) Q (ii) R128
6.3
The atomic number of an element P is 20.
(a) What is the electronic arrangement of a P atom?
(b) Would P conduct electricity? Why?
(c) Which of the following atoms would have chemical
properties similar to P?
(i) Q (ii) R128
Class practice 6.3
6.4 Groups — similarities and trends 6.4 —
Elements within the same group of the Periodic Table have
similar chemical properties. Yet there is also a gradual change
in chemical properties down a group. Let us take Group I,
Group VII and Group 0 as examples. I VII 0
33
Chapter 6 The Periodic Table
Group I: The alkali metals
Figure 6.3 shows the elements in Group I.
I
6.3 I
lithium
}sodium
potassium
rubidium
caesium
francium
silvery solids
Figure 6.3 Group I elements (thealkali metals).
I ( )
Li
Na
K
Rb
Cs
Fr
Similarities of Group I elements
1. All are soft metals.
2. All are silvery solids (when freshly cut).
3. All are reactive.
4. All have similar chemical properties.
5. All react with water, giving off hydrogen to form an alkaline
solution. That is why we call them alkali metals.
Difference in reactivity of Group I elements
Although all alkali metals are reactive, they differ in
reactivities.
I
1.
2.
3.
4.
5.
I
Reactivity of Group I elements increases down the group.
In fact, this rule also applies to Group II elements (the
alkaline earth metals).
Group VII: The halogens
Figure 6.4 shows the elements in Group VII.
I
I I
( )
VII
6.4 VII
34
Part II Microscopic World I I
fluorine (pale yellow gas)
(greenish yellow gas)
(dark red liquid)
(black solid)
(black solid)
chlorine
bromine
iodine
astatine
Figure 6.4 Group VII elements (the halogens).
VII ( )
F
Cl
Br
I
At
Similarities of Group VII elements
1. All are poisonous non-metals.
2. All are reactive.
3. All have similar chemical properties. For example, their
aqueous solutions are turned colourless by sodium sulphite
solution (p.26).
Difference in reactivity of Group VII elements
VII
1.
2.
3.
( 26 )
VII
Reactivity of Group VII elements decreases down the group.
Group 0: The noble gases
Figure 6.5 shows the elements in Group 0.
VII
0 ( )
6.5 0
helium
neon
argon
krypton
xenon Figure 6.5 Elements in Group 0 (the noble gases).
0 ( )
He
Ne
Ar
Kr
Xe
radon Rn
} colourless gases
0
1.
2.
Similarities of Group 0 elements
1. All are colourless gases.
2. All are very stable. They have little or no reaction with
other elements.
35
Chapter 6 The Periodic Table
6.5 Predicting chemical properties of anunfamiliar element
6.5
We can predict the chemical properties of an element from its
position in the Periodic Table.
1.
2. ( 2, 8,
18, 32, 18, 7) ( 2, 8,
18, 8, 2)
A.
B.
C.
D.
6.4
1. Can the chemical properties of an unfamiliar element be
deduced from its electronic arrangement? Why?
2. Which of the following correctly describes the elements
astatine (electronic arrangement 2,8,18,32,18,7) and
strontium (2,8,18,8,2) respectively?
A. A metal more reactive than magnesium
B. A metal less reactive than magnesium
C. A non-metal more reactive than chlorine
D. A non-metal less reactive than chlorine
Class practice 6.4
KK ee yy tt ee rr mm ssPage
1. group 28
2. main group 28
4. Periodic Table of Elements 27
5. reactivity 33
6. transition element 30
7. transition metal 30
3. period 28
36
Part II Microscopic World I I
SS uu mm mm aa rr yy6.1 Elements with similar chemical properties
1. Some elements show chemical properties.
6.2 The Periodic Table
2. In the modern Periodic Table, all elements are arranged in increasing order of.
3. (a) The Periodic Table consists of periods and groups.
(b) A horizontal row of elements is called a .
(c) A vertical column of elements is called a .
(d) For elements in the main groups:
(1) Period number of an element
= number of electron shells in an atom of the element
(2) Group number of an element
= number of electrons in an atom of the element
6.3 Patterns in the Periodic Table
4. Across a period from left to right, there is a change from metals, to and finallyto .
5. Elements within the same group of the Periodic Table have the number ofoutermost shell electrons in their atoms, therefore they have chemicalproperties. However, there is a gradual change in reactivity down a group.
6.4 Groups — similarities and trends
6. Group I elements are called the .
Group II elements are called the .
Reactivity of Group I and II elements down the group.
7. Group VII elements are called the .
Reactivity of Group VII elements down the group.
8. Group 0 elements are called the . They are all very unreactive.
6.5 Predicting chemical properties of an unfamiliar element
9. Chemical properties of an unfamiliar element can be predicted from its in thePeriodic Table.
37
Chapter 7 Chemical bonding: ionic bonding
Stability of noble gases
All noble gases (except helium) have 8 outermost shell
electrons in their atoms. Helium atom has 2 electrons in the
only one occupied shell. This suggests that a particle has great
stability if it attains
• an octet of electrons (i.e. 8 electrons in the outermost shell)
or
• a duplet of electrons (i.e. 2 electrons in the only one
occupied shell).
What is an ion?
( )
• (
)
• (
)
?
7.1 Formation of ions from atoms 7.1
A simple ion is derived from a single atom. A polyatomic
ion is derived from a group of atoms.
Examples of simple ions are sodium ion, lead(II) ion,
copper(II) ion, chloride ion and bromide ion. Examples of
polyatomic ions are ammonium ion, hydroxide ion, sulphate
ion, nitrate ion and permanganate ion.
Cations and anions
There are two kinds of ions: positively charged ions and
negatively charged ions. Positive ions are called cations — they
are attracted towards the cathode (negatively charged electrode
in electrolysis). Negative ions are called anions — they are
attracted towards the anode (positively charged electrode in
electrolysis). See Figure 7.1.
An ion is an atom or a group of atoms having an overall
electric charge.
(II)
(II)
— (
)
— (
) 7.1
38
Part II Microscopic World I I
Figure 7.1 Movement of cations andanions in electrolysis.
--
++ cation anion
anode
electron flow
electrolyte
cathode
(a) /
(b) /
7.1
Referring to the above discussion on cations and anions, delete
(cross out) the unsuitable words in the following statements:
(a) Cations are ions that usually come from metals/non-metals.
(b) Anions are ions that usually come from metals/non-metals.
Class practice 7.1
7.2 Colours and migration of ions 7.2
Colour of ions
Many ions are colourless. However, some ions are coloured.
We should notice that transition metals usually form
coloured ions; most of these are cations (e.g. copper(II) ion), but
a few are polyatomic anions (e.g. permanganate ion). On the
other hand, elements in the main groups in the Periodic Table
form colourless ions (not listed in Table 7.1).
Name Colour
(a) Copper(II) ion
(b) Iron(II) ion
(c) Iron(III) ion
(d) Cobalt(II) ion
(e) Nickel(II) ion
(f) Chromium(III) ion
(g) Chromate ion
(h) Dichromate ion
(i) Manganese(II) ion
(j) Permanganate ion
blue or green
pale green
yellow or brown
pink
green
green
yellow
orange
very pale pink
purple
(II)
(II)
(III)
(II)
(II)
(III)
(II) Table 7.1 The colours of someions in aqueous solution.
(
(II) )
( )
(
7.1 )
39
Chapter 7 Chemical bonding: ionic bonding
( )
(a)
(b)
(c)
(d) (II)
7.2
Predict the colour (if any) of each of the following solutions:
(a) Magnesium nitrate solution
(b) Sodium permanganate solution
(c) Ammonium chromate solution
(d) Iron(II) sulphate solution
Class practice 7.2
Gemstones and ions
Colours of gemstones
Gemstones are very rare minerals, usually coloured.
Coloured ions in gemstones
Colours of gemstones are due to traces of coloured ions. Someexamples are given in Table 7.2. 7.2
Gemstone
Amethyst
Emerald
Jade
Peridot
Topaz
Turquoise
Ion responsible for colour
manganese(III) ion (III)
chromium(III) ion (III)
chromium(III) ion (III)
iron(II) ion (II)
iron(III) ion (III)
copper(II) ion (II)
purple
green
green
light green
yellow
bluish green
Colour
Table 7.2 Coloured ions in some gemstones.
Migration of ions
We can observe the migration (movement) of coloured ions
during electrolysis, using the set-up as shown in Figure 7.2.7.2
40
Part II Microscopic World I I
dilute hydrochloric acid
this region slowly becomes orange due to the migration ofnegative dichromate ions towards the positive anode
carbon anode carbon cathode
dilute hydrochloric acid
this region slowly becomes blue due to themigration of positive copper(II) ions towardsthe negative cathode
(II)
a gel containing copper(II) ions and dichromate ions(II)
20 V d. c. supply 20 V
Figure 7.2 To show the migration of coloured ions during electrolysis (using a U-tube).
( U )
A simpler way of investigating the migration of coloured
ions under the influence of an electric field is shown in Figure
7.3.
7 . 3
small potassiumpermanganatecrystal
filter papermoistened withsodium sulphatesolution
purple spot microscope slide
anode cathode
20 V d.c. supply20 V
Figure 7.3 To show the migration of purple permanganate ions under the influence of an electric field (using a strip of filter paper on amicroscope slide).
( )
small potassiumpermanganate crystal
filter paper moistened withsodium sulphate solution
purple spot
microscope slideanode cathode
7.3
(a)
(b)
(c) (III)
7.3
Refer to Figure 7.3 again.
(a) Towards which electrode are potassium ions migrating?
Why?
(b) Can we see the movement of potassium ions? Why?
(c) If a chromium(III) sulphate crystal was used instead of a
potassium permanganate crystal, what would be observed?
Why?
Class practice 7.3
41
Chapter 7 Chemical bonding: ionic bonding
7.3 Formulae of ions 7.3
Formation of ions
An atom is overall electrically neutral, because it has the same
number of protons and electrons. But if the number of electrons
in an atom is increased or decreased, an ion is formed.
Example 7.1Understanding how an ion is formed
Explain, in terms of electronic arrangement and number ofprotons and electrons, the formation of
(a) a lithium ion (b) an oxide ion.
Solution
(a) Consider a lithium atom, Li.
Electronic arrangement: 2,1Number of protons = 3; number of electrons = 3Charge of the atom = (+1) � 3 + (–1) � 3 = 0
(i.e. the atom carries no charge)
To get the electronic arrangement of the nearest noblegas (helium) — 2 (which is a duplet), one electron hasto be removed. An ion is formed.
Number of electrons = 3 – 1 = 2Charge of the ion = (+1) � 3 + (–1) � 2 = +1 (writtenas 1+ or +)
The resulting positive ion is called lithium ion,represented by Li+.
(Note that ‘1’ is usually dropped out in writing thecharge on an ion. Thus we write Li+ instead of Li1+.)
(b) Consider an oxygen atom, O.
Electronic arrangement: 2,6Number of protons = 8; number of electrons = 8Charge of the atom = (+1) � 8 + (–1) � 8 = 0
(i.e. the atom carries no charge)
To get the electronic arrangement of the nearest noblegas neon — 2,8 (which is an octet), two electrons haveto be gained. An ion is formed.
Number of electrons = 8 + 2 = 10Charge of the ion = (+1) � 8 + (–1) � 10 = –2 (writtenas 2–)
The resulting negative ion is called oxide ion (notoxygen ion), represented by O2–.
7.1
(a) (b)
(a) Li
2,1
= 3 = 3
= (+1) � 3 + (–1) � 3= 0
( )
( )— 2 ( )
= 3 – 1 = 2
= (+1) � 3 + (–1) � 2= +1 ( 1+ +)
Li+
(1 Li+
Li1+ )
(b) O
2,6
= 8 = 8
= (+1) � 8 + (–1) � 8= 0
( )
( )— 2,8 ( )
= 8 + 2 = 10
= (+1) � 8 + (–1) �10 = –2 ( 2–)
O2–
42
Part II Microscopic World I I
Polyatomic ions are formed from a group of atoms.
However, their formation is not discussed here.
1.
(a)
(b)
2. 1
7.4
1. Write down the electronic arrangements of
(a) aluminium atom and aluminium ion
(b) chlorine atom and chloride ion
2. Put down the charge of each ion in Question 1.
Class practice 7.4
(a) H2 (b) H+
(c) H
(d) NH4+
(e) CCl4 (f) NH3
(g) H–
(h) NH2–
(i) OH–
(j) Mn2+
7.5
State which of the following formulae stand for simple ions and
polyatomic ions respectively.
(a) H2 (b) H+
(c) H (d) NH4+
(e) CCl4 (f) NH3 (g) H–
(h) NH2–
(i) OH–
(j) Mn2+
Class practice 7.5
What is a formula?
We can refer to an element, a compound or an ion by its name.
Alternatively, we can refer to it by its formula (plural:
formulae).
Names and formulae of common ions
Table 7.3 gives the names of some common ions with their
formulae.7.3
43
Chapter 7 Chemical bonding: ionic bonding
Anions
Charge Formula Name
H– hydride ion
Cl– chloride ion
Br– bromide ion
I– iodide ion
OH– hydroxide ion
NO3– nitrate ion
1– NO2– nitrite ion
HCO3– hydrogencarbonate ion
HSO4– hydrogensulphate ion
CN– cyanide ion
MnO4– permanganate ion
ClO3– chlorate ion
ClO– hypochlorite ion
O2– oxide ion
S2– sulphide ion
SO42– sulphate ion
SO32– sulphite ion
SiO32– silicate ion
2– CO32– carbonate ion
CrO42– chromate ion
Cr2O72– dichromate ion
N3– nitride ion
3– P3– phosphide ion
PO43– phosphate ion
Table 7.3 The names and formulae of some common ions.
Cations
Charge Formula Name
Na+ sodium ion
K+ potassium ion
Cu+ copper(I) ion(I)
Ag+ silver ion
Hg+ mercury(I) ion(I)
H+ hydrogen ion
1+ NH4+ ammonium ion
Mg2+ magnesium ion
Ca2+ calcium ion
Ba2+ barium ion
Pb2+ lead(II) ion(II)
Fe2+ iron(II) ion(II)
2+ Co2+ cobalt(II) ion(II)
Ni2+ nickel(II) ion(II)
Mn2+ manganese(II) ion(II)
Cu2+ copper(II) ion(II)
Zn2+ zinc ion
Hg2+ mercury(II) ion(II)
Al3+ aluminium ion
3+ Fe3+ iron(III) ion(III)
Cr3+ chromium(III) ion(III)
Cations
Charge Formula Name
44
Part II Microscopic World I I
Refer to Table 7.3. You should pay special attention to the
following points:
1. All simple metal ions (e.g. Na+, Mg2+) are cations.
2. All simple non-metal ions (except H+) and most polyatomic
ions (e.g. OH–, HCO3–) are anions (except NH4
+).
3. There is only one common polyatomic cation — NH4+.
4. Polyatomic ions usually consist of non-metals only (e.g.
NO3–, CO3
2–, SO42–), but some consist of a metal and a non-
metal (e.g. MnO4–, CrO4
2–, Cr2O72–).
5. When a metal forms only one cation, the ion has the same
name as the metal, e.g. sodium metal (Na) forms sodium ion
(Na+).
6. Transition metals can form more than one simple cation with
different charges. To name each ion, a Roman numeral
indicating the charge is written in brackets after the name
of the metal. For example, iron metal (Fe) can form iron(II)
ion Fe2+ and iron(III) ion Fe3+.
7. Simple anions have names ending in -ide, e.g. an oxygen
atom (O) forms an oxide ion (O2–); a sulphur atom (S)
forms a sulphide ion (S2–).
8. The polyatomic anion with more oxygen is named as -ate,
and that with less oxygen as -ite, e.g. SO42– sulphate ion,
SO32– sulphite ion; NO3
– nitrate ion, NO2– nitrite ion.
9. Ions with 4+ or 4– charges are uncommon. They are not
listed in the table.
7.3
1. ( Na+ Mg2+)
2. (H+ )
( OH– HCO3–)
(NH4+ )
3.
— NH4+
4.
( NO3– CO3
2– SO42–)
(
MnO4– CrO4
2– Cr2O72–)
5.
(Na) (Na+)
6.
(II) (III)
7.
CO32–
SO42–
SO32–
N O 3–
NO2–
8. 4+ 4–
45
Chapter 7 Chemical bonding: ionic bonding
7.4 Elements and ions 7.4
Which elements form ions?
A metal atom has few outermost shell electrons (usually 1 to 3).
To get a noble gas electronic arrangement, the easiest way is to
lose these electrons, forming a cation (positively charged). For
example, a Mg atom (2,8,2) forms a Mg2+ ion (2,8). See Figure
7.4a.
A non-metal atom has more outermost shell electrons. To
get a noble gas electronic arrangement, it is easier for the atom
to gain rather than to lose electrons. It thus gains electrons,
forming an anion (negatively charged). For example, an O
atom (2,6) forms an O2– ion (2,8). See Figure 7.4b.
( 1
3 )
( ) Mg (2,8,2)
Mg2+ (2,8) 7.4a
(
) O (2,6) O2–
(2,8) 7.4b
Mg Mg
Figure 7.4 Formation of ions.
OO
loses 2e–
2e–
gains 2e–
2e–
magnesium atom magnesium ion
2+ 2–
oxygen atom oxide ion
(a) Formation of a magnesium ion. (b) Formation of an oxide ion.
Relation between ionic charge and group number of anelement
Metals in Groups I, II and III, the number of positive charges on
an ion is equal to its group number.
For non-metals in Groups V, VI and VII, however, the
number of negative charges on an ion is usually equal to ‘8
minus group number’. For example, an atom of oxygen (a
Group VI element) gains (8 – 6) or 2 electrons to get an octet,
forming an O2– ion.
All metals form ions: they usually form cations. Some non-
metals form ions — most of these are anions.
I II III
V VI VII
( VI )
(8 – 6) 2
O2–
—
—
46
Part II Microscopic World I I
Class practice 7.6 7.6
1. Some elements are shown in the incomplete Periodic Table
below. Write the formulae of the corresponding ions.1.
2. The atomic numbers of strontium and astatine are 38 and 85
respectively. Write the formula of (a) strontium ion (b) astatide
ion.
(Refer to the Periodic Table for atomic symbols and group
numbers.)
2. 38 85
(a) (b)
( )
3
Group Period
I II III IV V VI VII 0
2 Li Be
4
Na Mg Al
K Ca
N O F
S Cl
Br
Comparing properties of an atom and its ion
An atom and its ion have different physical and chemical
properties. This is because they have different numbers of
electrons and therefore different electronic arrangements.
Example 7.2Statements about atom and ion
This question consists of two separate statements. Decidewhether each of the two statements is true or false; if bothare true, then decide whether or not the second statement isa correct explanation of the first statement.
‘A neon atom and an oxide ion have similar chemicalproperties.’
‘A neon atom and an oxide ion have the same electronicarrangement.’
Solution
The first statement is false, while the second statement istrue. In fact, a neon atom and an oxide ion behavedifferently because they have different numbers of protons.
7.2
47
Chapter 7 Chemical bonding: ionic bonding
7.5 Chemical bonds 7.5
Atoms can join together, by chemical bonds, to form millions
of different compounds.
Types of chemical bonds
There are three main types of chemical bonds:
1. Ionic (or electrovalent) bond
2. Covalent bond (to be discussed in Chapter 8)
3. Metallic bond (to be discussed in Chapter 9)
1. ( )
2. ( )
3. ( )
7.6 Ionic bond and ionic substances 7.6
Ionic bond
Formation of ionic bond between sodium andchlorine
A sodium atom Na has the electronic arrangement 2,8,1. It canlose one electron to get the stable octet 2,8, forming a Na+
ion.
On the other hand, a chlorine atom Cl has the electronicarrangement 2,8,7. It can gain one electron to get the stable octet2,8,8, forming a Cl–
ion.
Thus when a sodium atom and a chlorine atom react, thesodium atom loses one electron to the chlorine atom. As a resultof this transfer of electron, two ions are formed. See Figure 7.5.
2,8,12,8 NaNa+
2,8,7 2,8,8 ClC l –
7.5
electron
sodium atom (Na)(loses one electron)
(both unstable, therefore reactive)
transfer
chlorine atom (Cl)(gains one electron)
sodium ion (Na+) chloride ion (Cl
–)
–
(both stable)
Figure 7.5 Electron ‘dot/cross’ diagrams showing the transfer of an electron from a sodium atom to a chlorine atom in theformation of sodium chloride, NaCl.
/ NaCl
Na ClNa Cl+
48
Part II Microscopic World I I
In the electron ‘dot/cross’ diagrams (or simply electron
diagrams) given here, ions are put inside square brackets with
the charge written at the top right-hand corner.
Ionic bond is the strong non-directional electrostatic force
of attraction between oppositely charged ions.
An ionic bond can be formed by the transfer of one or more
electrons from one atom (or group of atoms) to another.
(
)
( )
In the above reaction between sodium and chlorine, only
the outermost shell electrons are involved. This is true for most
chemical reactions. So for electron diagrams in the rest of the
book, only the outermost shell will be drawn.
Thus Figure 7.5 can be simplified as:
or even more simply,
electron
2,8,1
Na
+
transferCl Na Cl
–
2,8,7 2,8 2,8,8
+
Cl +Na Na Cl
–+
Formation of ionic bond between magnesium andfluorine
In the reaction between magnesium and fluorine, a magnesiumatom loses 2 electrons, while a fluorine atom gains 1 electron.Therefore, each magnesium atom must combine with twofluorine atoms.
7.5
49
Chapter 7 Chemical bonding: ionic bonding
electron
fluorine atom
F
2+
transferMg
–
magnesium atom
+ F+ F Mg F
–
fluorine atom fluoride ion magnesium ion fluoride ion
2,7 2,8,2 2,7 2,8 2,8 2,8
(unstable atoms) (stable ions)
(
)
(a) (b)
7.7
Draw electron diagrams (showing electrons in the outermost
shell only) to show the bond formation in (a) potassium sulphide
and (b) calcium bromide.
Class practice 7.7
7.7 Structures of solid ionic compounds 7.7
In sodium chloride, cations (Na+) and anions (Cl–) are attracted
together by ionic bonds. They are packed regularly, so that each
ion is surrounded by six ions of the opposite charge (Figure
7.6).
This packing continues until a continuous, three-
dimensional structure called giant ionic structure is formed.
(Na+)
(Cl–)
( 7.6)
Figure 7.6 Sodium chloride has a giant ionic structure. It consists of Na+
and Cl–
ions held together by ionic bonds.
Na+
Cl–
centre of Cl–
ion–
–
– – –
–
– –
–
–
–
+ +
+ +
+
+
+
++centre of Na
+ion
+
Sodium chloride crystals
chloride ion
sodium ion
50
Part II Microscopic World I I
Sodium chloride consists of ions, so it is called an ionic
compound. Magnesium fluoride is another ionic compound.
An ionic compound (or ionic substance) is a compound
which consists of ions.( )
7.8 Formulae and names of ioniccompounds
7.8
Formulae of ionic compounds
The formula of an ionic compound is a symbol indicating the
types and numbers of atoms present in the compound.
Let us take sodium chloride as an example. When sodium
atom (Na) loses an electron and becomes an ion, it has a
positive charge. The symbol for sodium ion is Na+. On the
other hand, the symbol for a chlorine atom is Cl. When it
accepts an electron, it becomes a chloride ion (Cl–). The overall
charge of the sodium chloride compound should be zero
because the positive charge on the sodium ion balances the
negative charge on the chloride ion. See Figure 7.7.
To work out the formula, the symbol of positive ion should
be written down first, followed by the negative ion. So the
formula for sodium chloride is NaCl. The formula does not
show the charges on the sodium or chloride ions as the charges
cancel each other when they combine.
(Na)
Na +
Cl
(Cl–)
7.7
Figure 7.7 The overall charge of sodium chloride is zero.
Na+ Cl
–
Na+
Cl–
Charge: +1 –1: +1
51
Chapter 7 Chemical bonding: ionic bonding
Example 7.3Writing the formulae of some ionic compounds
7.3
two potassium ions
one oxide ion
this number written after the brackets shows the number ofpotassium ions present
one magnesium iontwo nitrate ions
this number written after the bracketsshows the number of nitrate ions present
Give the formulae of the following ionic compounds.
(a) potassium oxide (b) magnesium nitrate
(c) sodium hydroxide (d) calcium hydroxide
(e) iron(III) sulphate
Solution
(a) Potassium oxide
K+ ion carries 1 positive charge; O2– ion carries 2negative charges. To have electrical neutrality, the ratioof K+ ions: O2– ions must be 2 : 1.
Thus the ionic formula of potassium oxide is as shownbelow:
(K+) 2 O2–
(not K+O2–, K2+O2–, (K+)2(O2–))
The formula is K2O, not KO, K2O.
(b) Magnesium nitrate
Mg2+ ion carries 2 positive charges; NO3– ion carries 1
negative charge. To have electrical neutrality, the ratioof Mg2+ ions: NO3
– ions must be 1 : 2.
Thus the ionic formula of magnesium nitrate is asshown below:
Mg2+ (NO3–) 2
(not Mg2+NO3–, Mg2+NO3
–2)
The formula is Mg(NO3)2, not MgNO3, Mg2(NO3),MgNO32.
(c) Sodium hydroxide
The ionic formula is Na+OH–, not Na+(OH–)2, Na+(OH–),Na+(OH)–.
The formula is NaOH, not Na(OH)2, Na(OH).cont'd
(a) (b)(c) (d)(e) (III)
(a)
K+ O2–
K+ O2–
2 1
(K+) 2 O2–
( K + O 2 – K 2 + O 2 –
(K+)2(O2–))
K2O KOK2O
(b)
Mg2+ NO3–
M g 2 +
NO3– 1
2
Mg2+ (NO3–) 2
( Mg2+NO3– Mg2+NO3
–2)
Mg(NO3)2
MgNO3 Mg2(NO3) MgNO32
(c)
Na+OH–
N a + ( O H – ) 2 N a + ( O H – )Na+(OH)–
N a O HNa(OH)2 Na(OH)
52
Part II Microscopic World I I
(d) Calcium hydroxide
The ionic formula is Ca2+(OH–)2, not Ca2+OH–,Ca2+OH–
2, Ca2+(OH)–2.
The formula is Ca(OH)2, not Ca2(OH), CaOH2.
(e) Iron(III) sulphate
The ionic formula is (Fe3+)2(SO42–)3, not Fe3+SO4
2–,Fe3+
2(SO42–)3.
The formula is Fe2(SO4)3, not FeSO4, (Fe)2(SO4)3.
➲ Try Chapter Exercise Q20
(d)
Ca2+(OH–)2
C a 2 + O H – C a 2 + O H –2
Ca2+(OH)–2
Ca(OH)2
Ca2(OH) CaOH2
(e)
(Fe3+)2(SO42–)3
Fe3+SO42– Fe3+
2(SO42–)3
Fe2(SO4)3
FeSO4 (Fe)2(SO4)3
➲ 20
(a) (II) (b)
(c) (d)
7.8
Write the chemical formula of each of the following compounds:
(a) Copper(II) chloride (b) Calcium sulphide
(c) Aluminium hydroxide (d) Ammonium carbonate
Class practice 7.8
A short cut to predict formulae of ionic compounds
There is a short cut to predict the formula of an ionic
compound. Let us take the example of magnesium fluoride.
Predicting the formulae of ionic compounds
Step 1 Write the formulae of the two ions involved side by side.
Mg2+ F–
Step 2 Highlight the number of the charge on each ion.
Mg 2 + F 1 –
Step 3 Take the number of the charge on each ion across to theother.
Mg 2 + F 1 –
= Mg1 F2
Step 4 Combine the symbols and simplify the ratio.
MgF2
(Omit the number 1 for Mg)
Problem-solving strategy
1
Mg2+ F–
2
Mg 2 + F 1 –
3
Mg 2 + F 1 –
= Mg1 F2
4
MgF2
( Mg 1 )
53
Chapter 7 Chemical bonding: ionic bonding
Study more examples:
Al 3 + O 2 – Al2O3
Fe 3 + SO42 – Fe2(SO4)3
Ca 2 + O 2 – Ca2O2 CaO
Aluminium oxide
Iron(III) sulphate
Calcium oxide
(Note: The formula of calcium oxide is CaO but not Ca2O2. This
is because the formula of an ionic compound expresses the
simplest whole number ratio of the ions present. Therefore, the
ratio of 2 : 2 must be simplified to 1 : 1.)
Al 3 + O 2 – Al2O3
Fe 3 + SO42 – Fe2(SO4)3
Ca 2 + O 2 – Ca2O2 CaO
( CaO
Ca 2O 2
2 2 1 1 )
(a) (b)
(c) (II) (d)
7.9
Using the short-cut method, predict the chemical formula of
each of the following compounds:
(a) Magnesium hydroxide (b) Sodium oxide
(c) Lead(II) sulphate (d) Potassium dichromate
Class practice 7.9
Naming ionic compounds
We can name ionic compounds based on the following two
rules:
1. The cation is named first, followed by the anion. The word
‘ion’ is omitted. For example,
1.
(Na+)
(CO32–)
(Na2CO3)
54
Part II Microscopic World I I
Cation
Al3+
NH4+
Ca2+
Cu+
Cu2+
Pb2+
Al2(SO4)3
(NH4)2CO3
Ca(NO3)2
Cu2O
CuO
PbBr2
aluminium sulphate
ammonium carbonate
calcium nitrate
copper(I) oxide (I)
copper(II) oxide (II)
lead(II) bromide (II)
SO42–
CO32–
NO3–
O2–
O2–
Br–
Anion Formula of compound Name of compound
2. Some ionic compounds contain water of crystallization.
The number of molecules of water of crystallization (n) has
to be added at the end of the name as: -n-water. For
example, Na2CO3 · 10H2O is called sodium carbonate-10-
water.
2.
(n)
Na2CO3 ·
10H2O
(a) Ca(NO3)2 (b) FeCl3
(c) ZnSO4 · 7H2O (d) Cu(OH)2
7.10
Name the following compounds:
(a) Ca(NO3)2 (b) FeCl3
(c) ZnSO4 · 7H2O (d) Cu(OH)2
Class practice 7.10
55
Chapter 7 Chemical bonding: ionic bonding
1. anion 37
Page
3. chemical bond 47
6. giant ionic structure 49
7. ionic bond 48
8. ionic compound 50
4. electron ‘dot/cross’ diagram / 48
2. cation 37
5. formula 41
10. polyatomic ion 37
11. simple ion 37
12. transfer of electron 47
9. migration of ion 39
13. water of crystallization 54
KK ee yy tt ee rr mm ss
56
Part II Microscopic World I I
SS uu mm mm aa rr yy7.1 Formation of ions from atoms
1. Noble gases have great stability because their atoms have either an of electrons(8 electrons in the outermost shell), or a of electrons (2 electrons in the only oneoccupied shell) as in helium. Other atoms can also gain great stability if they can get an octet (orduplet).
2. An is an atom or a group of atoms having an overall electric charge. A is derived from a single atom. A is derived from a group of atoms. Positive ions (e.g. Na+,NH4
+) are called ; negative ions (e.g. Cl–, MnO4–) are called .
7.2 Colours and migration of ions
3. Colours of some ions in aqueous solution are listed in Table 7.1 on p.38.
4. Colours of some gemstones are due to traces of ions. Refer to Table 7.2 on p.39.
7.3 Formulae of ions
5. A represents the smallest unit (using chemical symbols and numbers) of asubstance or species under some specified conditions.
6. Names and formulae of common ions are listed in Table 7.3 on p.43.
7.4 Elements and ions
7. All metals form ions: they usually form . Some non-metals form ions — most ofthese are .
8. For metals in Groups I, II and III, the number of charges on an ion is equal to itsgroup number. For non-metals in Groups V, VI and VII, the number of chargeson an ion is usually equal to ‘8 minus group number’.
7.5 Chemical bonds
9. Atoms can join together by chemical bonds to form different compounds. There are three maintypes of chemical bonds, namely, bonds, bonds and
bonds.
57
Chapter 7 Chemical bonding: ionic bonding
7.6 Ionic bond and ionic substances
10. is the strong non-directional electrostatic force of attractionbetween oppositely charged ions.
11. When a metal (which tends to electrons) and a non-metal (which tends toelectrons) combine, they do so by the transfer of electrons, forming ions. The
ions are held together by ionic bonds.
For example,
7.7 Structures of solid ionic compounds
12. An (or ionic substance) is a compound which consists of ions.
7.8 Formulae and names of ionic compounds
13. The formulae of ionic compounds can often be predicted using a short-cut method:
X a Y b ⇒ XbYa
(where a, b = ionic charge)
e.g. Zn 2 + NO31 – ⇒ Zn(NO3)2
electron
2,8,1
Na
+
transfer
Electron diagram of sodium chloride
Cl Na Cl
–
2,8,7 2,8 2,8,8
+
58
Part II Microscopic World I I
Molecules in compounds and elements
Molecules in compounds
Compounds made up of non-metals only usually consist of
neutral particles called molecules.
Notice that a molecule of a compound consists of atoms of
different kinds. For example, carbon dioxide molecules consist
of two kinds of atoms (carbon and oxygen). Carbon dioxide
CO2, ammonia NH3, methane CH4 and hydrogen chloride HCl
are all molecules (Figure 8.1).
( ) CO2
NH3 CH4 HCl
( 8.1)
8.1 Covalent bonding and covalentsubstances
8.1
Figure 8.1 Molecules ofsome compounds.
carbon dioxide ammonia methane hydrogen chloride
O C CO H H
H H
H Cl
H
HH
N
Molecules in elements
Elements consist of either atoms or molecules. All metals consist
of atoms. All non-metals (except carbon) consist of discrete
(separate) molecules. For example, chlorine gas consists of
discrete chlorine molecules.
The number of atoms in a molecule of an element is called
its atomicity. In gaseous elements, the atomicity of chlorine
(Cl2), nitrogen (N2), oxygen (O2), fluorine (F2) and hydrogen
(H2) is 2; that of noble gases (e.g. Ar) is 1; that of ozone (O3) is 3.
In solid elements, the atomicity of yellow phosphorus (P4) is 4;
that of sulphur (S8) is 8. Thus argon (Ar) is monoatomic, oxygen
(O2) is diatomic, ozone (O3) triatomic and so on.
We can now define molecule.
( )
( )
(Cl2)
(N2) (O2) (F2) (H2)
2 ( Ar) 1
(O3) 3
(P4) 4 (S8)
8 (Ar)
(O2) (O3)
59
Chapter 8 Chemical bonding: covalent bonding
A molecule is the smallest part of an element or a
compound which can exist on its own under ordinary
conditions.
1.
B r 2 K+
B r Z n ( O H ) 2
C6H12O6 Ne Na NH3 CaO
2.
(a) (b) (c)
(d) (e) (f)
( )
8.1
1. Which of the following represent a molecule?
Br2, K+, Br, Zn(OH)2, C6H12O6, Ne, Na, NH3, CaO
2. Write the formulae for the following elements:
(a) neon (b) hydrogen (c) sodium
(d) nitrogen (e) fluorine (f) magnesium
(Refer to the Periodic Table if necessary.)
Class practice 8.1
Covalent bonding
Covalent bond formation in a chlorine molecule
A molecule usually consists of a number of atoms chemically
joined together.
Take the example of chlorine gas. The chlorine atom, Cl, is
very unstable. Its outermost shell contains only seven electrons
— one electron less than an octet. Electron transfer between
chlorine atoms is impossible here. This is because they all tend
to gain electrons, and no one would lose them. But by sharing
of electrons (one electron from each chlorine atom) in the
outermost shell, a chlorine molecule Cl2 is formed. In the
molecule, each chlorine atom has a stable octet (Figure 8.2).
Cl
( )
C l 2
( 8.2)
Figure 8.2 Electron diagrams showing the sharing of two electrons in the formation of a chlorine molecule(only the outermost shell electrons are shown).
( )
Cl Cl Cl Cl+
a shared pair of electrons forms a single covalent bond
electron
sharing
chlorine atom (Cl)(Cl)
chlorine atom (Cl)(Cl)
chlorine molecule (Cl2)(Cl2)
2,8,7 2,8,7 2,8,8 2,8,8(both unstable)
( )(more stable)
( )
60
Part II Microscopic World I I
Covalent bond is the strong directional electrostatic
attraction between the shared electrons (negatively
charged) and the two nuclei (positively charged) of the
bonded atoms.
A covalent bond is formed by the sharing of outermost shell
electrons between two atoms.
The molecular formula of a molecular substance is the
formula which shows the actual number of each kind of
atoms in one molecule of the substance.
The structural formula of a molecular substance is the
formula which shows how the constituent atoms are joined
up in one molecule of the substance.
A shared pair of electrons (or bond pair) makes a single
covalent bond. It is often represented by a stroke (–) between
the atomic symbols. So a chlorine molecule Cl2 can be written
as Cl–Cl. (The ‘–’ also indicates the direction of the
electrostatic attraction.) Cl2 is the molecular formula of
chlorine, while Cl–Cl is the structural formula of chlorine.
When we say the ‘formula’ of a molecular substance, we
usually mean its ‘molecular formula’.
Covalent bond formation in some molecules
Table 8.1 gives electron diagrams to show the covalent bond
formation in some simple molecules. All of them are molecules
of covalent substances.
It should now be obvious that a chlorine molecule must be
Cl2, and cannot possibly be Cl, Cl3 or Cl4.
A covalent substance is a non-ionic substance in which the
atoms are held together by covalent bonds.
( )
( )
( )
Cl2 Cl Cl (
) Cl2
Cl Cl
8.1
Cl2
Cl Cl3 Cl4
61
Chapter 8 Chemical bonding: covalent bonding
Table 8.1 Electron diagrams to show the formation of some simple molecules (only the outermost shell electrons are shown).( )
Electron diagrams to show covalent bond formation Molecular formula Structural formula
H H H H
2 hydrogen atoms2
Cl H ClH
1 hydrogen chloride molecule1
C
H
H
H H C
H
H
H H
1 carbon atom + 4 hydrogen atoms1 + 4
1 methane molecule1
NH H
H
NH H
H
1 nitrogen atom + 3 hydrogen atoms1 + 3
1 ammonia molecule1
OH H OH H
1 oxygen atom + 2 hydrogen atoms1 + 2
1 water molecule1
H2
HCl
H H
a single covalent bond
H Cl
a bond pair of electrons
a lone pairof electrons
CH4 H C
H
H
H
NH3 H N
H
H
H2O H O H
O C O O C O
1 carbon atom + 2 oxygen atoms1 + 2
1 carbon dioxide molecule1
CO2 O C O
a double covalent bond
N N N N
2 nitrogen atoms2
1 nitrogen molecule1
N2N N
a triple covalent bond
1 hydrogen molecule1
1 hydrogen atom + 1 chlorine atom1 + 1
62
Part II Microscopic World I I
When non-metal atoms combine with each other, there is
usually a sharing of electrons, forming covalent bonds.
Rules for forming covalent bonds
Table 8.2 lists out some rules for forming covalent bonds,
illustrated with a few examples. (Refer to Table 8.1 at the same
time.)
8.2
( 8.1 )
Rules Examples
(3) 2 atoms may share between them
• 1 electron pair (to form asingle covalent bond)
( )or • 2 electron pairs (to form a
double covalent bond)• ( )
or • 3 electron pairs (to form atriple covalent bond)
• ( )
contains 4 single covalent bonds;
O=C=O contains 2 double covalent bonds;O=C=O
contains 1 triple covalent bond.N NN N
H (CH4)H C
H
H
(1) An atom involved in covalent bondformation contributes n electron(s) forsharing.
n• For hydrogen atoms, n = 1
n = 1• For other atoms, n = 8 – group no.
of the elementn = 8 –
A hydrogen atom contributes 1 electron for sharing;
a carbon atom (Group IV) contributes (8 – 4) or 4 electrons for sharing;( IV ) (8 – 4) 4
a nitrogen atom (Group V) contributes (8 – 5) or 3 electrons forsharing;
( V ) (8 – 5) 3an oxygen atom (Group VI) contributes 2 electrons for sharing;
( VI ) (8 – 6) 2a fluorine atom (Group VII) contributes 1 electron for sharing
( VII ) (8 – 7) 1
(2) • For hydrogen and Group VIIelements, an atom shares electronswith one other atom in covalentbond formation.
V I I
• For other elements, an atom mayshare electrons with one or moreother atoms.
In , a chlorine atom shares electrons with a hydrogen atom.
In , a nitrogen atom shares electrons with another nitrogenatom.
In , a nitrogen atom shares electrons with 3 hydrogen
atoms.
H Cl
H ClN N
N N
HH N
H
HH N
H
63
Chapter 8 Chemical bonding: covalent bonding
Rules Examples
(4) A shared pair of electrons is known asa bond pair. Some atoms in a molecule may haveunshared pairs of outermost shellelectrons — known as lone pairs.
The nitrogen atom in an NH3 molecule has 3 bond pairs and 1 lonepair.
NH3
In a H2O molecule, the oxygen atom has 2 bond pairs and 2 lonepairs.
H2O
HH N
H
XX
HH OXX
XX
lone pair
bond pair
lone pair
lone pair bond pair
Table 8.2 Rules for forming covalent bonds.
Example 8.1Identifying some common substances
Given the names and formulae of the following substances:
tetrachloromethane (CCl4), silver (Ag), ammonium nitrate(NH4NO3), ethanoic acid (CH3COOH), lithium hydroxide(LiOH), heptane (C7H16), iodine (I2)
Which of them are
(a) ionic compounds
(b) covalent substances
(c) covalent compounds?
Solution
(a) Ammonium nitrate, lithium hydroxide
(b) Tetrachloromethane, ethanoic acid, heptane, iodine
(c) Tetrachloromethane, ethanoic acid, heptane
8.1
(CCl4) (Ag)(NH4NO3) (CH3COOH)
(LiOH) (C7H16) (I2)
(a)
(b)
(c)
(a)
(b)
(c)
64
Part II Microscopic World I I
1.
2. (a) (i)
()
(ii)
(b) (a)
(i)
(ii)
8.2
1. Fill in the blanks:
Metals tend to electrons, while non-
metals tend to or
electrons in chemical reactions.
2. (a) (i) Draw an electron diagram (showing electrons in
the outermost shell only) for a molecule of the
compound formed between nitrogen and
chlorine.
(ii) Find the number of bond pairs and lone pairs on
the nitrogen atom in this molecule.
(b) Give the
(i) molecular formula,
(ii) structural formula of the molecule in (a).
Class practice 8.2
Dative covalent bond
Atoms which have lone pairs of electrons may form dative
covalent bonds. Let us consider the following examples.
Dative covalent bond in ammonium ion (NH4+)
When ammonia reacts with hydrogen chloride to form
ammonium chloride, a dative covalent bond is formed between
the lone pair of electrons on the N atom in NH3 and a H+ ion
from HCl (Figure 8.3). The symbol ‘ ’ is used to represent
the dative covalent bond.
A dative covalent bond (or coordinate bond) is a bond
formed between two atoms where both electrons of the
shared pair are contributed by the same atom.
(NH4+)
NH3 N
HCl H+
( 8.3)
( )
65
Chapter 8 Chemical bonding: covalent bonding
The ammonium ion (NH4+) has an overall charge of +1
distributed all over the structure. Thus ammonium chloride
(NH4Cl) contains ionic bond (between NH4+ and Cl– ions) and
four covalent bonds (four N–H bonds) — three of the N–H
bonds are normal covalent bonds and one is dative covalent
bond.
It should be noticed that dative and normal covalent bonds
differ only in the way they are formed. Once a dative covalent
bond has formed, it cannot be distinguished from a normal
covalent bond.
Dative covalent bond in hydronium ion (H3O+)
When an acid is dissolved in water, hydrogen ions H+ are
formed. Take hydrochloric acid as an example. When hydrogen
chloride gas is passed into water, hydrogen chloride molecules
break down to give hydrogen ions H+ and chloride ions Cl–.
Each H+ ion is attracted to the unshared electrons of oxygen
atom of a water molecule, forming a dative covalent bond. A
more stable ion, hydronium ion H3O+, is obtained as a result.
See Figure 8.4.
H
NH
H
H
H
ClClHH
H
H
N
H
H
H N H Cl ClN HH
H
H
or
Figure 8.3 Electron ‘dot/cross’ diagram showing formation of ammonium chloride.
dative covalent bond
ammonium ion
(NH4+)
+1 (NH4Cl)
( NH4+ Cl– )
( N–H ) N–H
(H3O+)
H+
H +
Cl – H +
H+
H3O+ 8.4
66
Part II Microscopic World I I
H
H O H Cl ClHH
H
O
ClO HH
H
H ClOH
H
or
Figure 8.4 Electron ‘dot/cross’ diagram showing formation of hydronium ion.
Ionic bonding and covalent bonding in comparison
Ionic bonding and covalent bonding have many differences.
For example, an ionic bond is non-directional, while a covalent
bond is directional.
However, the two types of bonding have some common
features: the bonds are both strong, and the bonding forces are
electrostatic in nature.
dative covalent bond
hydronium ion
✘ The constituents of ammonium nitrate (NH4NO3) are all non-metals, so it is considered as a covalent compound.
✔ Although ammonium nitrate is made up of non-metals only, it isan ionic compound. It consists of ammonium ion (NH4
+) and
nitrate ion (NO3–).
Check your concept
✘ (NH4NO3)
✔
(NH4+) (NO3
–)
67
Chapter 8 Chemical bonding: covalent bonding
8.2 Prediction of formulae for covalentcompounds
8.2
We have used the ‘noble gas approach’ to work out the electron
diagrams of the molecules of a few covalent compounds. From
the electron diagram of a compound, we can deduce its
molecular formula and structural formula.
Alternatively, we can use a short cut similar to the one
used for ionic compounds. A few examples are shown in Table
8.3.
8.3
Compound Molecular formula Structural formula
Hydrogen sulphide
Tetrachloromethane
Ammonia
Carbon dioxide
Table 8.3 Predicting formulae of hydrogen sulphide, tetrachloromethane, ammonia and carbon dioxide (using a short cut).( )
H S H2S1 H2S21
C Cl C1Cl4 CCl4
14
N H N1H3 NH3
13
C O C2O4 CO2
24
H S H
Cl C
Cl
Cl
Cl
H HN
H
O C O
8 .3
(a) (b)
(c) (d)
8.3
Using the short cut as shown in Table 8.3, predict the molecular
formula for the compound formed between
(a) carbon and fluorine (b) hydrogen and oxygen
(c) phosphorus and hydrogen (d) silicon and chlorine.
Class practice 8.3
8.3 Particles that make up matter — asummary
8.3 —
Three types of particles that make up matter
The different types of particles are atoms, molecules and ions.
Molecules and ions, however, come from atoms.
Study the following example.
68
Part II Microscopic World I I
Example 8.2Identifying types of particles that make up substances
Complete the table below.
Solution
(a) element; (b) element; (c) compound; (d) molecules; (e) names; (f) CO; (g) CO2; (h) molecule; (i) mixture.
(a) (b) (c) (d) (e) (f) CO(g) CO2 (h) NaCl (i)
8.2
Substance/species Constituent particle(s) Formula Remarks
Nitrogen molecule N2nitrogen is an (a) ________________ .
(a)
Magnesium atom Mgmagnesium is an (b) ________________ .
(b)
Water H2Owater is a (c) _________ made of (d) __________ .
(c) (d)
Carbon monoxideCO
the formulae of some compounds can be guessedfrom their (e) _____________ : thus the formulaof carbon monoxide is (f) _____________ ; that ofcarbon dioxide is (g) _____________ .
(e) (f)
(g)
Carbon dioxideCO2
Sodium chloridedifferenttypes of ions
NaCl
there is no sodium chloride (h) ________________to represent the compound sodium chloride.
(h)
Hydroxide ionion OH– OH– is an ion OH–
Air notapplicable
air has no formula because it is a (i) ____________ of many substances.
(i)
different typesof molecules
molecule
molecule
molecule
Table 8.4 summarizes the constituent particles in various
substances.8 . 4
69
Chapter 8 Chemical bonding: covalent bonding
Table 8.4 Constituent particles of various substances.
PURESUBSTANCES
elements
compounds
metals
non-metals
compoundsmade up ofnon-metalsonly
compoundsmade up ofmetal(s) and non-metal(s)
atoms
molecules(exception:carbon)
( )
usuallymolecules
ions
Constituentparticles
copper (Cu)(Cu)
argon (Ar) chlorine (Cl2)sulphur (S8)
(Ar) (Cl2)(S8)
water (H2O)ammonia (NH3)
(H2O)(NH3)
potassium oxide (K2O)sodium chloride (NaCl)
(K2O)(NaCl)
Examples
(a) CHCl3 (b) Ar
(c) Cr2O72–
(d) Mg
(e) S8 (f) Ba2+
(g) I2 (h) P
8.4
Decide whether the following formulae stand for an atom, a
molecule or an ion.
(a) CHCl3 (b) Ar (c) Cr2O72–
(d) Mg
(e) S8 (f) Ba2+
(g) I2 (h) P
Class practice 8.4
8.4 Relative molecular mass and formulamass
8.4
For elements and compounds consisting of molecules, relative
molecular mass is the mass of one molecule of it on the 12C =
12.000 00 scale.
-12 =
12.000 00
70
Part II Microscopic World I I
Relative molecular mass can also be called molecular mass.Relative molecular mass carries no units.
Learning tip
Formula mass carries no units.
Learning tip
Relative molecular mass of = Sum of relative atomic masses
an element or a compound of all atoms present in a
molecule of the substance
For example, water (H2O) would have a relative molecular
mass of 1.0 � 2 + 16.0 = 18.0.
Some compounds (such as ionic compounds) do not
consist of molecules. For these, we use formula mass.
The formula mass of a substance (or species) is the mass of
one formula unit of it on the 12C = 12.000 00 scale.
Formula mass of a = Sum of relative atomic masses
substance (or species) of all atoms present in a formula
unit of the substance
(H2O)
1.0 � 2 + 16.0 = 18.0
( )
( ) -12 =
12.000 00
( )
Formula mass is a general term applicable to all substances
(or species) with a formula. In comparison, relative molecular
mass only applies to molecular substances. See Example 8.3.( )
8.3
71
Chapter 8 Chemical bonding: covalent bonding
Example 8.3Determining the formula masses of some substances/species
Calculate the formula mass of
(a) C6H12O6 (b) SO42– (c) Al2(SO4)3
SolutionC6 H12 O6
(a) Formula mass of = 12.0 � 6 + 1.0 � 12 + 16.0 � 6C6H12O6 = 180.0
Note: We can regard the formula mass of C6H12O6 as therelative molecular mass because the compoundactually consists of molecules.
S O4
(b) Formula mass of SO42– = 32.1 + 16.0 � 4
= 96.1
Al2 (SO4)3
(c) Formula mass of = 27.0 � 2 + (32.1 + 16.0 � 4) � 3 Al2(SO4)3 = 342.3
8.3
(a) C6H12O6 (b) SO42–
(c) Al2(SO4)3
(a) C6H12O6
C6 H12 O6
= 12.0 � 6 + 1.0 � 12 + 16.0 � 6 = 180.0
C6H12O6
S O4
(b) SO42– = 32.1 + 16.0 � 4
= 96.1
(c) Al2 (SO4)3
Al2 (SO4)3
= 27.0 � 2 + (32.1 + 16.0 � 4) � 3 = 342.3
1.
(a) CH4 (b) C2H6
(c) C12H22O11
2.
(a) NaCl (b) C2H6
(c) CO32–
(d) Cu(NO3)2 · 3H2O
8.5
1. What is the relative molecular mass of
(a) CH4 (b) C2H6 (c) C12H22O11?
2. Calculate the formula mass of
(a) NaCl (b) C2H6 (c) CO32–
(d) Cu(NO3)2 · 3H2O.
Class practice 8.5
72
Part II Microscopic World I I
KK ee yy tt ee rr mm ss
1. atomicity 58
Page
3. covalent bonding 60
6. formula mass 70
7. formula unit 70
8. lone pair 63
9. molecular formula 60
10. relative molecular mass 69
11. sharing of electrons 59
12. single covalent bond 60
13. structural formula 60
14. triple covalent bond 62
4. covalent substance 60
2. bond pair 60
5. double covalent bond 62
73
Chapter 8 Chemical bonding: covalent bonding
SS uu mm mm aa rr yy8.1 Covalent bonding and covalent substances
1. A is the smallest part of an element or a compound which can exist on its ownunder ordinary conditions.
2. Compounds made up of non-metals only usually consist of molecules.
Elements are made up of either atoms or molecules. All metals consist of . Allnon-metals (except carbon) consist of discrete .
3. Molecules can be represented by to show their shapes.
4. A is formed when one or more pairs of outermost shellelectrons are shared between two atoms. For example,
5. is the strong directional electrostatic attraction between theshared electrons and the two nuclei of the bonded atoms.
6. A shared pair of electrons ( ) makes a covalentbond, e.g. H – Cl.
2 shared pairs of electrons make a covalent bond, e.g. O = C = O.
a double covalent bond
3 shared pairs of electrons make a covalent bond, e.g. N ≡ N.
7. A (or coordinate bond) is a bond formedbetween two atoms where both electrons of the shared pair are contributed by the same atom.
8.2 Prediction of formulae for covalent compounds
8. The of a molecular substance shows the actual number ofeach kind of atoms in one molecule of the substance, e.g. CH4.
9. The of a molecular substance shows how the constituentatoms are joined up in one molecule of the substance, e.g.
ClHClH
1 hydrogen atom + 1 chlorine atom 1 hydrogen chloride molecule
H C
H
H
H
74
Part II Microscopic World I I
10. Some atoms have unshared pairs of outermost shell electrons. These are known as ,e.g.
11. The formulae of covalent compounds can often be predicted using a short-cut method:
X a Y b ⇒ XbYa
(where a, b = number of electrons contributed for sharing)
e.g. Si 4 H 1 ⇒ SiH4
8.3 Particles that make up matter — a summary
12. All matter is made up of particles: atoms, molecules or .
8.4 Relative molecular mass and formula mass
13. The 12C = 12.000 00 scale is used for comparing of atoms.
14. of an element or a compound= Sum of relative atomic masses of all atoms present in a molecule of the substance
15. of a substance (or species)= Sum of relative atomic masses of all atoms in a formula unit of the substance (or species)
OH H
a lone pair of electrons
1 water molecule
75
Chapter 9 Structures and properties of substances
Introduction
9.1 Structure of substances 9.1
The study of structures is important, since physical
properties of a substance are closely related to its structure.
The structure of a substance is a description of what its
constituent particles are, and about how they are arranged
or packed together.
Classification of substances according to structure
Under ordinary conditions, all substances exist as either
molecular structures or giant structures.
Molecular structures
There are two types of molecular structures, depending on the
molecular size:
• Simple molecular structures, which may be solids, liquids
or gases;
• Macromolecules, which are all solids at room conditions.
Giant structures
In a giant structure, all particles (trillions of atoms or ions) are
joined together by strong chemical bonds. A continuous giant
lattice is formed, in which no discrete molecules exist. All
substances having giant structures are solids at room
conditions.
A classification of substances according to structure,
together with some examples, is shown in Figure 9.1.
•
•
(
)
9.1
( )
76
Part II Microscopic World I I
EXAMPLES
Elements Compounds Non-metals Metals Covalent Ionic
hydrogen H2
iodine I2
H2
I2
water H2Ocarbondioxide CO2
H2OCO2
polyethene —(CH2CH2 ) —n
—(CH2CH2 ) —n
Molecularstructures
sodiumchloride NaCl
NaCl
diamond,
graphite
(different forms
of carbon)
( )
silicon(IV)oxide SiO2
(IV)SiO2
Giantstructures
Simplemolecularstructures
Macro-molecules
Giant ionicstructures
Giantcovalentstructures
copper Cuiron Fe
CuFe
Giantmetallicstructures
SUBSTANCES
Figure 9.1 Classif ication ofsubstances according to structure.
(a) (b)
9.1
Name the structures possible for (a) non-metal elements (b)
covalent compounds.
Class practice 9.1
For polyethene, the formula is represented by –(CH2CH2)–, where nis a whole number from 100 to 30 000. Each molecule is very large(hence called macromolecule); it consists of many, usuallythousands, of –CH2CH2– groups joined together.
Learning tip
–(CH2CH2)n– n 100 30 000(
) ( ) –CH2CH2–
n
77
Chapter 9 Structures and properties of substances
9.2 Simple molecular structures 9.2
Most non-metals and covalent compounds are composed of
simple, discrete molecules. These substances have a simple
molecular structure. The atoms within a molecule are strongly
bonded together (by covalent bonds). However, each molecule is
attracted to neighbouring molecules by weak intermolecular
forces only.
Structure of carbon dioxide
Each carbon dioxide molecule consists of one carbon atom and
two oxygen atoms covalently bonded together. Under room
conditions, carbon dioxide is a gas. Since weak intermolecular
forces (called van der Waals’ forces) always exist between
molecules, each CO2 molecule is attracted to neighbouring
molecules.
In general, the larger the molecular size, the greater will be
the van der Waals’ forces between molecules.
When carbon dioxide gas is placed under temperatures
below –78.5°C, it changes to solid called dry ice directly without
going through the liquid state. The structure of dry ice is
shown in Figure 9.2.
1 2
( )
CO2
–78.5°C
9.2
Figure 9.2 The structure of dry ice. The molecules are held by van der Waals' forces in thestructure.
78
Part II Microscopic World I I
Structure of iodine
In an iodine crystal, I2 molecules are packed closely together,
but they are still discrete molecules. The molecules are held by
van der Waals’ forces. See Figure 9.3.
Figure 9.3 The crystal structure ofiodine. indicates the position ofan I2 molecule. Here the moleculesare packed in a regular pattern.Repetition of this pattern trillions oftimes would result in a crystal.
I 2
9.3
I2
(
)
9.2
Explain why iodine is a solid, bromine is a liquid, while chlorine
and fluorine are gases at room conditions. (Hint: You may
answer the question according to the van der Waals’ forces
between the molecules.)
Class practice 9.2
Properties of simple molecular substances
1. Simple molecular substances have low melting points and
boiling points. Because the molecules are held together only
by weak intermolecular forces (such as van der Waals’
forces), little heat energy is needed to separate the
molecules.
1.
( )
A volatile liquid evaporates quickly under room conditions.
Learning tip
79
Chapter 9 Structures and properties of substances
2. Simple molecular solids are soft. Intermolecular forces are
weak. It is easy to separate molecules and break down the
crystal structure.
3. They are usually insoluble in water, but soluble in non-aqueous
solvents such as methylbenzene and heptane.
4. They are non-conductors of electricity, whether as solids,
liquids or in aqueous solution. This is because they do not
contain ions or freely moving electrons to conduct
electricity.
Note: The aqueous solutions of a few molecular substances
conduct electricity and can be electrolysed. This is because
mobile ions are formed during the dissolution process.
Examples include sulphuric acid and ammonia.
2.
3.
( )
4.
Solvents other than water are called non-aqueous solvents.
Learning tip
(a)
(b)
(c)
(d) (i) (ii) (
)
9.3
Answer the following questions.
(a) Is sulphur high-melting or low-melting?
(b) Are sulphur crystals hard?
(c) Does molten sulphur conduct electricity?
(d) Is sulphur soluble in (i) water (ii) carbon disulphide, a non-
aqueous solvent?
Class practice 9.3
9.3 Macromolecules 9.3
Macromolecules are very large molecules, each containing
thousands of atoms. Examples are plastics, proteins and some
carbohydrates like starch. ( )
80
Part II Microscopic World I I
9.4 Giant ionic structures 9.4
An ionic compound is usually formed by combining a metal
with a non-metal. Ionic crystals consist of positive and negative
ions held together by strong non-directional electrostatic
attractions (ionic bonds). The ions are regularly packed to form
a continuous, three-dimensional giant ionic structure. (There
are no discrete molecules.)
Structure of caesium chloride
Since caesium ion is larger in size than the sodium ion, each
caesium ion is surrounded by eight chloride ions and each
chloride ion is in turn surrounded by eight caesium ions.
Therefore, the structure of caesium chloride CsCl is different
from that of sodium chloride. See Figure 9.4.
( )
( )
8 8
CsCl
9.4
Figure 9.4 Caesium chloride has a giantionic structure. It consists of Cs
+and Cl
–
ions held together by ionic bonds.
Cs+
Cl–
Cs+
Cl–
In the structure of sodium chloride, each sodium ion is surroundedby six chloride ions.
Learning tip
6
Properties of ionic compounds
1. All ionic compounds are solids. The oppositely charged ions
are attracted together by strong ionic bonds.
2. They usually have high melting points and boiling points. A
lot of heat energy is required to overcome the strong
attractive forces (ionic bonds) between ions in melting and
boiling.
1.
2.
( )
or
81
Chapter 9 Structures and properties of substances
3. Most of them are soluble in water, but insoluble in non-
aqueous solvents such as heptane.
4. They conduct electricity when molten or in aqueous solution.
They are non-conductors when solid. This is because in
solid state, the ions present are not mobile; when molten or
in aqueous solution, the ions become mobile and can
conduct electricity. They are therefore electrolytes.
3.
4.
An electrolyte is a compound which, when molten or in aqueoussolution, conducts electricity, and is decomposed at the same time.
Learning tip
Example 9.1Statements about ionic bonds and covalent bonds
This question consists of two separate statements. Decidewhether each of the two statements is true or false; if bothare true, then decide whether or not the second statement isa correct explanation of the first statement.
‘Sodium chloride is a high-melting solid, whereas chlorine isa gas.’
‘Ionic bonds are strong while covalent bonds are weak.’
Solution
The first statement is true. The second statement is false,because both ionic and covalent bonds are strong.
Sodium chloride is an ionic compound. The strong ionicbonds between ions must be overcome before thecompound can melt. This requires a lot of heat energy.Hence, sodium chloride is a high-melting solid.
Chlorine is a molecular substance. To boil chlorine, onlythe weak van der Waals’ forces between chlorine moleculesmust be overcome; the strong covalent bond within eachmolecule is not broken. Thus chlorine is a gas.
9.1
82
Part II Microscopic World I I
9.5 Giant covalent structures 9.5
In a few elements and compounds, non-metal atoms are joined
by covalent bonds to form a giant network, called a giant
covalent structure. Covalent bonds extend throughout the
whole structure. There are no discrete molecules.
Carbon atoms can be joined in two different ways, to form
diamond or graphite. Diamond and graphite have very
different physical properties because of their different
structures.
Structure and properties of diamond
Diamond is one form of carbon. It has a giant covalent
structure. Each carbon atom is covalently bonded to four other
carbon atoms, forming a three-dimensional giant network. See
Figure 9.5.
9.5
carbon atoms
covalentbonds
Figure 9.5 The three-dimensional structureof diamond.
To break the structure, numerous very strong covalent
bonds between carbon atoms must be broken. This explains the
extreme hardness and very high melting and boiling points of
diamond.
Diamond cannot conduct electricity, because it contains no
ions or freely moving electrons to carry electric charges.
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Chapter 9 Structures and properties of substances
Two main uses of diamond:
(a) Jewellery
(b) Diamond cutter (used for cutting glass)
Structure and properties of graphite
In graphite, the carbon atoms are arranged in flat, parallel layers.
Each layer contains many six-membered carbon rings (Figure
9.6).
(a)
(b) ( )
( 9.6)
Figure 9.6 The structure of graphite.(The ‘lead’ pencil is graphite mixedwith some clay.)
(
)
strong covalent bonds (within layers)( )
weak van der Waals’ forces(between layers)
()
Each carbon atom is covalently bonded to only three other
carbon atoms in its layer, and one outer electron of each carbon
atom is ‘free’. Those electrons are not attached to any particular
atoms but belong to the whole structure (i.e. the electrons are
delocalized). They are free to move from one six-membered
carbon ring to the next within a layer. Thus, graphite can
conduct electricity.
Since only van der Waals’ forces exist between adjacent
layers, these weak forces make the graphite crystal easy to
cleave, and explain its softness and lubricating property. On the
other hand, graphite has a very high melting point, since this
involves the breaking of strong covalent bonds within the
layers.
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Part II Microscopic World I I
Some physical properties of diamond and graphite are
summarized in Table 9.1.
Property
Appearance
Hardness
Electricalconductivity
Melting point (°C)°
Graphite
black solid
soft, brittle
conductor (conducts in the directionparallel to hexagonal planes)
()
3730
Diamond
colourless
non-conductor
3550
9.1
hardest naturalsubstance on the Earth
Table 9.1 Some propertiesof diamond and graphite.
Structure and properties of silicon(IV) oxide
Both elements and compounds may form giant covalent
structures. An example of a compound having a giant covalent
structure is silicon(IV) oxide (or silicon dioxide) SiO2.
In the structure of silicon(IV) oxide, each silicon atom is
covalently bonded to four oxygen atoms. Each oxygen atom is
bonded to two silicon atoms (Figure 9.7). Silicon and oxygen
atoms are joined together by covalent bonds throughout the
whole structure.
(IV)
(IV) ( ) SiO2
(IV)
( 9.7)
Figure 9.7 The giant covalent structure of silicon(IV) oxide. Note that this represents only a verysmall part of the lattice, which extends in all directions.
(IV)
silicon atom
oxygen atom
85
Chapter 9 Structures and properties of substances
There are no discrete SiO2 molecules in silicon(IV) oxide.
Thus SiO2 is only an empirical formula, not a molecular formula.
This formula shows that the simplest whole number ratio of Si :
O atoms in the compound is 1 : 2.
Because of its structure, silicon(IV) oxide has a very high
melting point (1610°C) and boiling point. Also, it does not
conduct electricity whether it is in the solid state or molten.
(IV) SiO2
SiO2
Si O
1 2
(IV) (1610°C)
✘ Silicon(IV) oxide, with a formula SiO2, has a simple molecularstructure.
✔ Silicon(IV) oxide has a giant covalent structure. The formulaSiO2 only represents the composition of the elements in thelattice.
Check your concept
✘
✔
Properties of giant covalent structures
1. Giant covalent structures are all solids with very high melting
points and boiling points. To break the lattice in melting or
boiling the solid, a lot of heat energy must be supplied.
This is because a great number of strong covalent bonds
must be broken.
2. All (except graphite) are hard.
3. They are insoluble in any solvent.
4. All (except graphite) are non-conductors of electricity.
1.
( )
2. (
)
3.
4. (
)
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Part II Microscopic World I I
Example 9.2Statements about covalent substances
This question consists of two separate statements. Decidewhether each of the two statements is true or false; if bothare true, then decide whether or not the second statement isa correct explanation of the first statement.
‘Covalent substances are all gases, liquids or low-meltingsolids.’
‘Covalent bonds are weak.’
Solution
Both statements are false!
1st statement: This is true only for simple molecularstructures; covalent substances with a giant covalentstructure are solids with very high melting points.
2nd statement: Covalent bonds are strong forces of attraction.
➲ Try Chapter Exercise Q17
9.2
➲ 17
(a) ( (IV) )
(i) (ii)
(b)
9.4
(a) Is tridymite (a form of silicon(IV) oxide) soluble in
(i) water (ii) heptane?
(b) Does molten tridymite conduct electricity?
Class practice 9.4
9.6 Giant metallic structures 9.6
Giant metallic structures
A metal consists of atoms packed closely together. Take sodium
as an example. A sodium atom has the electronic arrangement
2,8,1. This single outermost shell electron is far away from the
nucleus, so it can escape easily to leave a positive sodium ion.
The outermost shell electrons of all sodium atoms move freely
and randomly in the sodium metal. These are delocalized
electrons, since each electron no longer holds onto the nucleus
of its original atom. What is formed is a giant metallic
structure — a giant lattice of metal ions surrounded by a ‘sea’
of freely moving electrons.
2, 8, 1
87
Chapter 9 Structures and properties of substances
Metallic bond
Metal atoms are joined to one another in a giant metallic
structure by metallic bonds, which result from the
attraction between a ‘sea’ of delocalized electrons and metal
ions.
2,8,8,2
(a)
(b)
9.5
Calcium has the electronic arrangement 2,8,8,2.
(a) How many outermost shell electrons does a calcium atom
have?
(b) How many delocalized electrons does each calcium atom
in the metal contribute?
Class practice 9.5
Properties of metals explained by structure and bonding
We can explain the common physical properties of metals by
their special structure and bonding.
• Metals are good conductors of electricity. In a piece of metal,
the delocalized electrons move freely in all directions.
However, when both ends of the metal piece are connected
to a battery, the delocalized electrons move towards the
positive pole of the battery, leaving the metal. At the same
time, an equal number of electrons move into the other end
of the metal from the negative pole. An electrical circuit is
complete.
• Metals are good conductors of heat. When one end of a piece
of metal is heated, the delocalized electrons there get more
energy. They move faster, colliding with the neighbouring
electrons. Heat is transferred in the collisions.
•
•
88
Part II Microscopic World I I
• Most metals are solids with high melting points and boiling
points. A lot of energy is required to break the strong
metallic bonds in a giant metallic structure.
• Most metals have high densities. This can be explained by
the close packing of metal atoms in a regular arrangement.
• Metals are malleable (can be rolled into sheets and other
shapes) and ductile (can be pulled out into wires). The
atoms in a metal are packed in layers. When we apply force
to a piece of metal, the layers of atoms can slip over one
another. As a result, atoms settle into new positions and the
piece of metal takes up a new shape. The metal piece does
not break. This is because the non-directional metallic bonds
continue to hold the metal atoms together.
•
(
)
•
• ( )
( )
9.7 Comparison of structures and propertiesof substances
9.7
The bonding, structures and properties of substances with
simple molecular, giant ionic, giant covalent and giant metallic
structures are summarized in Table 9.2.
9.2
Simple molecularstructure
Giant ionicstructure
Giant covalentstructure
Giant metallicstructure
(1) Examples H2, I2, H2O, NH3, CCl4 NaCl, CaO, KOH C (diamond), C (graphite), SiO2
C ( ), C ( ), SiO2
All metals
(2) Structure small discretemolecules e.g. H2
H2
giant lattice of ionse.g. NaCl
NaCl
giant lattice of atomse.g. C (diamond)
C ()
metal ions,surrounded by a‘sea’ of freelymoving electrons
89
Chapter 9 Structures and properties of substances
Simple molecularstructure
Giant ionicstructure
Giant covalentstructure
Giant metallicstructure
(3) Bonds holdingconstituentparticles
strong covalent bondsbind atoms togetherwithin a molecule;separate molecules areattracted by weakintermolecular forces(e.g. van der Waals’forces)
()
ionic bonds linkoppositely chargedions throughout thestructure
covalent bonds linkatoms throughoutthe networkstructure
metallic bonds linkthe metal ions(positivelycharged) and the‘sea’ of electrons(negativelycharged)
()
( )
gases, volatile liquids,or solids of low melting points
solids solids solids (exceptmercury)
( )
low high very high usually high
soft hard usually high usually high
(i) most are insoluble
(ii) generally soluble
(i) most are soluble
(ii) insoluble
(i) insoluble
(ii) insoluble
(i) insoluble
(ii) insoluble
(4) Physical properties
(a) State at roomconditions
(b) M.p. and b.p.
(d) Solubility in
(i) water
(ii) non-aqueoussolvents (e.g.heptane)
( )
(c) Hardness of solidform
non-conductorsNote: A few (e.g.sulphuric acid) reactwith water to form asolution whichconducts electricity
( ( )
)
non-conductorswhen solid; goodconductors whenmolten or inaqueous solution
non-conductors(except graphite)
( )
good conductors(e) Conduction ofelectricity
Table 9.2 Comparison of different structures and properties.
90
Part II Microscopic World I I
9.8 Predicting structure from physicalproperties
9.8
The flow chart as shown in Figure 9.8 may help us to predict
structures other than the metallic structure.
9 . 8
Physical properties Structure
Is the substance a gas orliquid at room conditions?
Does the solid have a lowmelting point?
Does the substance conductelectricity when molten andin aqueous solution?
Does the substance have avery high melting point?
Simple molecular structure
Giant ionic structure
Giant covalent structure
yes
yes
yes
yes
no
no
no
Figure 9.8 Predicting the structure of a substance from its physical properties.
Metals have a giant metallic structure. Usually we can tell whether asubstance is a metal from its electrical conductivity andappearance. At room temperature and pressure, mercury is theonly liquid with a giant metallic structure.
Learning tip
91
Chapter 9 Structures and properties of substances
Class practice 9.6 9.6
The following table gives information about some properties of
substances A to D.A D
Answer the following questions, and explain your answers.
(a) Which substance has a giant metallic structure?
(b) Which substance has a giant ionic structure?
(c) Which substance has a simple molecular structure?
(d) Which substance has a giant covalent structure?
(e) Which substance is likely to be soluble in heptane?
(a)
(b)
(c)
(d)
(e)
Substance M.p. (°C)°
Electrical conductivity
solid molten
A 70 poor poor
B 375 poor good
C 98 good good
D 1610 poor poor
9.9 Predicting physical properties frombonding and structure
9.9
Suppose we know what elements make up a given compound.
From the group number of the elements, we can predict the
bonding and structure of the compound. We can then predict
its physical properties. See Example 9.3. 9.3
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Part II Microscopic World I I
Example 9.3Predicting physical properties of compounds from theirbonding
9.3
Predict the (i) formula (ii) structure (iii) physical properties(melting point, boiling point, hardness, solubility behaviourand electrical conductivity) of the compound formed between
(a) potassium and sulphur
(b) nitrogen and fluorine.
Solution
(a) (i) The compound formed between a metal(potassium) and a non-metal (sulphur) is an ioniccompound.
Potassium (Group I) forms K+ ions;
sulphur (Group VI) forms S2– ions.
The formula of the compound is thus K2S.
(ii) It has a giant ionic structure.
(iii) Its physical properties:
(1) A solid with a high melting point and boilingpoint.
(2) Hard.
(3) Soluble in water, insoluble in most non-aqueous solvents.
(4) Non-conductor of electricity when solid;conductor when molten and in aqueoussolution.
(b) (i) The compound formed between non-metals(nitrogen and fluorine) is a molecular compound.
Nitrogen (Group V) contributes 3 electrons forsharing;
fluorine (Group VII) contributes 1 electron forsharing.
(See p.62 for explanations.)
Using the short-cut method to predict itsmolecular formula:
3 1
N F N1F3 NF3
The formula of the compound is thus NF3.
cont'd
(i) (ii) (iii) (
)
(a)
(b)
(a) (i) ( ) ( )
( I ) K+
( VI ) S2–
K2S
(ii)
(iii)
(1)
(2)
(3)
(4)
(b) (i) ( )
( V ) 3
( VII ) 1
( 62 )
3 1
N F N1F3 NF3
NF3
93
Chapter 9 Structures and properties of substances
(ii) It has a simple molecular structure.
(iii) Its physical properties:
(1) A substance with a low melting point andboiling point.
(2) In solid state, the compound is soft.
(3) Insoluble in water, soluble in non-aqueoussolvents.
(4) Non-conductor of electricity no matter solidor liquid.
(ii)
(iii)
(1)
(2)
(3)
(4)
9.3 X Y(a) (b)
(c) (X II YVII )
9.7
Using Example 9.3 as reference, predict the (a) formula, (b)
structure and (c) physical properties of a compound formed
between two elements X and Y. (X belongs to Group II; Y
belongs to Group VII.)
Class practice 9.7
9.10 Applications of substances according totheir structures
9.10
Substances of different properties are used in our daily life for
different purposes. For example, graphite is a covalent
substance with high melting point and boiling point. On the
other hand, it can conduct electricity. Because of these
properties, it is widely used as electrodes in many cases.
94
Part II Microscopic World I I
KK ee yy tt ee rr mm ss
1. delocalized electron 86
Page
3. giant metallic structure 86
4. giant network 82
7. macromolecule 75
8. metallic bond 87
9. molecular structure 75
10. non-directional 80
11. simple molecular structure 75
12. van der Waals’ forces 77
5. giant structure 75
2. giant covalent structure 82
6. intermolecular forces 77
SS uu mm mm aa rr yy9.1 Structure of substances
1. The of a substance is a description of what its constituent particles are, andabout how they are arranged or packed together.
9.2 Simple molecular structures
2. In some substances, atoms within a molecule are bound together by strong covalent bonds andeach molecule is attracted to other neighbouring molecules by weak
.
95
Chapter 9 Structures and properties of substances
9.3 Macromolecules
3. are very large molecules, each containing thousands of atoms. Examplesinclude plastics, proteins and some carbohydrates.
9.4 Giant ionic structures
4. In ionic compounds, crystals consisting of positive and negative ions are held together by strongnon-directional electrostatic attractions. The ions are regularly packed to form a continuous,three-dimensional .
9.5 Giant covalent structures
5. In a few elements and compounds, the non-metal atoms join together by covalent bonds to form agiant network called .
9.6 Giant metallic structures
6. Metal atoms are joined to one another in a by, which result from the attraction between a ‘sea’ ofand metal ions.
9.7 Comparison of structures and properties of substances
7. The structure, bonding and physical properties of simple molecular structure, giant ionicstructure, giant covalent structure and giant metallic structure are summarized in Table 9.2 onp.88.
9.8 Predicting structure from physical properties
8. It is possible to predict the structure of a substance from its properties.
(Refer to the flow chart in Figure 9.8 on p.90.)
9.9 Predicting physical properties from bonding and structure
9. It is possible to predict the physical properties of a substance from its bonding and. (Refer to Example 9.3 on p.92.)
9.10 Applications of substances according to their structures
10. Some specialized new materials have been created on the basis of the findings of research on thestructure, chemical bonding, and other properties of matter.