OxfordAQA International GCSE Chemistry: Chemical bonds: ionic, covalent and metallic (9202)

This guide teaches section 3.2.1, Chemical bonds: ionic, covalent and metallic, of the OxfordAQA International GCSE Chemistry (9202) specification, Version 6.2, for exams May/June 2018 onwards. All ten lettered statements, a to j, are taught here. There are no tiers in this qualification, so none of the content is optional. No required practical is attached to this section.

Electron arrangements come from the earlier unit on atoms and shells, so revise that first if shells feel shaky. Course hub: OxfordAQA International GCSE Chemistry. Tick topics off on the printable checklist, and spot weak areas with a free 10-minute diagnostic.

What this unit covers

Statement (3.2.1)	Your task
a	Say what a compound is
b	Explain bonding as transferring or sharing outer-shell electrons to reach a noble gas arrangement
c	Describe how positive and negative ions form; draw an ion’s shells in square brackets with its charge; link charge to group number
d	State that Group 1 metals form ionic compounds with non-metals, with a 1+ metal ion
e	State that Group 7 elements form ionic compounds with metals, with a 1− halide ion
f	Describe a giant ionic lattice and ionic bonding; know the sodium chloride structure; draw ion formation diagrams
g	Describe covalent bonds; draw the listed molecules; tell simple molecules from giant covalent structures
h	State that compounds of non-metals consist of molecules held together by covalent bonds
i	Describe how the atoms in a metal are arranged
j	Explain metallic bonding with delocalised electrons and draw the standard diagram

How these bonds explain melting points, conductivity and hardness belongs to section 3.2.2, the next unit.

Compounds and why atoms bond (a, b)

A compound is a substance in which atoms of two or more elements are chemically combined. Water is a compound. Oxygen gas, O₂, is not: both bonded atoms belong to one element.

Atoms bond by moving or sharing the electrons in their highest occupied energy level (outer shell). Each atom ends up with the electron arrangement of a noble gas (Group 0). There are two ways to get there:

• Transfer of electrons, which makes ions (ionic bonding).
• Sharing pairs of electrons (covalent bonding).

Ions and their charges (c)

Atoms that transfer electrons become ions. Losing electrons gives a positive ion (more protons than electrons); gaining electrons gives a negative ion. Metals form positive ions; non-metals form negative ions. Compounds made from a metal and a non-metal consist of ions.

The specification shows a sodium ion as a shell diagram: the nucleus, a circle for each shell with its electrons as crosses (2 then 8), all inside square brackets with + at the top right. In writing, the same ion can be given as [2,8]⁺. The proton number never changes.

Worked example 1. Give the electron arrangements of the ions formed by lithium (atomic number 3) and sulfur (atomic number 16).

Lithium atom: 2,1   loses 1 electron
  3 protons, 2 electrons -> charge +1
  Ion: [2]+    same arrangement as helium

Sulfur atom: 2,8,6  gains 2 electrons
  16 protons, 18 electrons -> charge 2-
  Ion: [2,8,8]2-   same arrangement as argon

Lithium’s ion copies helium, not neon: losing one electron empties the second shell.

Charge and group number

The charge on a simple ion follows from the group number.

Group	Electrons lost or gained	Ion charge	Examples
1	lose 1	1+	Li⁺, Na⁺, K⁺
2	lose 2	2+	Mg²⁺, Ca²⁺
3	lose 3	3+	Al³⁺
5	gain 3	3−	N³⁻, P³⁻
6	gain 2	2−	O²⁻, S²⁻
7	gain 1	1−	F⁻, Cl⁻, Br⁻, I⁻

For metals, charge = group number. For non-metals in Groups 5 to 7, charge = group number − 8. Elements in Groups 4 and 0 do not usually form simple ions.

Group 1 and Group 7 compounds (d, e)

Group 1 elements, the alkali metals, all react with non-metal elements to give ionic compounds. In every one, the metal ion has a single positive charge (1+), because each atom has one outer electron to lose. The same 1+ ion forms when an alkali metal reacts with water, for example 2Na + 2H₂O → 2NaOH + H₂.

Group 7 elements, the halogens, all react with metals to give ionic compounds. In every one, the halide ion has a single negative charge (1−), because each atom needs one more electron to fill its outer shell. A halogen also displaces a less reactive halogen from its compounds; this unit only needs the ions involved.

Worked example 2. Lithium is heated in iodine vapour. Write the balanced equation and name the ions in the product.

Lithium: Group 1 -> Li+
Iodine:  Group 7 -> I-  (iodide)
One 1+ balances one 1-, so the formula is LiI

Iodine is diatomic: I2
2Li + I2 -> 2LiI

The product, lithium iodide, is made of Li⁺ ions and I⁻ ions. Any Group 1 metal with any halogen gives a 1:1 formula.

Ionic bonding and the sodium chloride lattice (f)

An ionic compound is a giant structure of ions. What holds it together is the strong electrostatic attraction that each ion feels towards the oppositely charged ions next to it. Because this attraction works in all directions through the lattice, the bonding is not a set of separate links. That is ionic bonding.

Sodium chloride

Sodium chloride is the one ionic structure you need. It is a cube-shaped lattice in which sodium and chloride ions take turns:

• Each Na⁺ ion is surrounded by six Cl⁻ ions.
• Each Cl⁻ ion is surrounded by six Na⁺ ions.
• The pattern repeats in three dimensions throughout the crystal, so there are equal numbers of each ion and the formula is NaCl.

You may be given part of the lattice to complete or label. Keep the alternation going in every direction, along rows, columns and layers.

Dot and cross diagrams for ion formation

A dot and cross diagram shows where each electron goes. Electrons from one atom are dots; electrons from the other are crosses. You may show just the outer shell. Each finished ion goes in square brackets with its charge.

Worked example 3. Show how lithium and sulfur form lithium sulfide, and give its formula.

Before:   Li •        Li •          S  (6 crosses)

Each lithium atom gives its one outer electron (•) to sulfur.
Sulfur needs two, so two lithium atoms are needed.

After:    [Li]+       [Li]+         [ S: 6 x + 2 • ]2-

Charges: 2 x (+1) + (-2) = 0   ->   Li2S

Draw the sulfide ion with eight electrons in its outer shell: its own six crosses and the two dots it received. The lithium ions are shown with no outer electrons left in the second shell.

Covalent bonding (g, h)

When atoms share pairs of electrons, they form covalent bonds. Each shared pair is one bond, and these bonds are strong. Covalent bonding happens between non-metal atoms. Compounds formed only from non-metals consist of molecules, with the atoms in each molecule held together by covalent bonds.

Covalent substances come in two kinds.

Kind	What it is	Named examples
Simple molecules	Small, separate groups of atoms	H₂, Cl₂, O₂, N₂, HCl, H₂O, NH₃, CH₄
Giant covalent structures (macromolecules)	Huge networks where every atom is bonded to neighbours throughout	diamond, silicon dioxide

In diamond, every carbon atom is joined to four other carbon atoms. In silicon dioxide, every silicon atom is joined to four oxygen atoms, and every oxygen atom to two silicon atoms. That gives the 1:2 ratio in SiO₂, even though there are no separate SiO₂ molecules. Notice that silicon dioxide, a compound of two non-metals, is giant rather than made of small molecules.

Three ways to show a covalent molecule

The specification illustrates ammonia, NH₃, in three forms:

1. Dot and cross with circles: shared pairs sit where the outer shells overlap.
2. Dot and cross without circles: the same electrons written around the symbols.
3. Displayed formula: each shared pair becomes a line.

The specification names water, ammonia, hydrogen, hydrogen chloride, methane and oxygen as molecules to represent in these forms, so be ready to draw any of them.

Molecule	Shared pairs (bonds)	Non-bonding pairs	Displayed formula
H₂	1	0	H–H
HCl	1	3 on Cl	H–Cl
H₂O	2	2 on O	H–O–H
NH₃	3	1 on N	N with three H–N lines
CH₄	4	0	C with four C–H lines
O₂	2 (a double bond)	2 on each O	O=O

Check any diagram: total outer electrons of the atoms = 2 × (shared pairs + non-bonding pairs). For oxygen, 6 + 6 = 12 = 2 × (2 + 4).

Worked example 4. Phosphorus is in Group 5 and bromine is in Group 7. Draw a dot and cross diagram for phosphorus tribromide, PBr₃.

P has 5 outer electrons (crosses); each Br has 7 (dots).
P needs 3 more; each Br needs 1 more.
So P shares one pair with each of the 3 Br atoms.

Each P-Br overlap:     one x (from P) + one • (from Br)
Left on P:             one pair, xx
Left on each Br:       three pairs, ••  ••  ••
Displayed formula:     Br-P-Br, with the third Br bonded to P

Shared pairs: 3
Non-bonding pairs: 1 on P + 3 on each Br = 10
Check: 5 + 7 + 7 + 7 = 26 = 2 x (3 + 10)

Phosphorus looks like the nitrogen in ammonia, but each bromine also keeps three non-bonding pairs.

Worked example 5. Carbon is in Group 4, oxygen and sulfur in Group 6. Carbonyl sulfide is O=C=S. Explain the bonding.

C: 4 outer electrons, needs 4 more
O and S: 6 outer electrons each, each needs 2 more

C shares 2 pairs with O and 2 pairs with S -> two double bonds
O and S each keep 2 non-bonding pairs

Check: 6 + 4 + 6 = 16 = 2 x (4 + 4)

Every atom now has eight outer electrons.

Recognising structures from diagrams

Worked example 6. Diagram A shows groups of two atoms, each pair joined by one line, with gaps between the pairs. Diagram B shows atoms each joined by lines to four others, with the pattern running off every edge. Identify each kind of structure.

A is simple molecular: the bonds stop at each pair, and the diagram shows many separate molecules. B is a giant covalent structure: the bonded network continues in all directions, so there is no end to a “molecule”. If bonds run off the edges of a diagram, the structure is giant.

Metallic bonding (i, j)

A metal is a giant structure in which the atoms are packed in a regular pattern. Each atom’s outer-shell electrons are delocalised: no longer attached to a particular atom, they can move anywhere in the metal.

The picture is a lattice of positive ions with delocalised electrons between them. The electrons hold the ions together by strong electrostatic attractions between the positive ions and the negative electrons. The specification’s diagram shows rows of circles marked +, sitting in a shaded background that represents the delocalised electrons, labelled “delocalised electrons”. Drawing small e⁻ symbols between the ions shows the same idea.

Worked example 7. Describe the bonding in solid calcium. Calcium is in Group 2.

Each Ca atom (2,8,8,2) releases its 2 outer electrons
  -> Ca2+ ions, arrangement [2,8,8]
The released electrons are delocalised through the lattice
Strong electrostatic attraction: Ca2+ ions <-> delocalised electrons

The ions attract the delocalised electrons, not each other.

Common errors

• Writing that sodium chloride is made of “molecules” or that one Na⁺ bonds to one Cl⁻. It is a giant lattice of ions.
• Leaving the brackets or charge off an ion.
• Giving a non-metal ion a positive charge, or a metal ion a negative one.
• Drawing a shared pair with both electrons from the same atom.
• Leaving out the non-bonding pairs on oxygen in water or on nitrogen in ammonia.
• Saying metals contain “free atoms” or “free protons” instead of delocalised electrons.

Where to go next

Run through the revision notes, then try the practice questions, which have fully worked answers.

Official syllabus

Source document: the OxfordAQA International GCSE Chemistry (9202) specification, Version 6.2, for exams May/June 2018 onwards, from OxfordAQA. This page follows its section 3.2.1, Chemical bonds: ionic, covalent and metallic.