HKDSE Chemistry Bridging Programe 1B (1)

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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


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


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


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


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


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


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


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


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


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


Example 5.4Statements about oxygen atom


‘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


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


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


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


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


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


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


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


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


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


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


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


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


(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


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


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


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


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


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


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


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


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


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


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


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


( )

(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


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


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


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


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


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


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


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


‘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


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


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


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


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


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


(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


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


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


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


56


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


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


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


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


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


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


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


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


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


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


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


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


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


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


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



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


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


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


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


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


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


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


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


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


‘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


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.

83


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.

84


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


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. (

)

86


Example 9.2Statements about covalent substances


‘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


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


• 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


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


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


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

92


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


(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



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


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.

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HKDSE Chemistry Bridging Programe 1B (1)