HKDSE Chemistry Bridging Programe 1B (1)

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First published July, 2009

A teacher’s book is available for use by teachers.

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 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

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

Chapter 5 Atomic structure

5.1

What is an element?

5.1

In Chapter 1, we have defined that an element is a pure substance which cannot be broken down into anything simpler by chemical methods.

5.2


5.2

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.

5.3


11

2

5.3

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

)

♦ ♦

( (

) )

1

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

Property

Metals

State at room temperature and

solids (except mercury — a liquid)

pressure

Non-metals gases or solids (except bromine — a liquid)

shiny; mostly silvery white in

usually dull in appearance (when

colour (except copper and gold)

solid); in various colours

Melting point and boiling point

usually high

usually low

Hardness and strength

hard and strong

not uniform in hardness and

Appearance

strength

malleable and ductile

Malleability and ductility

brittle i.e. easily broken when hit (when solid); not malleable and not ductile

Density Thermal

conductivity

electrical conductivity

and

usually high

low

good conductors of heat and

poor conductors of heat;

electricity

non-conductors of electricity

Table 5.1 Some typical differences in physical properties of metals and non-metals.

2


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

100°C

non-metal carbon (in the form of graphite). It is a good electrical conductor, shiny, and has a very high melting point

(

)

(3730°C). (3730°C)

Example 5.1 Metal or non-metal?

5.1

A reddish brown solid element X conducts electricity well. Is X a metal or non-metal? Give reasons.

X X

Solution X is a metal. All metals conduct electricity while non-metals cannot. (An exception to this rule is the non-metal graphite (a form of carbon), but its colour is black, not reddish brown.)

X ( (

) )

5.1

Class practice 5.1 What characteristics do the two elements mercury and bromine have in common?

Check your concept

✘ ✔

All elements can be classified as metals or non-metals.

✘

Many, but not all elements, can be classified as metals or nonmetals. A few elements have properties in between those of metals and non-metals — they are classified as semi-metals.

✔

3


I

The in-between elements — the semi-metals A few elements, called semi-metals (or metalloids), have properties in between those of metals and non-metals. Examples of semi-metals include boron and silicon.

( )

(

)

Most semi-metals have important uses in industry. An example is silicon, a semi-conductor. It is widely used in making transistors and silicon chips.

Example 5.2 Statements about graphite

5.2

This question consists of two separate statements. Decide whether each of the two statements is true or false; if both are true, then decide whether or not the second statement is a 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) is a non-metal. The second statement is true.

( )

5.2

Class practice 5.2 1.

2.

Would you classify the following elements/compounds as a metal or non-metal? Why?

1.

(a) Water

(a)

(b) Graphite

(b)

(c) Mercury

(c)

Decide which is the odd one in each of the following groups of elements. Give reason(s) for your choice in each case.

2. ( )

(a) Iron, copper, mercury, silver

4

(b) Magnesium, sulphur, lead, tin

(a)

(c) Iodine, oxygen, nitrogen, argon

(b)

(d) Phosphorus, bromine, helium, carbon (in the form of graphite)

(c) (d)

(

)


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

5.1

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.

)

6 V battery 6V carbon (graphite) rods ( )

light bulb

crucible solid piece under test

crocodile clip

solid powder (or liquid) under test (

(b)

(a) Figure 5.1 Testing electrical conductivity of substances (a) in form of solid piece and (b) in form of solid powder or liquid.

5.4


)

(a) (b)

5.4

It is useful to give each element a chemical symbol. Chemical symbols of some common metals, non-metals and semi-metals

5.2

are given in Table 5.2.

5


I

Metal Element [Latin name] Aluminium (

)

Barium (

) Beryllium ( ) Calcium ( ) Chromium ( ) Cobalt ( ) Copper ( ) [Cuprum] Gold ( ) [Aurum] Iron ( ) [Ferrum] Lead ( ) [Plumbum] Lithium ( ) Magnesium ( ) Manganese ( ) Mercury ( ) [Hydrargyrum] Nickel ( ) Platinum ( ) Potassium ( ) [Kalium] Silver ( ) [Argentum] Sodium ( ) [Natrium] Tin ( ) [Stannum] Zinc ( )

Non-metal

Semi-metal Chemical Symbol

Chemical Symbol

Element

Al

Boron (

Ba

Silicon (

) )(

)

Chemical Symbol

Element

B

Argon (

)

Si

Bromine (

Be

Carbon (

Ca

Chlorine (

Ar

)

Br

)

C

) Fluorine ( ) Helium ( ) Hydrogen ( ) Iodine ( ) Neon ( ) Nitrogen ( ) Oxygen ( ) Phosphorus ( ) Sulphur ( )

Cr Co Cu Au Fe Pb Li Mg Mn

Cl F He H I Ne N O P S

Hg Ni Pt K Ag Na Sn Zn

Table 5.2 Chemical symbols of some common elements (classified into metals, semi-metals and non-metals). ( )

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.

5.3

Class practice 5.3 5.2

Referring to Table 5.2,

6

(a) Give the chemical symbols for (i) magnesium, (ii) silver and (iii) sodium.

(a)

(i)

(ii)

(iii)

(b) Give the chemical symbols for the noble gases (i) argon, (ii) helium and (iii) neon.

(b)

(i)

(ii)

(iii)

(c) Write the names of (i) F, (ii) Br and (iii) Hg.

(c)

(i) F

(ii) Br

(iii) Hg


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 –8

–23

10 cm. They have masses of around 10

g.

10 cm

10

–23

–8

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,

2008

corresponding to the 118 different elements.

1

118

118

Symbols for atoms You have learnt chemical symbols of some elements on p.6 —

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

C

the atomic symbol for a carbon atom.

5.4

Class practice 5.4 (a) What is the total number of atomic symbols at present?

(a)

(b) What is the chemical symbol for the element bromine? (c) What is the atomic symbol for a nitrogen atom?

(b)

(d) What does Cu stand for?

(c) (d) Cu 7


5.6

I

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-

5.3

atomic particles.

Sub-atomic particle

Symbol

Mass (in g) (g)

Relative mass

Electric charge (relative to that on a proton)

Position within the atom

–24

1

+1

inside nucleus

–24

1

0

inside nucleus

–1

space outside nucleus

Proton

p

1.6725 10

Neutron

n

1.6748 10

negligible

Electron

–

e

–28

9.109 10

1 ( ) 1837

Table 5.3 Data on the three fundamental sub-atomic particles.

Building up different atoms from protons, neutrons and electrons Different atoms have different numbers of protons, neutrons and electrons.

8

)


The hydrogen atom is the simplest of all atoms. The commonest type of hydrogen atoms consists of 1 proton and 1

1

electron (with no neutron). The next simplest one is helium atom, with 2 protons, 2 neutrons and 2 electrons (Figure 5.2).

(

1

)

2

2

2

(

5.2)

neutron

electron

} proton

Figure 5.2 Diagrammatic representations of a hydrogen atom and a helium atom.

nucleus

helium atom

hydrogen atom

Table 5.4 gives the number of protons, neutrons and

5.4

20

electrons in the 20 simplest atoms. Atom

Number of

Symbol protons

Hydrogen (

)

electrons

H

1

0

1

Helium (

)

He

2

2

2

Lithium (

)

Li

3

4

3

Be

4

5

4

B

5

6

5

C

6

6

6

N

7

7

7

Beryllium ( Boron (

)

)

Carbon (

)

Nitrogen (

)

Oxygen (

)

O

8

8

8

Fluorine (

)

F

9

10

9

Ne

10

10

10

Na

11

12

11

Neon (

)

Sodium (

)

Magnesium (

)

Mg

12

12

12

Aluminium (

)

Al

13

14

13

Si

14

14

14

P

15

16

15

Silicon (

)

Phosphorus ( Table 5.4 Number of protons, neutrons and electrons in the 20 simplest atoms. 20

neutrons

)

Sulphur (

)

S

16

16

16

Chlorine (

)

Cl

17

18

17

Ar

18

22

18

K

19

20

19

Ca

20

20

20

Argon (

)

Potassium ( Calcium (

) )

9


I

Atoms are electrically neutral Although an atom contains electrically charged particles, the atom itself has no overall charge. That is, an atom is electrically neutral. This is because in an atom, the number of protons is equal to the number of electrons. On the other hand, the number of neutrons may not be equal to the number of protons (look at Table 5.4 again).

5.4

5.5

Class practice 5.5 (a) All atoms (except one) are made up of protons, electrons

(a)

and neutrons. Which atom does not contain any neutron at all? (b) A certain atom contains 91 protons. How many electrons

(b)

91

and neutrons does it have? (c) A certain particle has 8 protons, 8 neutrons and 10 electrons. Is it an atom? Why?

5.7


(c)

8

8

10

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

47

number of silver is therefore 47.

47

5.6

Class practice 5.6 Refer to Table 5.4. What would happen if an atom with 12 protons were changed to one with 17 protons?

10

5.4 12

17


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.

( 12

)

11 11 + 12 = 23

Learning tip The electrons in an atom have almost no mass. So the mass of an atom is nearly all due to protons and neutrons. For this reason, the number of protons plus the number of neutrons in an atom is called the mass number.

The atomic number (Z) and mass number (A) of an atom

(Z)

(A)

are usually shown in a full atomic symbol as follows: mass number = number of protons + number of neutrons

A Z

X

=

EXAMPLE

+

Atomic symbol

atomic number = number of protons = number of electrons of a neutral atom

mass number

4

atomic number

2

= =

Example 5.3 Working out the number of protons, electrons and neutrons in an atom Consider

35 17 Cl.

He

Work out the number of protons, electrons

5.3

35 17

Cl

and neutrons in the atom. Solution Atomic number (Z) = 17, so number of protons = 17 (by definition) As an atom is electrically neutral, number of electrons = number of protons = 17 Mass number (A) = 35, ∴ number of neutrons = mass number – number of protons = 35 – 17 = 18

(

∴

(Z) = 17 )

= 17

= = 17 (A) = 35 = – = 35 – 17 = 18 11


I

Example 5.4 Statements about oxygen atom

5.4

This question consists of two separate statements. Decide whether each of the two statements is true or false; if both are true, then decide whether or not the second statement is a correct explanation of the first statement. ‘The atomic number of ‘An

16 8O

16 8O

is 8.’

atom contains 8 neutrons.’

16 8O

8

16 8O

8

Solution Both statements are true, but the second statement does not explain the first one. A correct explanation would be: ‘An 168 O atom contains 8 protons.’

16 8O

5.7


Fill in the following blanks:

1. (a)

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.

8

47 (b) 47

47

A particular atom of an element (atomic number = 13) has a

(c)

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.

2.

( 27 (a) (b)

(c)

(i) (iii)

12

= 13)

(ii)


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

5.3

5.5

hydrogen atoms, as shown in Figure 5.3 and Table 5.5. They all have the same number of protons (same atomic number) and

(

) 1 1

electrons but different numbers of neutrons. Therefore, hydrogen has 3 isotopes:

1 2 1 H, 1 H

and

3 1 H.

H

2 1

H

3 1H

electron proton

neutron

Figure 5.3 The three isotopes of hydrogen.

1 1H

2 1H

Number of

Isotope

Table 5.5 Number of protons, electrons and neutrons in the three isotopes of hydrogen.

3 1H

p

e

–

n

1 1

H

1

1

0

2 1

H

1

1

1

3 1

H

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)

13


I

Element

Hydrogen

Carbon

Oxygen

Table 5.6 Isotopes of some elements in Nature.

Sodium Chlorine

Isotopes

Atomic number

Mass number

1 1

H

1

1

99.984

2 1

H

1

2

0.016

3 1

H

1

3

very small percentage

12 6

C

6

12

98.892

13 6

C

6

13

1.108

14 6

C

6

14

very small percentage

16 8

O

8

16

99.76

17 8

O

8

17

0.04

18 8

O

8

18

0.20

23 11

Na

11

23

100

35 17

Cl

17

35

75.4

37 17

Cl

17

37

24.6

5.8


Refer to Table 5.6. (a) How many natural isotopes does oxygen have?

(a)

(b) Which is the most abundant isotope of oxygen?

(b)

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.

14

% abundance of isotopes in Nature


5.9


5.9

Relative isotopic mass The carbon-12 scale

-12

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 6

6

)

12.000 units (atomic mass unit, a.m.u.). Masses of all other atoms are compared with this reference standard to give their

(

12.000

(

a.m.u.)

relative masses. 12

On the C = 12.000 00 scale, the relative masses of a proton and a neutron are both very close to 1; the relative mass of an

12

C = 12.000 00

electron is nearly 0. Thus the relative isotopic mass of an

1

isotope is roughly equal to its mass number.

0

Learning tip ‘Relative isotopic mass’ and ‘relative atomic mass’ are both relative values; they carry no units.

Relative isotopic mass ≈ mass number

≈

5.9

Class practice 5.9 What is the relative isotopic mass of (a)

37 17 Cl

(b)

35 17 Cl

(c)

4

He

(d)

238

U

(e)

19 K?

(a)

37 17 Cl

(d)

238

U

(b) (e)

35 17 Cl

(c)

4

He

19 K

15


I

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)

The relative atomic mass of an element is the weighted average of the relative isotopic masses of its natural isotopes

( 12

12

on the C = 12.000 00 scale.

C = 12.000 00

)

For example, for an element consisting of three isotopes A, B and C:

A

B

C

= a% MA + b% MB

Relative atomic mass = a% MA + b% MB + c% MC

+ c% MC

where a%, b%, c% = percentage abundance of isotopes A, B and

a%

b%

c% =

MA

MB

MC =

A

B

C

C respectively MA, MB, MC = relative isotopic masses of isotopes A, B and C respectively

16

A

B

C


Example 5.5 Calculating relative atomic mass and percentage abundance 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 of chlorine. (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 two isotopes in natural bromine.

5.5

35

(a)

Cl 75.4%

37

Cl 24.6%

79

(b) 81

Br

Br ( = 79.9

)

Solution (a) By approximation, relative isotopic mass of 35Cl isotope = its mass no. = 35

(a) 35

relative isotopic mass of 37Cl isotope = its mass no. = 37

Cl

=

37

Cl

= weighted average of the relative isotopic masses 75.4 24.6 35 + 37 = 35.5 100 100

= 37 12

=1

(Note: The relative atomic mass of 35.5 is not the relative mass of any one chlorine atom, but the weighted average of all the chlorine atoms present.) 79

Br and

81

Br be y%

=

75.4 100

35 +

∴ y = 55

24.6 100

37 = 35.5

( 35.5

79y + 81(100 – y) 100

7990 = 79y + 8100 – 81y

C = 12.000 00

=

Relative atomic mass of bromine = weighted average of the relative isotopic masses 79.9 =

= 35 37

= average mass of 1 chlorine atom on the 12C = 12.000 00 scale

(b) Let the percentage abundance of and (100 – y)% respectively.

=

Cl

Relative atomic mass of chlorine

=

Cl

35

35.5 (b)

)

79

Br 81 Br y% (100 – y) % =

Thus the percentage abundance of 79Br is 55% and that of 81Br is 45%. 79.9 =

79y + 81(100 – y) 100

7990 = 79y + 8100 – 81y ∴ y = 55 79

Br 55%

81

Br 45%

17


I

5.10


There is only one kind of sodium atoms in nature, i.e.

23 11 Na.

23 11 Na

1.

What is the relative atomic mass of sodium? 2.

Neon in air contains 90% of

20 10 Ne

and 10% of

22 10 Ne.

2.

90% 22 10 Ne

10%

Calculate the relative atomic mass of neon.

20 10

The accurate relative atomic masses of elements are very seldom whole numbers (why?).

(

)

Check your concept

✘ ✔

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.

✔

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

1

2

3

(

4 )

electrons (Table 5.7).

(

The arrangement of electrons in a sodium atom can be

5.4

shown by Figure 5.4. Shell number,

2

Maximum number of electrons (= 2n )

n

Table 5.7 Maximum number of electrons the first four shells can hold.

18

1 2 3 4 . . .

2 8 18 32 . . .

5.7)

Ne


1st shell (innermost shell) ( ) nucleus

2nd shell

electron

Figure 5.4 Arrangement of electrons in a sodium atom.

3rd shell (outermost shell in sodium atom) (

)

Electrons in an atom are arranged in shells. The distribution of electrons in the various shells is called electronic arrangement (or electronic configuration).

(

)

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

1.

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

2.

innermost shell. When a certain shell is ‘full’ (refer to Table

(

5.7 again), any remaining electrons would go into the next

5.7)

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).

) 2, 8, 1

(

5.5) 19


I

Electronic arrangement of sodium atom:

2, 8, 1 Figure 5.5 Showing the electronic arrangement of a sodium atom by numbering.

Number of electrons in:

1st shell

2nd shell

3rd 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

5.6

electron diagram of a sodium atom.

Na

Figure 5.6 The electron diagram of a sodium atom.

5.11

Class practice 5.11 Draw electron diagrams for the following atoms: (a) Helium

(b) Oxygen

(c) Silicon

(d) Calcium

(a)

(b)

(c)

(d)

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).

20

1 – 20

(

5.8)


Element

Symbol

Atomic Number of number electrons

No. of electrons in electron shells 1st

2nd

3rd

Electronic arrangement

4th

Hydrogen

H

1

1

1

1

Helium

He

2

2

2

2

Lithium

Li

3

3

2

1

2, 1

Beryllium

Be

4

4

2

2

2, 2

Boron

B

5

5

2

3

2, 3

Carbon

C

6

6

2

4

2, 4

Nitrogen

N

7

7

2

5

2, 5

Oxygen

O

8

8

2

6

2, 6

Fluorine

F

9

9

2

7

2, 7

Neon

Ne

10

10

2

8

2, 8

Sodium

Na

11

11

2

8

1

2, 8, 1

Magnesium

Mg

12

12

2

8

2

2, 8, 2

Aluminium

Al

13

13

2

8

3

2, 8, 3

Si

14

14

2

8

4

2, 8, 4

Phosphorus

P

15

15

2

8

5

2, 8, 5

Sulphur

S

16

16

2

8

6

2, 8, 6

Chlorine

Cl

17

17

2

8

7

2, 8, 7

Argon

Ar

18

18

2

8

8

2, 8, 8

Potassium

K

19

19

2

8

8

1

2, 8, 8, 1

Calcium

Ca

20

20

2

8

8

2

2, 8, 8, 2

Silicon

(

)

Table 5.8 The electronic arrangements (by numbering) of the elements with atomic numbers 1 – 20. 1 – 20 ( )

5.12

Class practice 5.12 (a) What is the atomic number of chlorine? (See Table 5.8)

(a)

(b) Show the electronic arrangement of a chlorine atom by (i) numbering

(ii) an electron diagram.

5.11 Stability of noble gases related to their electronic arrangements The term ‘noble gases’ is a collective name for Group 0

( 5.8

(b)

) (i)

(ii)

5.11 0

elements, which are very unreactive.

21


I

The exceptional stability of noble gases is related to their electronic arrangements: Helium

(He)

2

(He)

2

Neon

(Ne)

2, 8

(Ne)

2, 8

Argon

(Ar)

2,8, 8

(Ar)

2, 8, 8

Krypton

(Kr)

2,8,18, 8

(Kr)

2, 8, 18, 8

Xenon

(Xe)

2,8,18,18, 8

(Xe)

2, 8, 18, 18, 8

Radon

(Rn)

2,8,18,32,18, 8

(Rn)

2, 8, 18, 32, 18, 8

All noble gases (except helium) have 8 outermost shell

(

)

8

electrons in their atoms. Helium atom has 2 electrons in the only one occupied shell. This suggests that a particle has great

2

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).

•

( 2

Atoms of elements other than noble gases are usually not stable. They will become stable if they attain an octet or a duplet.

22

8

)


Key terms Page 1. atom

7

2. atomic number

10

3. chemical symbol

5

4. duplet

22

5. electron diagram

20

6. electron shell

(

)

18

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

23


I

Summary 5.1

What is an element?

1.

An chemical methods.

5.2


2.

Elements can be classified based on solids, liquids or gases. metal) are the only two liquid elements.

5.3


3.

Elements can be classified into

4.

All metals conduct . All non-metals (except carbon in the form of graphite) do not conduct . To tell whether an element is a metal or non-metal, a simple but effective way is to test whether it conducts .

is a pure substance which cannot be broken down into anything simpler by

states, that is, whether the elements are (a silvery metal) and (a dark red non-

,

and

.

(Refer to Table 5.1 on p.2 for some typical differences in physical properties between metals and non-metals.) 5.4


5.

Chemists use chemical to represent elements. Chemical symbols of most elements come from their English names. (Refer to Table 5.2 on p.6 for chemical symbols of some common metals, non-metals and semimetals.)

5.5

Atoms

6.

An element.

7.

An 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.

5.6

Structure of atoms

8.

(a) An atom consists of three types of sub-atomic particles — and . (b)

Sub-atomic particle

,

Relative mass

Relative charge

Proton (p)

1

+1

Neutron (n)

1

0

Electron (e–)

24

is the smallest part of an element which has the chemical properties of that

negligible (

1 ) 1837

–1


(c) An atom has an extremely small centre called are in the nucleus.

. The protons and neutrons

(d) Electrons move around the nucleus in (e) An atom is electrically 5.7

. .


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 the atom

11.

Full atomic symbol mass number = number of protons + number of neutrons

A Z

X

Atomic symbol

EXAMPLE

mass number

4

atomic number

2

He

atomic number = number of protons = number of electrons of a neutral atom

5.8

Isotopes

12.

are different atoms of the same element, with the same number of protons (and electrons) but different numbers of neutrons. Different isotopes of the same element have the same chemical properties but slightly different physical properties.

5.9


13.

≈ mass number

14.

of an element = weighted average of the relative isotopic masses 12 of its natural isotopes on the C = 12.000 00 scale.

5.10 Arrangement of electrons 15.

The of an atom is the distribution of electrons in the various shells 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 one occupied shell) as in helium. 25


6.1

I


6.1

Grouping elements 92

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

A.

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

B.

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. C. Action of sodium sulphite solution on aqueous chlorine 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.

26

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

1869 63

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 (

6.1)

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.

GROUPS

atomic number

relative atomic mass

electronic arrangement

Noble gases

Halogens

Alkaline earth metals

Alkali metals

PERIODS

Transition elements

Keys:

metal

semi-metal

non-metal

gas

liquid

solid

main groups

Table 6.1 Part of the modern Periodic Table. (A complete Periodic Table is shown on the inside front cover.) ( )

27


I

The elements are arranged in periods and groups of the Periodic Table.

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.

Period number = number of occupied electron shells

=

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).

I

II

III

IV

We should note that Group I elements have one outermost

I

shell electron, Group VII elements have seven outermost shell

VII

V

VI

VII

0

electrons, and so on.

Group number = number of electrons in the outermost shell

=

Learning tip There are exceptions to this rule: (1) Hydrogen does not belong to any group. (2) For Group 0 elements, helium has two electrons in the outermost shell, while all the others have eight.

28

(1) (2)

0

2 8


Figure 6.1 illustrates the above two rules for period

6.1

number and group number. Some of the groups have special names: Group I

:

Alkali metals

I

Group II

:

Alkaline earth metals

II

Group VII

:

Halogens

VII

Group 0

:

Noble gases

0

Electronic arrangement of a chlorine atom:

2, 8, 7 Figure 6.1 The relation among electronic arrangement, period number and group number.

no. of occupied electron shells = 3 = period no. = 3

no. of electrons in the outermost shell = 7 = group no. (VII) =7=

(VII)

=

Example 6.1 Identifying an unknown element based on its atomic number

6.1 X

Element X has an atomic number of 15.

(a)

15 X

(a) Deduce the electronic arrangement of an atom of X.

(b) X

(b) In which (i) group (ii) period of the Periodic Table should X be placed?

(c)

X

(a) (b) (c)

2,8,5 V (i)

(c)

(i)

(ii)

Is X a metal or a non-metal?

Solution (a) 2,8,5 (c) Non-metal

(b)

(i) Group V

(ii) Period 3

(ii)

29


I

6.1

Class practice 6.1 Element W has the electronic arrangement of 2, 8, 18, 32, 18, 8,

W

2. (a) To which period and group of the Periodic Table does W

2, 8, 18, 32, 18,

8, 2 (a) W

belong? What is the special name of the group?

W

(b) By referring to the Periodic Table, name element W. (c) Predict whether W can conduct electricity. Give your

(b)

reason.

W

(c)

W

The elements in between Group II and Group III are called

II

the transition elements (or transition metals). Many common

III (

)

metals such as iron (Fe) and copper (Cu) are transition

(Fe)

(Cu)

elements.

6.3


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

more metallic

Group

I

II

III

IV

V

VI

VII

0

2

Li

Be

B

C

N

O

F

Ne

3

Na

Mg

Al

Si

P

S

Cl

Ar

reactive non-metals

noble gases

Period

Figure 6.2 Elements change from metals to non-metals across Period 2 and Period 3 of the Periodic Table.

reactive metals

less reactive metals

semi-metals

more non-metallic

30

less reactive non-metals

)


6.2

Class practice 6.2 Element X belongs to Group II of the Periodic Table.

X

Element Y is a Group 0 element.

Y

0

Element Z is a Group IV element.

Z

IV

Try to classify X, Y and Z as a metal or a non-metal.

X

II

Y

Z

Electronic arrangement and chemical properties Electronic arrangements of some elements in the Periodic Table are given below:

Group I Period 2 Period 3 Period 4 Period 5 Period 6

Li Na K Rb Cs

I

2, 1 2,8, 1 2,8,8, 1 2,8,18,8, 1 2,8,18,18,8, 1

Group VII F Cl Br I At

VII

2, 7 2,8, 7 2,8,18, 7 2,8,18,18, 7 2,8,18,32,18, 7

Let us take Group I as an example. All Group I elements

Group 0 Ne Ar Kr Xe Rn

I

0

2, 8 2,8, 8 2,8,18, 8 2,8,18,18, 8 2,8,18,32,18, 8

I

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.

Check your concept

✘

Elements in the same group have the same chemical properties.

✔

Elements in the same group have similar chemical properties.

✘ ✔

31


I

Example 6.2 Deciding which elements show similar chemical properties

6.2

Which of the following pairs of atoms would have similar chemical properties? Explain your answer. A.

6X

and

15 Y

B.

4X

and

20 Y

C.

9X

and

16 Y

D.

7X

and

17 Y

A.

6X

15 Y

B.

4X

20 Y

C.

9X

16 Y

D.

7X

17 Y

Solution The subscripts stand for atomic numbers. Electronic arrangements of the atoms: A.

6X

(2, 4) and

B.

4 X (2, 2) and

20 Y (2, 8, 8, 2)

C.

9X

(2, 7) and

16 Y

(2, 8, 6)

D.

7X

(2, 5) and

17 Y

(2, 8, 7)

In B,

4

X and

20 Y

15 Y

(2, 8, 5)

have the same number of outermost shell

A.

6X

(2, 4)

15 Y

(2, 8, 5)

B.

4X

(2, 2)

20 Y

(2, 8, 8, 2)

C.

9X

(2, 7)

16 Y

(2, 8, 6)

D.

7X

(2, 5)

17 Y

(2, 8, 7)

electrons, so they should have similar chemical properties.

4

The atomic number of an element P is 20.

P

(a) What is the electronic arrangement of a P atom?

(a)

(b) Would P conduct electricity? Why?

(b) P

(c) Which of the following atoms would have chemical properties similar to P?

(c)

6.4

8Q

(ii)

12 R


20 Y

6.3

Class practice 6.3

(i)

X

P

P (i)

6.4

20

8Q

(ii)

12 R

—

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.

32

I

VII

0


Group I: The alkali metals

I

Figure 6.3 shows the elements in Group I.

6.3

Figure 6.3 Group I elements (the alkali metals). I ( )

lithium

Li

sodium

Na

potassium

K

rubidium

Rb

caesium

Cs

francium

Fr

}

I

silvery solids

Similarities of Group I elements

I

1.

All are soft metals.

1.

2.

All are silvery solids (when freshly cut).

2.

3.

All are reactive.

3.

4.

All have similar chemical properties.

4.

5.

All react with water, giving off hydrogen to form an alkaline

5.

solution. That is why we call them alkali metals.

Difference in reactivity of Group I elements

I

Although all alkali metals are reactive, they differ in reactivities.

Reactivity of Group I elements increases down the group.

I

In fact, this rule also applies to Group II elements (the

II

alkaline earth metals).

(

Group VII: The halogens

VII

Figure 6.4 shows the elements in Group VII.

6.4

)

VII

33


I

fluorine

F

(pale yellow gas)

chlorine

Cl

(greenish yellow gas)

bromine

Br

(dark red liquid)

iodine Figure 6.4 Group VII elements (the halogens). VII ( )

astatine

I

(black solid)

At

(black solid)

Similarities of Group VII elements

VII

1.

All are poisonous non-metals.

1.

2.

All are reactive.

2.

3.

All have similar chemical properties. For example, their

3.

aqueous solutions are turned colourless by sodium sulphite solution (p.26).

(

Difference in reactivity of Group VII elements

VII

Group 0: The noble gases

0

Figure 6.5 shows the elements in Group 0.

6.5 helium

He

neon

Ne

argon

Ar

krypton

Kr

xenon

Xe

radon

Rn

}

colourless gases

Similarities of Group 0 elements

34

0

1.

All are colourless gases.

1.

2.

All are very stable. They have little or no reaction with

2.

other elements.

)

VII

Reactivity of Group VII elements decreases down the group.

Figure 6.5 Elements in Group 0 (the noble gases). 0 ( )

26

( 0

)


6.5


6.5

We can predict the chemical properties of an element from its position in the Periodic Table.

6.4


Can the chemical properties of an unfamiliar element be

1.

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?

2.

(

A.

A metal more reactive than magnesium

18, 32, 18, 7) 18, 8, 2)

B. C.

A metal less reactive than magnesium A non-metal more reactive than chlorine

A.

D. A non-metal less reactive than chlorine

(

2, 8, 2, 8,

B. C. D.

Key terms Page 1. group

28

2. main group

28

3. period

28

4. Periodic Table of Elements

27

5. reactivity

33

6. transition element

30

7. transition metal

30

35


I

Summary 6.1


1.

Some elements show

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.

chemical properties.

(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


4.

Across a period from left to right, there is a change from metals, to to .

5.

Elements within the same group of the Periodic Table have the outermost shell electrons in their atoms, therefore they have properties. However, there is a gradual change in reactivity down a group.

6.4


6.

Group I elements are called the

.

Reactivity of Group I and II elements

36

number of chemical

.

Group II elements are called the

7.

and finally

down the group.

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


9.

Chemical properties of an unfamiliar element can be predicted from its Periodic Table.

in the

Chapter 7 Chemical bonding: ionic bonding

7.1


7.1

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?

?

An ion is an atom or a group of atoms having an overall electric charge.

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,

(II)

copper(II) ion, chloride ion and bromide ion. Examples of

(II)

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.

— )

( 7.1 37


I

electron flow

anode

cathode

Figure 7.1 Movement of cations and anions in electrolysis.

electrolyte anion

-

+

-

cation

+

7.1

Class practice 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.

7.2


(a)

/

(b)

/

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

(II)

form colourless ions (not listed in Table 7.1).

(

) ) (

7.1

)

Name

Table 7.1 The colours of some ions in aqueous solution.

(a) Copper(II) ion

(II)

blue or green

(b) Iron(II) ion

(II)

pale green

(c) Iron(III) ion

(III)

yellow or brown

(d) Cobalt(II) ion

(II)

pink

(e) Nickel(II) ion

(II)

green

(f) Chromium(III) ion

(III)

green

(g) Chromate ion

yellow

(h) Dichromate ion

orange

(i) Manganese(II) ion (j) Permanganate ion

38

Colour

(II)

very pale pink purple


7.2

Class practice 7.2 Predict the colour (if any) of each of the following solutions:

(

(a) Magnesium nitrate solution

(a)

(b) Sodium permanganate solution

(b)

(c) Ammonium chromate solution

(c)

(d) Iron(II) sulphate solution

(d)

)

(II)

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. Some examples are given in Table 7.2.

Gemstone

7.2

Ion responsible for colour

Colour

Amethyst

purple

manganese(III) ion

Emerald

green

chromium(III) ion

(III)

Jade

green

chromium(III) ion

(III)

Peridot Topaz Turquoise

light green yellow bluish green

iron(II) ion iron(III) ion copper(II) ion

(III)

(II) (III) (II)

Table 7.2 Coloured ions in some gemstones.

Migration of ions We can observe the migration (movement) of coloured ions

7.2

during electrolysis, using the set-up as shown in Figure 7.2.

39


I

20 V d. c. supply 20 V carbon cathode

carbon anode

dilute hydrochloric acid dilute hydrochloric acid this region slowly becomes orange due to the migration of negative dichromate ions towards the positive anode

this region slowly becomes blue due to the migration of positive copper(II) ions towards the negative cathode (II)

a gel containing copper(II) ions and dichromate ions (II)

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

7.3

ions under the influence of an electric field is shown in Figure 7.3. small potassium permanganate crystal small potassium permanganate crystal

filter paper moistened with sodium sulphate solution

purple spot

microscope slide filter paper moistened with sodium sulphate solution

purple spot anode anode

cathode

cathode microscope slide 20 V d.c. supply 20 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 a microscope slide). ( )

7.3

Class practice 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? 40

7.3 (a) (b) (c)

(III)


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 7.1 Understanding how an ion is formed

7.1

Explain, in terms of electronic arrangement and number of protons and electrons, the formation of (a) a lithium ion

(a)

(b)

(b) an oxide ion.

Solution

(a)

Li 2,1

(a) Consider a lithium atom, Li.

=3

Electronic arrangement: 2,1 Number of protons = 3; number of electrons = 3 Charge of the atom = (+1) 3 + (–1) 3 = 0 (i.e. the atom carries no charge)

= (+1) 3 + (–1) 3 =0 ( —2(

)

)

=3–1=2

Number of electrons = 3 – 1 = 2 Charge of the ion = (+1) 3 + (–1) 2 = +1 (written as 1+ or +)

= (+1) 3 + (–1) 2 = +1 ( 1+ +)

The resulting positive ion is called lithium ion, + represented by Li .

+

Li (

(Note that ‘1’ is usually dropped out in writing the + 1+ charge on an ion. Thus we write Li instead of Li .) Electronic arrangement: 2,6 Number of protons = 8; number of electrons = 8 Charge of the atom = (+1) 8 + (–1) 8 = 0 (i.e. the atom carries no charge)

)

(

To get the electronic arrangement of the nearest noble gas (helium) — 2 (which is a duplet), one electron has to be removed. An ion is formed.

(b) Consider an oxygen atom, O.

=3

Li

1+

1 )

Li

(b)

+

O 2,6 =8

=8

= (+1) 8 + (–1) 8 =0

To get the electronic arrangement of the nearest noble gas neon — 2,8 (which is an octet), two electrons have to be gained. An ion is formed.

)

( ( — 2,8 (

Number of electrons = 8 + 2 = 10 Charge of the ion = (+1) 8 + (–1) 10 = –2 (written as 2–)

) )

= 8 + 2 = 10

The resulting negative ion is called oxide ion (not 2– oxygen ion), represented by O .

= (+1) 8 + (–1) 10 = –2 ( 2–) 2–

O

41


I

Polyatomic ions are formed from a group of atoms. However, their formation is not discussed here.

7.4


2.

Write down the electronic arrangements of

1.

(a) aluminium atom and aluminium ion

(a)

(b) chlorine atom and chloride ion

(b) 2.

Put down the charge of each ion in Question 1.

1

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).

7.5

Class practice 7.5 State which of the following formulae stand for simple ions and polyatomic ions respectively. +

(a) H2

(b) H

(c) H

(e) CCl4

(f)

NH3

(g) H

(j)

Mn

(i)

OH

–

–

+

(d) NH4

–

(h) NH2

2+

(a) H2 + (d) NH4 –

(g) H 2+ (j) Mn

Names and formulae of common ions Table 7.3 gives the names of some common ions with their formulae.

42

7.3

+

(b) H (e) CCl4 –

(h) NH2

(c) H (f) NH3 (i)

OH

–


Cations Cations Charge Charge

Formula Formula Na

+

K

+

Cu

+

Ag

+

+

H

1+

Name Name

Charge

+

NH4

Formula –

sodium ion

+

Hg

Anions

H

hydride ion

–

potassium ion

Cl

copper(I) ion (I) silver ion

Br

mercury(I) ion (I) hydrogen ion

OH

ammonium ion

chloride ion

–

bromide ion

–

I

iodide ion –

hydroxide ion

–

NO3

1–

nitrate ion

–

NO2

nitrite ion –

HCO3

–

HSO4 –

CN

–

–

ClO3 ClO 2+

2+

Ca Ba

2+ 2+

Pb

2+

Fe

2+

2+

Co

2+

Ni

2+

Mn

2+

Cu

2+

Zn

2+

Hg Al

3+

3+

3+

Fe

3+

Cr

O

calcium ion

S

barium ion

SO4

lead(II) ion (II) iron(II) ion (II) cobalt(II) ion (II) nickel(II) ion (II) manganese(II) ion (II) copper(II) ion (II) zinc ion

SO3

–

permanganate ion

hypochlorite ion oxide ion

2–

sulphide ion 2–

sulphate ion

2–

sulphite ion

2–

SiO3

2–

hydrogensulphate ion

chlorate ion

2–

magnesium ion

hydrogencarbonate ion

cyanide ion

MnO4

Mg

Name

2–

CO3

silicate ion carbonate ion

2–

CrO4

2–

Cr2O7

chromate ion dichromate ion

mercury(II) ion (II) 3–

aluminium ion iron(III) ion (III) chromium(III) ion (III)

N

3–

nitride ion

3–

P

phosphide ion 3–

PO4

phosphate ion

Table 7.3 The names and formulae of some common ions. 43


I

Refer to Table 7.3. You should pay special attention to the following points:

7.3 1.

+

All simple metal ions (e.g. Na , Mg ) are cations.

2.

All simple non-metal ions (except H ) and most polyatomic

+

–

2.

(H

+

ions (e.g. OH , HCO3 ) are anions (except NH4 ).

(

There is only one common polyatomic cation — NH4 .

4.

Polyatomic ions usually consist of non-metals only (e.g. 2–

+

2–

metal (e.g. MnO4 , CrO4 , Cr2O7 ). 5.

)

4. (

When a metal forms only one cation, the ion has the same

NO3

MnO4

+

Transition metals can form more than one simple cation with

–

CO3

2–

2–

SO4 ) (

–

(Na ).

2–

2–

CrO4

Cr2O7 )

5.

different charges. To name each ion, a Roman numeral of the metal. For example, iron metal (Fe) can form iron(II)

+

(Na)

indicating the charge is written in brackets after the name 2+

–

HCO3 )

+

name as the metal, e.g. sodium metal (Na) forms sodium ion

6.

) –

— NH4

2–

2–

2+

Mg )

3.

NO3 , CO3 , SO4 ), but some consist of a metal and a non–

+

+

OH

(NH4

+

3.

–

Na

2+

1.

–

(

(Na )

6.

3+

ion Fe and iron(III) ion Fe . 7.

Simple anions have names ending in -ide, e.g. an oxygen 2–

atom (O) forms an oxide ion (O ); a sulphur atom (S) 2–

forms a sulphide ion (S ). 8.

(II)

The polyatomic anion with more oxygen is named as -ate, and that with less oxygen as -ite, e.g. SO4 2–

–

2–

sulphate ion,

(III)

7. CO 3

–

SO3 sulphite ion; NO3 nitrate ion, NO2 nitrite ion. 9.

2–

Ions with 4+ or 4– charges are uncommon. They are not listed in the table. SO 4 SO 3 NO3

2–

–

–

NO2 8.

44

4+

4–

2–


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

3

1

)

lose these electrons, forming a cation (positively charged). For 2+

example, a Mg atom (2,8,2) forms a Mg

ion (2,8). See Figure

(

7.4a.

Mg

)

2+

Mg

(2,8)

(2,8,2)

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

(

2–

atom (2,6) forms an O ion (2,8). See Figure 7.4b.

)

O

(2,8)

Mg

magnesium atom

Figure 7.4 Formation of ions.

2–

–

gains 2e – 2e

Mg

O

magnesium ion

oxygen atom

(a) Formation of a magnesium ion.

O

7.4b

2+

–

loses 2e – 2e

(2,6)

2–

O

oxide ion

(b) Formation of an oxide ion.

All metals form ions: they usually form cations. Some non-

—

metals form ions — most of these are anions.

—

Relation between ionic charge and group number of an element I

II

III

Metals in Groups I, II and III, the number of positive charges on an ion is equal to its group number. V

VI

VII

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, 2–

( (8 – 6)

VI

)

2

2–

O

forming an O ion. 45


I


7.6

Some elements are shown in the incomplete Periodic Table

1.

below. Write the formulae of the corresponding ions.

Group

I

II

2

Li

Be

3

Na

Mg

4

K

Ca

Period

2.

III

IV

V

VI

VII

N

O

F

S

Cl

Al

The atomic numbers of strontium and astatine are 38 and 85 respectively. Write the formula of (a) strontium ion (b) astatide ion.

Br

2.

38 (a) (

)

numbers.)

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.

This question consists of two separate statements. Decide whether each of the two statements is true or false; if both are true, then decide whether or not the second statement is a correct explanation of the first statement. ‘A neon atom and an oxide ion have similar chemical properties.’ ‘A neon atom and an oxide ion have the same electronic arrangement.’ Solution The first statement is false, while the second statement is true. In fact, a neon atom and an oxide ion behave differently because they have different numbers of protons.

46

85

(b)

(Refer to the Periodic Table for atomic symbols and group

Example 7.2 Statements about atom and ion

0

7.2


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

1.

2.

Covalent bond (to be discussed in Chapter 8)

2.

(

)

3.

Metallic bond (to be discussed in Chapter 9)

3.

(

)

7.6


(

)

7.6

Ionic bond Formation of ionic bond between sodium and chlorine

2,8,1

A sodium atom Na has the electronic arrangement 2,8,1. It can + lose one electron to get the stable octet 2,8, forming a Na ion. On the other hand, a chlorine atom Cl has the electronic arrangement 2,8,7. It can gain one electron to get the stable octet – 2,8,8, forming a Cl ion.

2,8 + Na

Na

2,8,7

2,8,8 Cl – Cl

Thus when a sodium atom and a chlorine atom react, the sodium atom loses one electron to the chlorine atom. As a result of this transfer of electron, two ions are formed. See Figure 7.5. 7.5

electron

Na

+

Cl

Na

Cl

transfer

sodium atom (Na) (loses one electron)

chlorine atom (Cl) (gains one electron)

(both unstable, therefore reactive)

+

–

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 the formation of sodium chloride, NaCl. / NaCl 47


I

In the electron ‘dot/cross’ diagrams (or simply electron

(

diagrams) given here, ions are put inside square brackets with

)

the charge written at the top right-hand corner.

Ionic bond is the strong non-directional electrostatic force of attraction between oppositely charged ions. An ionic bond can be formed by the transfer of one or more electrons from one atom (or group of atoms) to another.

(

)

In the above reaction between sodium and chlorine, only the outermost shell electrons are involved. This is true for most chemical reactions. So for electron diagrams in the rest of the book, only the outermost shell will be drawn. Thus Figure 7.5 can be simplified as:

7.5

Na

+

–

+

electron Cl

Na

Cl

2,8

2,8,8

transfer 2,8,1

2,8,7

or even more simply,

Na

+

Cl

Na +

Formation of ionic bond between magnesium and fluorine In the reaction between magnesium and fluorine, a magnesium atom loses 2 electrons, while a fluorine atom gains 1 electron. Therefore, each magnesium atom must combine with two fluorine atoms.

48

– Cl


–

electron F

+

+

Mg

F

2+

–

F

Mg

F

transfer fluorine atom

magnesium atom

fluorine atom

fluoride ion

magnesium ion

fluoride ion

2,8,2

2,7

2,8

2,8

2,8

2,7

(unstable atoms)

(stable ions)

7.7

Class practice 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.

7.7

(a)

Structures of solid ionic compounds +

(b)

7.7

–

+

In sodium chloride, cations (Na ) and anions (Cl ) are attracted

(Na ) –

(Cl )

together by ionic bonds. They are packed regularly, so that each ion is surrounded by six ions of the opposite charge (Figure

(

7.6).

7.6)

This packing continues until a continuous, threedimensional structure called giant ionic structure is formed.

– –

+ +

–

centre of Cl ion

+ –

+

+

–

–

–

centre of Na ion

–

chloride ion

+ – + + –

+ +

+

sodium ion –

–

Sodium chloride crystals +

–

Figure 7.6 Sodium chloride has a giant ionic structure. It consists of Na and Cl ions held together by ionic bonds. + – Na Cl

49


I

Sodium chloride consists of ions, so it is called an ionic compound. Magnesium fluoride is another ionic compound.

An ionic compound (or ionic substance) is a compound

(

)

which consists of ions.

7.8


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

(Na)

atom (Na) loses an electron and becomes an ion, it has a +

positive charge. The symbol for sodium ion is Na . On the

Na

other hand, the symbol for a chlorine atom is Cl. When it

+

Cl

–

–

accepts an electron, it becomes a chloride ion (Cl ). The overall

(Cl )

charge of the sodium chloride compound should be zero because the positive charge on the sodium ion balances the

7.7

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. Charge: +1 : +1

–1

–

+

Cl

Na

Na

+

–

Cl

Figure 7.7 The overall charge of sodium chloride is zero.

50


Example 7.3 Writing the formulae of some ionic compounds

7.3

Give the formulae of the following ionic compounds. (a) potassium oxide

(b) magnesium nitrate

(c)

sodium hydroxide

(d) calcium hydroxide

(e)

iron(III) sulphate

(a) (c) (e)

(b) (d) (III)

Solution (a)

(a) Potassium oxide +

K

2–

K ion carries 1 positive charge; O ion carries 2 negative charges. To have electrical neutrality, the ratio + 2– of K ions: O ions must be 2 : 1.

+

+

K 2 1

Thus the ionic formula of potassium oxide is as shown below:

+

+

2–

(K )

2

O

O

(K )

one oxide ion

2–

2–

O

2

+

2–

O

this number written after the brackets shows the number of potassium ions present

two potassium ions +

2–

2+

2–

+

( K O + 2– (K )2(O ))

2–

(not K O , K O , (K )2(O )) The formula is K2O, not KO, K2O.

2–

K

2+

O

K 2O

2–

KO

K2O

(b) Magnesium nitrate –

2+

Mg ion carries 2 positive charges; NO3 ion carries 1 negative charge. To have electrical neutrality, the ratio – 2+ of Mg ions: NO3 ions must be 1 : 2.

(b) Mg

Thus the ionic formula of magnesium nitrate is as shown below:

–

2+

Mg (NO3 ) one magnesium ion 2+

NO 3

2+

–

2+

–

1

2

–

2+

Mg (NO3 )

two nitrate ions –

NO 3 Mg

this number written after the brackets shows the number of nitrate ions present

2

2+

2

–

(not Mg NO3 , Mg NO3 2) The formula is Mg(NO 3 ) 2 , not MgNO 3 , Mg2(NO 3 ), MgNO32. (c)

2+

(

Mg NO3

Sodium hydroxide +

–

+

–

+

MgNO 3 MgNO32

–

The ionic formula is Na OH , not Na (OH )2, Na (OH ), + – Na (OH) .

–

2+

–

Mg NO3 2)

Mg(NO 3 ) 2 Mg2(NO 3 )

(c)

The formula is NaOH, not Na(OH)2, Na(OH). cont'd

+

–

Na OH + – Na (OH )2 Na (OH ) + – Na (OH) +

Na(OH)2

–

NaOH Na(OH) 51


I

(d) Calcium hydroxide

(d) 2+

–

2+

–

2+

The ionic formula is Ca (OH ) 2 , not Ca OH , 2+ – 2+ – Ca OH 2, Ca (OH) 2. The formula is Ca(OH)2, not Ca2(OH), CaOH2. (e)

–

Ca(OH) 2 CaOH2

Iron(III) sulphate 3+

2–

3+

Ca2(OH)

2–

The ionic formula is (Fe ) 2 (SO 4 ) 3 , not Fe SO 4 , 2– 3+ Fe 2(SO4 )3.

–

Ca (OH ) 2 2+ – Ca OH Ca OH 2 2+ – Ca (OH) 2 2+

(e) 2–

3+

The formula is Fe2(SO4)3, not FeSO4, (Fe)2(SO4)3. 3+

Fe SO4

➲ Try Chapter Exercise Q20

FeSO4

2–

(Fe )2(SO4 )3 2– 3+ Fe 2(SO4 )3

Fe2(SO4)3 (Fe)2(SO4)3

➲

20

7.8

Class practice 7.8 Write the chemical formula of each of the following compounds: (a) Copper(II) chloride

(b) Calcium sulphide

(a)

(c) Aluminium hydroxide

(d) Ammonium carbonate

(c)

(II)

(b) (d)

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.

Problem-solving strategy Predicting the formulae of ionic compounds Step 1 Write the formulae of the two ions involved side by side. Mg

2+

1

–

F

Mg

Step 2 Highlight the number of the charge on each ion. Mg

2 +

F

1 –

=

2 +

Mg1

F

–

F

2

Step 3 Take the number of the charge on each ion across to the other. Mg

2+

Mg

2 +

F

1 –

1 –

3

F2

Step 4 Combine the symbols and simplify the ratio.

Mg

MgF2 (Omit the number 1 for Mg)

=

2 +

F

Mg1

1 –

F2

4 MgF2

52

(

Mg

1

)


Study more examples: Aluminium oxide Al

3+

O

2 –

Al2O3

Al

3+

O

2 –

Al2O3

Iron(III) sulphate Fe

3+

2 –

SO4

Fe2(SO4)3

Fe

3+

2 –

SO4

Fe2(SO4)3

Calcium oxide Ca

2 +

O

2 –

Ca2O2

CaO

Ca

(Note: The formula of calcium oxide is CaO but not Ca2O2. This

(

is because the formula of an ionic compound expresses the

C a 2O 2

2 +

O

2 –

Ca2O2

CaO

CaO

simplest whole number ratio of the ions present. Therefore, the ratio of 2 : 2 must be simplified to 1 : 1.)

2

2

1

1

)

7.9

Class practice 7.9 Using the short-cut method, predict the chemical formula of each of the following compounds: (a) Magnesium hydroxide

(b) Sodium oxide

(a)

(c) Lead(II) sulphate

(d) Potassium dichromate

(c)

(b) (II)

(d)

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

1.

‘ion’ is omitted. For example,

+

(Na ) 2–

(CO3 ) (Na 2 CO 3 )

53


I

Cation Al

3+

NH4

+

2+

Ca

Cu Cu Pb

2.

+

2+

2+

Anion SO4

Formula of compound

Name of compound

2–

Al2(SO4)3

aluminium sulphate

2–

(NH4)2CO3

ammonium carbonate

–

Ca(NO3)2

calcium nitrate

CO3

NO3 2–

Cu2O

copper(I) oxide

(I)

2–

CuO

copper(II) oxide

(II)

–

PbBr2

lead(II) bromide

(II)

O O

Br

Some ionic compounds contain water of crystallization.

2.

The number of molecules of water of crystallization (n) has

(n)

to be added at the end of the name as: -n-water. For example, Na2CO3 · 10H2O is called sodium carbonate-10-

Na 2 CO 3 · 10H2O

water.

7.10

Class practice 7.10 Name the following compounds:

54

(a) Ca(NO3)2

(b) FeCl3

(a) Ca(NO3)2

(b)

FeCl3

(c) ZnSO4 · 7H2O

(d) Cu(OH)2

(c) ZnSO4 · 7H2O

(d)

Cu(OH)2


Key terms Page 1. anion

37

2. cation

37

3. chemical bond

47

4. electron ‘dot/cross’ diagram

/

48

5. formula

41

6. giant ionic structure

49

7. ionic bond

48

8. ionic compound

50

9. migration of ion

39

10. polyatomic ion

37

11. simple ion

37

12. transfer of electron

47

13. water of crystallization

54

55


I

Summary

56

7.1


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 one occupied shell) as in helium. Other atoms can also gain great stability if they can get an octet (or duplet).

2.

An A A + NH4 ) are called

7.2


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

7.3

Formulae of ions

5.

A represents the smallest unit (using chemical symbols and numbers) of a substance 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 these are .

8.

For metals in Groups I, II and III, the number of charges on an ion is equal to its group number. For non-metals in Groups V, VI and VII, the number of charges on 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 main types of chemical bonds, namely, bonds, bonds and bonds.

is an atom or a group of atoms having an overall electric charge. is derived from a single atom. + is derived from a group of atoms. Positive ions (e.g. Na , – – ; negative ions (e.g. Cl , MnO4 ) are called .

ions. Refer to Table 7.2 on p.39.

. Some non-metals form ions — most of


7.6


10.

is the strong non-directional electrostatic force of attraction between oppositely charged ions.

11.

When a metal (which tends to electrons) and a non-metal (which tends to electrons) combine, they do so by the transfer of electrons, forming ions. The ions are held together by ionic bonds. For example, +

–

electron

+

Na

Cl

Na

Cl

transfer 2,8,7

2,8,1

2,8 2,8,8 Electron diagram of sodium chloride

7.7

Structures of solid ionic compounds

12.

An

7.8


13.

The formulae of ionic compounds can often be predicted using a short-cut method:

(or ionic substance) is a compound which consists of ions.

X

a

Y

b

⇒

XbYa

(where a, b = ionic charge) e.g.

Zn

2 +

NO3

1–

⇒

Zn(NO3)2

57


8.1

I


8.1

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).

(

)

CO 2

NH3 (

CH4

HCl

8.1) H

O

C

O

H

N

H

Figure 8.1 Molecules of some compounds. carbon dioxide

C

H

H

H

H

Cl

H

ammonia

methane

hydrogen chloride

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;

(Cl2) (N2)

(O2) 2

( (O 3 )

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.

58

(F2)

(P4) 8

(H2)

Ar)

1

3 4

(S8) (Ar)

(O2)

(O3)

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.

8.1


Which of the following represent a molecule?

1.

+

2.

Br2, K , Br, Zn(OH)2, C6H12O6, Ne, Na, NH3, CaO

Br

Write the formulae for the following elements:

C6H12O6

2

K

+

Br Ne

Zn(OH)

Na

NH3

2

CaO

2.

(a) neon

(b) hydrogen

(c) sodium

(d) nitrogen

(e) fluorine

(f) magnesium

(Refer to the Periodic Table if necessary.)

(a)

(b)

(c)

(d)

(e)

(f)

(

)

Covalent bonding Covalent bond formation in a chlorine molecule A molecule usually consists of a number of atoms chemically joined together. Cl

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

(

) Cl2

outermost shell, a chlorine molecule Cl 2 is formed. In the

(

molecule, each chlorine atom has a stable octet (Figure 8.2).

8.2)

a shared pair of electrons forms a single covalent bond

Cl

Cl

+

electron

Cl

Cl

sharing chlorine atom (Cl) (Cl)

chlorine atom (Cl) (Cl)

2,8,7

2,8,7 (both unstable) (

)

chlorine molecule (Cl2) (Cl2) 2,8,8 2,8,8 (more stable) ( )

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). ( )

59


I

It should now be obvious that a chlorine molecule must be

Cl 2

Cl2, and cannot possibly be Cl, Cl3 or Cl4.

Cl

Cl3

Cl4

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.

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

Cl2

Cl

electrostatic attraction.) Cl 2 is the molecular formula of Cl

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.

When we say the ‘formula’ of a molecular substance, we usually mean its ‘molecular formula’.

Covalent bond formation in some molecules

formation in some simple molecules. All of them are molecules of covalent substances.

A covalent substance is a non-ionic substance in which the atoms are held together by covalent bonds.

60

( ) Cl 2

chlorine, while Cl–Cl is the structural formula of chlorine.

Table 8.1 gives electron diagrams to show the covalent bond

Cl

8.1

Cl


Electron diagrams to show covalent bond formation

H

H

H

Molecular formula

Structural formula

H

H

H2 2 hydrogen atoms 2

H

a single covalent bond

1 hydrogen molecule 1

Cl

H

H Cl H

HCl 1 hydrogen atom + 1 chlorine atom 1 +1

Cl

1 hydrogen chloride molecule 1

H H H

C

H

C

H

H

H CH4

H

H

C

H

H

H 1 carbon atom + 4 hydrogen atoms 1 +4

1 methane molecule 1 a lone pair of electrons

H

N

H

N

H

H

H H

a bond pair of electrons

NH3

H

N

H

H 1 nitrogen atom + 3 hydrogen atoms 1 +3

H

O

H

1 oxygen atom + 2 hydrogen atoms 1 +2

1 ammonia molecule 1

H O

H

H2O

H

O

H

1 water molecule 1 a double covalent bond

O

C

O

O

C

O CO2

1 carbon atom + 2 oxygen atoms 1 +2

O

C

O

1 carbon dioxide molecule 1 a triple covalent bond

N

N

N

N N2

2 nitrogen atoms 2

1 nitrogen molecule 1

N

N

Table 8.1 Electron diagrams to show the formation of some simple molecules (only the outermost shell electrons are shown). ( ) 61


I

When non-metal atoms combine with each other, there is usually a sharing of electrons, forming covalent bonds.

Rules for forming covalent bonds Table 8.2 lists out some rules for forming covalent bonds,

8.2 (

illustrated with a few examples. (Refer to Table 8.1 at the same

8.1

)

time.)

Rules (1)

An atom involved in covalent bond formation contributes n electron(s) for sharing.

• •

(2)

•

•

n For hydrogen atoms, n = 1 n=1 For other atoms, n = 8 – group no. of the element n = 8 –

Examples 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 for sharing; ( V ) (8 – 5) 3 an oxygen atom (Group VI) contributes 2 electrons for sharing; ( VI ) (8 – 6) 2 a fluorine atom (Group VII) contributes 1 electron for sharing ( VII ) (8 – 7) 1

For hydrogen and Group VII In H elements, an atom shares electrons H with one other atom in covalent In N bond formation. atom. VII N

Cl , a chlorine atom shares electrons with a hydrogen atom. Cl N , a nitrogen atom shares electrons with another nitrogen N

In H N For other elements, an atom may H share electrons with one or more atoms. other atoms. H

N

H , a nitrogen atom shares electrons with 3 hydrogen

H

H

(3)

2 atoms may share between them •

or

•

or

• • •

62

H

1 electron pair (to form a H C H (CH4) contains 4 single covalent bonds; single covalent bond) H ( ) 2 electron pairs (to form a O=C=O contains 2 double covalent bonds; O=C=O double covalent bond) ( ) N N contains 1 triple covalent bond. 3 electron pairs (to form a N N triple covalent bond) ( )


Rules (4)

Examples

A shared pair of electrons is known as The nitrogen atom in an NH3 molecule has 3 bond pairs and 1 lone a bond pair. pair. Some atoms in a molecule may have NH3 unshared pairs of outermost shell lone pair electrons — known as lone pairs. XX

H

N

H

bond pair

H

In a H2O molecule, the oxygen atom has 2 bond pairs and 2 lone pairs. H2O lone pair H

XX

O

H

XX

bond pair

lone pair

Table 8.2 Rules for forming covalent bonds.

Example 8.1 Identifying some common substances

8.1

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)

(CCl 4 ) (NH4NO3) (LiOH)

(Ag) (CH3COOH) (C7H16) (I2)

Which of them are (a) ionic compounds (b) covalent substances (c)

covalent compounds?

Solution

(a) (b) (c)

(a) Ammonium nitrate, lithium hydroxide (b) Tetrachloromethane, ethanoic acid, heptane, iodine (c)

Tetrachloromethane, ethanoic acid, heptane

(a) (b) (c)

63


I

8.2


1.

Fill in the blanks: Metals tend to metals tend to

electrons, while nonor

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)

2.

(a) (i) (

Find the number of bond pairs and lone pairs on

)

the nitrogen atom in this molecule.

(ii)

(b) Give the (i)

molecular formula,

(ii)

structural formula of the molecule in (a).

(b)

(a) (i) (ii)

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.

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 )

(NH4 )

When ammonia reacts with hydrogen chloride to form ammonium chloride, a dative covalent bond is formed between

NH 3

+

the lone pair of electrons on the N atom in NH3 and a H ion from HCl (Figure 8.3). The symbol ‘ the dative covalent bond.

64

’ is used to represent

HCl (

8.3)

N H

+


H

H H

Cl

H

N

H

H

N

Cl

H

H ammonium ion

H

H or

H

H

Cl

H

N H

N

H

Cl

H dative covalent bond

Figure 8.3 Electron ‘dot/cross’ diagram showing formation of ammonium chloride.

+

+

The ammonium ion (NH4 ) has an overall charge of +1 distributed all over the structure. Thus ammonium chloride +

–

(NH 4 ) +1

(NH4Cl) contains ionic bond (between NH4 and Cl ions) and

(

four covalent bonds (four N – H bonds) — three of the N – H

(

(NH 4 Cl) NH 4

+

Cl

N– H

–

) 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. +

(H3O )

+

Dative covalent bond in hydronium ion (H3O ) H

+

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 .

H Cl

–

H

+

+

+

+

Each H ion is attracted to the unshared electrons of oxygen atom of a water molecule, forming a dative covalent bond. A

H

+

more stable ion, hydronium ion H3O , is obtained as a result. See Figure 8.4.

+

H3O

+

8.4

65


H

I

O

H

Cl

H

O

H

Cl

H

H

hydronium ion

or

H

O

H

Cl

H

H

O

H

Cl

H dative covalent bond

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.

Check your concept

✘ ✔

The constituents of ammonium nitrate (NH4NO3) are all nonmetals, so it is considered as a covalent compound.

✘

Although ammonium nitrate is made up of non-metals only, it is + an ionic compound. It consists of ammonium ion (NH4 ) and – nitrate ion (NO3 ).

✔

(NH4NO3)

+

(NH4 )

66

–

(NO3 )


8.2


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

Hydrogen sulphide

1

2

4

1

H S

H 2 S1

Structural formula H 2S

H

S

H

Cl

Tetrachloromethane

C Cl

C1Cl4

CCl4

Cl

C

Cl

Cl 3

1

Ammonia

N H

Carbon dioxide

C O

4

2

N 1H 3

NH3

C2O4

CO2

H H

N

H

O

C

O

Table 8.3 Predicting formulae of hydrogen sulphide, tetrachloromethane, ammonia and carbon dioxide (using a short cut). ( )

8.3

Class practice 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.

8.3


8.3

(a)

(b)

(c)

(d)

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. 67


I

8.2

Example 8.2 Identifying types of particles that make up substances Complete the table below. Substance/species

Nitrogen

Magnesium

Water

Carbon monoxide

Carbon dioxide

Constituent particle(s)

Formula

Remarks

molecule

N2

nitrogen is an (a) ________________ . (a)

atom

Mg

magnesium is an (b) ________________ . (b)

molecule

H2O

water is a (c) _________ made of (d) __________ . (c) (d)

molecule

CO

the formulae of some compounds can be guessed from their (e) _____________ : thus the formula of carbon monoxide is (f) _____________ ; that of carbon dioxide is (g) _____________ .

molecule

CO2

(e) (f) (g)

Sodium chloride

Hydroxide ion

Air

different types of ions

ion different types of molecules

NaCl

OH

–

not applicable

there is no sodium chloride (h) ________________ to represent the compound sodium chloride. (h)

–

–

OH is an ion OH

air has no formula because it is a (i) ____________ of many substances. (i)

Solution (a) element; (b) element; (c) compound; (d) molecules; (e) names; (f) CO; (g) CO2; (h) molecule; (i) mixture.

Table 8.4 summarizes the constituent particles in various substances.

68

(a) (d) (g) CO2

(b) (e) (h) NaCl

8.4

(c) (f) CO (i)


Constituent particles

Examples

metals

atoms

copper (Cu) (Cu)

non-metals

molecules (exception: carbon) (

argon (Ar) chlorine (Cl2) sulphur (S8) (Ar) (Cl2) (S8)

elements

PURE SUBSTANCES

)

compounds made up of non-metals only

usually molecules

water (H2O) ammonia (NH3) (H2O) (NH3)

compounds made up of metal(s) and non-metal(s)

ions

potassium oxide (K2O) sodium chloride (NaCl) (K2O) (NaCl)

compounds

Table 8.4 Constituent particles of various substances.

8.4

Class practice 8.4 Decide whether the following formulae stand for an atom, a molecule or an ion. (a) CHCl3 (e) S8

8.4

(b) (f)

Ar 2+

Ba

2–

(c)

Cr2O7

(d)

Mg

(g)

I2

(h)

P

(a) (c) (e) (g)


CHCl3 2– Cr2O7 S8 I2

-12 =

12

12.000 00 scale.

Ar Mg 2+ Ba P

8.4

For elements and compounds consisting of molecules, relative molecular mass is the mass of one molecule of it on the C =

(b) (d) (f) (h)

12.000 00 69


I

Learning tip Relative molecular mass can also be called molecular mass. Relative molecular mass carries no units.

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.

(H 2 O) 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 12

one formula unit of it on the C = 12.000 00 scale.

(

)

-12 =

12.000 00

Formula mass of a

=

substance (or species)

Sum of relative atomic masses

(

)

of all atoms present in a formula unit of the substance

Learning tip Formula mass carries no units.

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

70

)


Example 8.3 Determining the formula masses of some substances/ species Calculate the formula mass of (a) C6H12O6

(b) SO4

2–

(c)

Al2(SO4)3

8.3

(a) C6H12O6 (c) Al2(SO4)3

(b) SO4

2–

Solution C6 H12 O6 (a) Formula mass of = 12.0 6 + 1.0 12 + 16.0 6 C6H12O6 = 180.0 Note: We can regard the formula mass of C6H12O6 as the relative molecular mass because the compound actually consists of molecules. S O4 2– (b) Formula mass of SO4 = 32.1 + 16.0 4 = 96.1 (c)

(a) C6H12O6 C6 H12 O6 = 12.0 6 + 1.0 12 + 16.0 6 = 180.0 C 6 H 12 O 6

(b) SO4

Al2 (SO4)3 Formula mass of = 27.0 2 + (32.1 + 16.0 4) 3 Al2(SO4)3 = 342.3

(c)

2–

Al2 (SO4)3 Al2 (SO4)3 = 27.0 2 + (32.1 + 16.0 4) 3 = 342.3

8.5


What is the relative molecular mass of (a) CH4

2.

(b) C2H6

1. (c) C12H22O11? (a) CH4

Calculate the formula mass of (a) NaCl

S O4 = 32.1 + 16.0 4 = 96.1

(b) C2H6

(d) Cu(NO3)2 · 3H2O.

(b) C2H6

(c) C12H22O11

2–

(c) CO3

2. (a) NaCl (b) C2H6 2– (c) CO3 (d) Cu(NO3)2 · 3H2O

71


I

Key terms Page

72

1. atomicity

58

2. bond pair

60

3. covalent bonding

60

4. covalent substance

60

5. double covalent bond

62

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


Summary 8.1


1.

A is the smallest part of an element or a compound which can exist on its own under 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 non-metals (except carbon) consist of discrete .

. All

3.

Molecules can be represented by

4.

A is formed when one or more pairs of outermost shell electrons are shared between two atoms. For example,

H

Cl

1 hydrogen atom + 1 chlorine atom

to show their shapes.

H Cl 1 hydrogen chloride molecule

5.

is the strong directional electrostatic attraction between the shared electrons and the two nuclei of the bonded atoms.

6.

A shared pair of electrons ( bond, e.g. H – Cl. 2 shared pairs of electrons make a

) makes a

covalent

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 formed between two atoms where both electrons of the shared pair are contributed by the same atom.

8.2


8.

The of a molecular substance shows the actual number of each kind of atoms in one molecule of the substance, e.g. CH4.

9.

The of a molecular substance shows how the constituent atoms are joined up in one molecule of the substance, e.g. H

H

C

H

H

73


10.

I

Some atoms have unshared pairs of outermost shell electrons. These are known as e.g. a lone pair of electrons

H O

H

1 water molecule

11.

The formulae of covalent compounds can often be predicted using a short-cut method: X

a

Y

b

(where a, b e.g.

74

Si

4

H

1

⇒

XbYa

= number of electrons contributed for sharing) ⇒

SiH4

8.3


12.

All matter is made up of particles: atoms, molecules or

8.4


13.

The C = 12.000 00 scale is used for comparing

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)

12

.

of atoms.

,

Chapter 9 Structures and properties of substances

9.1


9.1

Introduction The structure of a substance is a description of what its constituent particles are, and about how they are arranged or packed together.

The study of structures is important, since physical properties of a substance are closely related to its structure.

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 (

)

75


I

EXAMPLES Elements Non-metals Simple molecular structures

Compounds Metals

Covalent

Ionic

water H2O carbon dioxide CO2 H2O CO2

hydrogen H2 iodine I2 H2 I2

Molecular structures

polyethene —(CH2CH2 )— n

Macromolecules

—(CH2CH2 )— n SUBSTANCES


Giant structures


sodium chloride NaCl NaCl diamond, (different forms of carbon) (

Figure 9.1 Classification of substances according to structure.

silicon(IV) oxide SiO2 (IV) SiO2

graphite


)

copper Cu iron Fe Cu Fe

9.1

Class practice 9.1 Name the structures possible for (a) non-metal elements (b) covalent compounds.

(a)

(b)

Learning tip For polyethene, the formula is represented by –(CH2CH2)– n , where n is a whole number from 100 to 30 000. Each molecule is very large (hence called macromolecule); it consists of many, usually thousands, of –CH2CH2– groups joined together.

76

–(CH 2CH 2)– n ) –CH2CH2–

n

100 (

30 000 ( )


9.2


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

1

2

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

(

)

CO 2

molecules, each CO2 molecule is attracted to neighbouring molecules. In general, the larger the molecular size, the greater will be

–78.5°C

the van der Waals’ forces between molecules. When carbon dioxide gas is placed under temperatures

9.2

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.

Figure 9.2 The structure of dry ice. The molecules are held by van der Waals' forces in the structure.

77


I

Structure of iodine In an iodine crystal, I2 molecules are packed closely together,

I2

but they are still discrete molecules. The molecules are held by van der Waals’ forces. See Figure 9.3.

9.3

Figure 9.3 The crystal structure of iodine. indicates the position of an I2 molecule. Here the molecules are packed in a regular pattern. Repetition of this pattern trillions of times would result in a crystal. I2

9.2

Class practice 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.)

( )

Properties of simple molecular substances 1.

Simple molecular substances have low melting points and

1.

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.

Learning tip A volatile liquid evaporates quickly under room conditions.

78

(

)


2.

Simple molecular solids are soft. Intermolecular forces are

2.

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.

3. (

)

They are non-conductors of electricity, whether as solids, liquids or in aqueous solution. This is because they do not

4.

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.

Learning tip Solvents other than water are called non-aqueous solvents.

9.3

Class practice 9.3 Answer the following questions. (a) Is sulphur high-melting or low-melting?

(a)

(b) Are sulphur crystals hard?

(b)

(c) Does molten sulphur conduct electricity?

(c)

(d) Is sulphur soluble in (i) water (ii) carbon disulphide, a nonaqueous solvent?

(d)

9.3

Macromolecules

(i)

(ii)

(

)

9.3

Macromolecules are very large molecules, each containing thousands of atoms. Examples are plastics, proteins and some carbohydrates like starch.

(

)

79


9.4

I


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.

8

8

Therefore, the structure of caesium chloride CsCl is different

CsCl

from that of sodium chloride. See Figure 9.4.

Figure 9.4 Caesium chloride has a giant– + ionic structure. It consists of Cs and Cl ions held together by ionic bonds. Cs

–

9.4

or

Cs

–

Cl

+

Cl

Learning tip In the structure of sodium chloride, each sodium ion is surrounded by six chloride ions.

6

Properties of ionic compounds 1.

All ionic compounds are solids. The oppositely charged ions

1.

are attracted together by strong ionic bonds. 2.

They usually have high melting points and boiling points. A

2.

lot of heat energy is required to overcome the strong attractive forces (ionic bonds) between ions in melting and boiling.

80

+

(

)


3.

Most of them are soluble in water, but insoluble in non-

3.

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

4.

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.

Learning tip An electrolyte is a compound which, when molten or in aqueous solution, conducts electricity, and is decomposed at the same time.

Example 9.1 Statements about ionic bonds and covalent bonds

9.1

This question consists of two separate statements. Decide whether each of the two statements is true or false; if both are true, then decide whether or not the second statement is a correct explanation of the first statement. ‘Sodium chloride is a high-melting solid, whereas chlorine is a 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 ionic bonds between ions must be overcome before the compound 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, only the weak van der Waals’ forces between chlorine molecules must be overcome; the strong covalent bond within each molecule is not broken. Thus chlorine is a gas.

81


9.5

I


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

covalent bonds

Figure 9.5 The three-dimensional structure of diamond.

To break the structure, numerous very strong covalent bonds between carbon atoms must be broken. This explains the extreme hardness and very high melting and boiling points of diamond. Diamond cannot conduct electricity, because it contains no ions or freely moving electrons to carry electric charges.

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Two main uses of diamond: (a) Jewellery

(a)

(b) Diamond cutter (used for cutting glass)

(b)

(

)

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).

(

9.6)

strong covalent bonds (within layers) ( )

Figure 9.6 The structure of graphite. (The ‘lead’ pencil is graphite mixed with some clay.) (

weak van der Waals’ forces (between layers) ( )

)

Each carbon atom is covalently bonded to only three other carbon atoms in its layer, and one outer electron of each carbon atom is ‘free’. Those electrons are not attached to any particular atoms but belong to the whole structure (i.e. the electrons are delocalized). They are free to move from one six-membered carbon ring to the next within a layer. Thus, graphite can conduct electricity. Since only van der Waals’ forces exist between adjacent layers, these weak forces make the graphite crystal easy to cleave, and explain its softness and lubricating property. On the other hand, graphite has a very high melting point, since this involves the breaking of strong covalent bonds within the layers.

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Some physical properties of diamond and graphite are

9.1

summarized in Table 9.1.

Property

Diamond

Appearance

colourless

Hardness

black solid

hardest natural substance on the Earth

soft, brittle

3550

3730

non-conductor

conductor (conducts in the direction parallel to hexagonal planes) ( )

Melting point (°C) °

Table 9.1 Some properties of diamond and graphite.

Graphite

Electrical conductivity

Structure and properties of silicon(IV) oxide

(IV)

Both elements and compounds may form giant covalent structures. An example of a compound having a giant covalent

(IV) (

) SiO 2

structure is silicon(IV) oxide (or silicon dioxide) SiO2. In the structure of silicon(IV) oxide, each silicon atom is

(IV)

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

(

9.7)

whole structure.

silicon atom

oxygen atom

Figure 9.7 The giant covalent structure of silicon(IV) oxide. Note that this represents only a very small part of the lattice, which extends in all directions. (IV)

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There are no discrete SiO2 molecules in silicon(IV) oxide.

(IV)

Thus SiO2 is only an empirical formula, not a molecular formula.

SiO 2

SiO 2

This formula shows that the simplest whole number ratio of Si :

Si

O atoms in the compound is 1 : 2.

1

O

2

Because of its structure, silicon(IV) oxide has a very high melting point (1610°C) and boiling point. Also, it does not

(IV)

(1610°C)

conduct electricity whether it is in the solid state or molten.

Check your concept

✘ ✔

Silicon(IV) oxide, with a formula SiO2, has a simple molecular structure.

✘

Silicon(IV) oxide has a giant covalent structure. The formula SiO2 only represents the composition of the elements in the lattice.

✔

Properties of giant covalent structures 1.

Giant covalent structures are all solids with very high melting

1.

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.

2.

( )

3.

4.

( )

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Example 9.2 Statements about covalent substances

9.2

This question consists of two separate statements. Decide whether each of the two statements is true or false; if both are true, then decide whether or not the second statement is a correct explanation of the first statement. ‘Covalent substances are all gases, liquids or low-melting solids.’ ‘Covalent bonds are weak.’ Solution Both statements are false! 1st statement: This is true only for simple molecular structures; covalent substances with a giant covalent structure are solids with very high melting points. 2nd statement: Covalent bonds are strong forces of attraction.

➲ Try Chapter Exercise Q17

➲

17

9.4

Class practice 9.4 (a) Is tridymite (a form of silicon(IV) oxide) soluble in (i) water (ii) heptane? (b) Does molten tridymite conduct electricity?

(a)

( (i)

(b)

9.6


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. 86

2, 8, 1

(IV) (ii)

)


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.

9.5

Class practice 9.5 Calcium has the electronic arrangement 2,8,8,2.

2,8,8,2 (a)

(a) How many outermost shell electrons does a calcium atom have?

(b)

(b) How many delocalized electrons does each calcium atom in the metal contribute?

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.

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• 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


The bonding, structures and properties of substances with

9.7 9.2

simple molecular, giant ionic, giant covalent and giant metallic structures are summarized in Table 9.2.

Simple molecular structure

Giant ionic structure

Giant covalent structure

(1) Examples

H2, I2, H2O, NH3, CCl4

NaCl, CaO, KOH

C (diamond), C (graphite), SiO2 C( ), ), SiO2 C(

All metals

(2) Structure

small discrete molecules e.g. H2

giant lattice of ions e.g. NaCl

giant lattice of atoms e.g. C (diamond)

metal ions, surrounded by a ‘sea’ of freely moving electrons

H2

NaCl

C ( )

88

Giant metallic structure





Giant metallic structure

strong covalent bonds bind atoms together within a molecule; separate molecules are attracted by weak intermolecular forces (e.g. van der Waals’ forces)

ionic bonds link oppositely charged ions throughout the structure

covalent bonds link atoms throughout the network structure

metallic bonds link the metal ions (positively charged) and the ‘sea’ of electrons (negatively charged) ( ) ( )

gases, volatile liquids, or solids of low melting points

solids

solids

solids (except mercury) (

(b) M.p. and b.p.

low

high

very high

usually high

(c) Hardness of solid form

soft

hard

usually high

usually high

(i) most are insoluble

(i) most are soluble

(i) insoluble

(i) insoluble

(ii) generally soluble

(ii) insoluble

(ii) insoluble

(ii) insoluble

non-conductors Note: A few (e.g. sulphuric acid) react with water to form a solution which conducts electricity

non-conductors when solid; good conductors when molten or in aqueous solution

non-conductors (except graphite)

good conductors

(3) Bonds holding constituent particles

( )

(4) Physical properties (a) State at room conditions

)

(d) Solubility in (i) water (ii) non-aqueous solvents (e.g. heptane) ( (e) Conduction of electricity

)

(

(

(

)

) )

Table 9.2 Comparison of different structures and properties.

89


9.8

I


9.8

The flow chart as shown in Figure 9.8 may help us to predict

9.8

structures other than the metallic structure.

Physical properties Is the substance a gas or liquid at room conditions?

Structure yes


no Does the solid have a low melting point?

yes

no Does the substance conduct electricity when molten and in aqueous solution?

yes


no Does the substance have a very high melting point?

yes


Figure 9.8 Predicting the structure of a substance from its physical properties.

Learning tip Metals have a giant metallic structure. Usually we can tell whether a substance is a metal from its electrical conductivity and appearance. At room temperature and pressure, mercury is the only liquid with a giant metallic structure.

90


Class practice 9.6

9.6

The following table gives information about some properties of substances A to D. Substance

M.p. (°C) °

A

D

Electrical conductivity solid

A

70

poor

B

375

poor

C

98

D

1610

molten poor good

good

good

poor

poor

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)

9.9


9.9

Suppose we know what elements make up a given compound. From the group number of the elements, we can predict the bonding and structure of the compound. We can then predict its physical properties. See Example 9.3.

9.3

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Example 9.3 Predicting physical properties of compounds from their bonding Predict the (i) formula (ii) structure (iii) physical properties (melting point, boiling point, hardness, solubility behaviour and electrical conductivity) of the compound formed between (a) potassium and sulphur

9.3

(i)

(ii)

(

)

(b) nitrogen and fluorine.

(a)

Solution

(b)

(a) (i)

(iii)

The compound formed between a metal (potassium) and a non-metal (sulphur) is an ionic compound.

(a) (i)

(

)

(

+

Potassium (Group I) forms K ions; 2–

sulphur (Group VI) forms S ions. The formula of the compound is thus K2S. (ii) It has a giant ionic structure. (1) A solid with a high melting point and boiling point.

S

(2) (3) (4)

The compound formed between non-metals (nitrogen and fluorine) is a molecular compound. Nitrogen (Group V) contributes 3 electrons for sharing;

(b) (i)

(

fluorine (Group VII) contributes 1 electron for sharing. (See p.62 for explanations.) Using the short-cut method to predict its molecular formula:

F

2–

)

(1)

(4) Non-conductor of electricity when solid; conductor when molten and in aqueous solution.

N

VI

K

(iii)

(3) Soluble in water, insoluble in most nonaqueous solvents.

1

(

+

)

(ii)

(2) Hard.

3

I

K2S

(iii) Its physical properties:

(b) (i)

(

(

V

(

VII

)

3 )

( N1F3

)

1 62

)

NF3

The formula of the compound is thus NF3. cont'd

3

1

N

F

NF3 92

N1F3

NF3

)


(ii)

It has a simple molecular structure.

(ii)

(iii)

Its physical properties:

(iii)

(1) A substance with a low melting point and boiling point.

(1)

(2) In solid state, the compound is soft.

(2)

(3) Insoluble in water, soluble in non-aqueous solvents.

(3)

(4) Non-conductor of electricity no matter solid or liquid.

(4)

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.)

9.10 Applications of substances according to their structures

X (a)

(c)

(X VII

Y (b) II

Y

)

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.

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Key terms Page 1. delocalized electron

86

2. giant covalent structure

82

3. giant metallic structure

86

4. giant network

82

5. giant structure

75

6. intermolecular forces

77

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

Summary

94

9.1


1.

The of a substance is a description of what its constituent particles are, and about how they are arranged or packed together.

9.2


2.

In some substances, atoms within a molecule are bound together by strong covalent bonds and each molecule is attracted to other neighbouring molecules by weak .


9.3

Macromolecules

3.

are very large molecules, each containing thousands of atoms. Examples include plastics, proteins and some carbohydrates.

9.4


4.

In ionic compounds, crystals consisting of positive and negative ions are held together by strong non-directional electrostatic attractions. The ions are regularly packed to form a continuous, three-dimensional .

9.5


5.

In a few elements and compounds, the non-metal atoms join together by covalent bonds to form a giant network called .

9.6


6.

Metal atoms are joined to one another in a by , which result from the attraction between a ‘sea’ of and metal ions.

9.7


7.

The structure, bonding and physical properties of simple molecular structure, giant ionic structure, giant covalent structure and giant metallic structure are summarized in Table 9.2 on p.88.

9.8


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


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 the structure, chemical bonding, and other properties of matter.

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HKDSE Chemistry Bridging Programe 1B (1)

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