Pearson Edexcel International GCSE Chemistry 4CH1: Ionic, covalent and metallic bonding – Study Guide

This guide teaches sub-topics 1(f) Ionic bonding, 1(g) Covalent bonding and 1(h) Metallic bonding (points 1.37 to 1.54C) of the Pearson Edexcel International GCSE Chemistry (4CH1) specification, Issue 3 (September 2024). The qualification is untiered. Points 1.37 to 1.51 are assessed on both papers. Points 1.52C, 1.53C and 1.54C carry a C reference, so they are Paper 2 only, and they are labelled below.

The short overview of the whole topic is in the Principles of chemistry study guide. This page goes into the bonding sub-topics in depth. For quick recall use the bonding revision notes, then test yourself with the bonding practice questions. Course hub: Edexcel IGCSE Chemistry; printable checklist.

What this unit covers

Point	What you must be able to do	Paper
1.37	Explain how ions form by losing or gaining electrons	Both
1.38	Know the charges of the listed ions	Both
1.39	Write formulae for compounds of those ions	Both
1.40	Draw dot-and-cross diagrams for ionic compounds (Groups 1, 2, 3 with 5, 6, 7)	Both
1.41	Explain ionic bonding as electrostatic attraction	Both
1.42	Explain why giant ionic lattices have high melting and boiling points	Both
1.43	Know when ionic compounds conduct	Both
1.44-1.45	Define a covalent bond and explain it as electrostatic attraction	Both
1.46	Draw dot-and-cross diagrams for the listed molecules	Both
1.47-1.48	Explain the properties of simple molecular substances and the trend with relative molecular mass	Both
1.49-1.50	Explain giant covalent structures, including diamond, graphite and C60 fullerene	Both
1.51	Know that covalent compounds do not usually conduct	Both
1.52C-1.54C	Draw a metallic lattice, explain metallic bonding and the properties of metals	Paper 2 only

1(f) Ionic bonding

How ions form (1.37)

Metal atoms lose their outer electrons and become positive ions. Non-metal atoms gain electrons to fill their outer shell and become negative ions. In both cases the ion ends up with the same electronic configuration as a noble gas.

Worked example 1. Describe how an aluminium atom forms an aluminium ion.

Aluminium: atomic number 13, configuration 2.8.3
It loses its 3 outer electrons
13 protons, 10 electrons  ->  charge 13 - 10 = +3
Ion: Al3+, configuration 2.8

The number of protons never changes.

Charges you must know (1.38)

Ion type	Charge	Examples
Group 1 metals	1+	Li⁺, Na⁺, K⁺
Group 2 metals	2+	Mg²⁺, Ca²⁺
Group 3 metals	3+	Al³⁺
Group 5 non-metals	3−	N³⁻
Group 6 non-metals	2−	O²⁻, S²⁻
Group 7 non-metals	1−	F⁻, Cl⁻, Br⁻
Listed metal ions	as shown	Ag⁺, Cu²⁺, Fe²⁺, Fe³⁺, Pb²⁺, Zn²⁺
Listed other ions	as shown	H⁺, OH⁻, NH₄⁺, CO₃²⁻, NO₃⁻, SO₄²⁻

Learn the last two rows by heart.

Writing formulae (1.39)

The total positive charge must equal the total negative charge. Find the lowest common multiple of the two charges, then work out how many of each ion you need. Put brackets round a group ion such as NO₃⁻ or SO₄²⁻ when you need more than one of it.

Worked example 2. Magnesium nitride.

Mg2+ and N3-
LCM of 2 and 3 = 6
3 x Mg2+ = +6,  2 x N3- = -6
Formula: Mg3N2

Worked example 3. Iron(III) sulfate.

Fe3+ and SO4 2-
LCM of 3 and 2 = 6
2 x Fe3+ = +6,  3 x SO4 2- = -6
Formula: Fe2(SO4)3

The Roman numeral in a name gives the charge on the metal ion: iron(II) is Fe²⁺ and iron(III) is Fe³⁺.

Dot-and-cross diagrams for ionic compounds (1.40)

The diagram shows electron transfer. Only outer electrons need be shown. Use dots for one atom’s electrons and crosses for the other’s, so you can see where each electron went. Each ion goes in square brackets with its charge at the top right.

Worked example 4. Sodium oxide, Na₂O.

Before:  Na (2.8.1)  Na (2.8.1)  O (2.6)
Each Na loses 1 electron; O gains 2

After:
[Na]+    [Na]+    [ O: 6 dots + 2 crosses ]2-

Two sodium atoms are needed because each gives only one electron and oxygen needs two. Draw the oxide ion with eight outer electrons: its own six (dots) and the two it gained (crosses). Each sodium ion can be drawn with an empty outer shell or with its full second shell. Either way, the brackets and the 1+ charge must be shown.

Ionic bonding and lattices (1.41-1.43)

An ionic bond is the strong electrostatic attraction between oppositely charged ions. In a solid ionic compound the ions are arranged in a regular three-dimensional pattern called a giant ionic lattice. Each ion attracts the oppositely charged ions around it in every direction.

• High melting and boiling points (1.42). Melting means overcoming the attractions between the ions. There are a very large number of strong electrostatic attractions throughout the lattice, so a lot of energy is needed.
• Conductivity (1.43). Ionic compounds do not conduct when solid. They do conduct when molten or in aqueous solution.

The full explanation belongs to electrolysis (1.56C, Paper 2 only): ions must be free to move to carry charge, and in the solid they are held in fixed positions.

1(g) Covalent bonding

What a covalent bond is (1.44-1.45)

A covalent bond is a shared pair of electrons between two atoms. It forms between non-metal atoms. The bond holds the atoms together because of electrostatic attraction between the shared pair of electrons (negative) and the nuclei of both atoms (positive). Writing “the atoms share electrons” alone describes the bond, but does not explain why it holds.

Two shared pairs make a double bond; three make a triple bond.

Dot-and-cross diagrams for molecules (1.46)

Draw overlapping outer-shell circles with each shared pair in the overlap, one electron from each atom. After bonding, every atom except hydrogen should have eight outer electrons; hydrogen has two. Show the lone pairs (non-bonding pairs) as well.

Worked example 5. Water, H₂O.

      x x
  H •x O x• H
      x x

Oxygen (2.6) has six outer electrons (crosses). Each hydrogen brings one (dot). There are two shared pairs and two lone pairs on oxygen, giving oxygen eight outer electrons and each hydrogen two.

Worked example 6. Oxygen, O₂.

   x x            • •
   O    (xx ••)    O
   x x            • •

Each oxygen needs two more electrons, so the atoms share two pairs: a double bond (the four electrons in brackets sit in the overlap of the two circles). Each oxygen keeps two lone pairs. Check: 2 shared pairs plus 4 lone pairs = 12 electrons, which equals 6 + 6.

Check any molecule: total outer electrons = 2 × shared pairs + 2 × lone pairs.

Molecule	Shared pairs	Lone pairs (whole molecule)
H₂	1	0
N₂	3 (triple bond)	2
Cl₂	1	6
HCl	1	3 (on Cl)
NH₃	3	1 (on N)
CO₂	4 (two double bonds)	4 (two on each O)
CH₄	4	0
C₂H₆	7	0
C₂H₄	6 (one C=C)	0
CH₃Cl	4	3 (on Cl)

In any halogen-containing molecule the halogen forms one shared pair and keeps three lone pairs.

Simple molecular substances (1.47-1.48)

A simple molecular substance is made of small molecules. Inside each molecule the covalent bonds are strong. Between the molecules there are only weak intermolecular forces of attraction. When the substance melts or boils, only these weak forces are overcome. The covalent bonds do not break. Little energy is needed, so the substance is a gas, a liquid, or a solid with a low melting point.

As relative molecular mass increases, the intermolecular forces get stronger, so more energy is needed to separate the molecules. Melting and boiling points therefore increase, in general.

Worked example 7. Explain the boiling points of the halogens.

Halogen	Mr	Boiling point / °C	State at 25 °C
F₂	38	−188	gas
Cl₂	71	−34	gas
Br₂	160	59	liquid
I₂	254	184	solid

All four are simple molecular. From F₂ to I₂ the relative molecular mass increases, so the intermolecular forces become stronger. More energy is needed to overcome them, so the boiling point rises. The covalent bonds inside the molecules are not broken by boiling.

Giant covalent structures (1.49-1.50)

In a giant covalent structure, every atom is joined to its neighbours by covalent bonds in a continuous network. To melt it, covalent bonds must be broken, and there are a very large number of strong bonds. So giant covalent substances are solids with high melting and boiling points.

Form of carbon	Structure	Hardness	Conducts?
Diamond	Giant covalent; each C bonded to 4 others	Very hard: rigid 3-D network	No: all outer electrons in bonds
Graphite	Giant covalent layers of hexagons; each C bonded to 3 others	Soft and slippery: weak forces between layers let them slide	Yes: one delocalised electron per carbon moves along the layers
C60 fullerene	Simple molecular: separate C60 molecules, each C bonded to 3 others	Soft: only weak intermolecular forces between molecules	Poor: electrons cannot move freely from one molecule to the next

C60 is carbon only, but it is not a giant structure: each molecule is a hollow ball of 60 atoms (20 hexagons, 12 pentagons). Separating the molecules needs far less energy than breaking diamond’s covalent network.

Conductivity of covalent compounds (1.51)

Covalent compounds do not usually conduct electricity. Graphite is the exception you must know. Explaining why covalent compounds do not conduct (1.55C) belongs to electrolysis and is Paper 2 only.

1(h) Metallic bonding – Paper 2 only

The 2-D lattice diagram (1.52C)

Draw rows of circles, closely packed and regularly arranged, each marked with a + charge to show a positive metal ion. Draw electrons, labelled e⁻, scattered in the spaces between the ions. Label them “delocalised electrons” and label the circles “positive ions”. For magnesium you may mark each ion 2+.

 (+) e- (+) e- (+) e- (+)
 e- (+) e- (+) e- (+) e-
 (+) e- (+) e- (+) e- (+)

Metallic bonding (1.53C)

A metal is a giant lattice of positive metal ions surrounded by delocalised electrons. The outer electrons have left the atoms and are free to move through the whole structure. Metallic bonding is the electrostatic attraction between the positive metal ions and the delocalised electrons.

Properties of metals (1.54C)

• Electrical conductivity. The delocalised electrons are free to move through the lattice. When a potential difference is applied, they flow and carry charge. This works in the solid and the liquid.
• Malleability. The positive ions are arranged in layers. When a force is applied, the layers can slide over one another. The delocalised electrons still attract the ions in their new positions, so the metal changes shape instead of breaking.
• High melting points (typical). The many strong attractions need a lot of energy to overcome.

Common errors

• Saying a metal ion forms by “gaining protons”. It loses electrons.
• Writing Al₂(SO₄)₃ without brackets as Al₂SO₄₃.
• Drawing an ionic compound with shared electrons in overlapping circles. Ionic diagrams show transfer and brackets.
• Leaving off the lone pairs on N, O or Cl in a covalent diagram.
• Saying covalent bonds break when a simple molecular substance boils.
• Calling C60 a giant covalent structure.
• Saying graphite is soft because its covalent bonds are weak. The forces between layers are weak.
• Calling the particles in a metal lattice “atoms” or “protons” in a bonding answer. Use “positive ions”.

Next steps

Try the bonding practice questions and the wider Principles of chemistry practice set. The free 4CH1 all-topics diagnostic shows where to focus.

Official syllabus

Pearson Edexcel International GCSE in Chemistry (4CH1) specification, Issue 3, September 2024, published by Pearson Education Limited – Topic 1 Principles of chemistry, sub-topics (f) Ionic bonding, (g) Covalent bonding and (h) Metallic bonding, points 1.37 to 1.54C.