Pearson Edexcel International GCSE Chemistry 4CH1: Ionic, covalent and metallic bonding – Revision Notes

For full explanations and worked examples, use the bonding study guide. These notes condense 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 1.52C to 1.54C on metallic bonding are Paper 2 only.

The topic-wide summary is in the Principles of chemistry revision notes. Course hub: Edexcel IGCSE Chemistry; printable checklist.

Definitions to learn word for word

Term	Definition
Ion	A charged particle formed when an atom loses or gains electrons
Ionic bond	Strong electrostatic attraction between oppositely charged ions
Giant ionic lattice	A regular 3-D arrangement of oppositely charged ions
Covalent bond	A shared pair of electrons between two atoms
Why a covalent bond holds	Electrostatic attraction between the shared pair and the nuclei of both atoms
Intermolecular forces	Weak forces of attraction between molecules
Metallic bonding (Paper 2 only)	Electrostatic attraction between positive metal ions and delocalised electrons

1(f) Ionic bonding

Forming ions (1.37)

• Metals lose outer electrons → positive ions.
• Non-metals gain electrons → negative ions.
• The ion has a full outer shell (a noble gas configuration). Protons do not change.

Atom	Configuration	Ion	Ion configuration
Li	2.1	Li⁺	2
Mg	2.8.2	Mg²⁺	2.8
Al	2.8.3	Al³⁺	2.8
K	2.8.8.1	K⁺	2.8.8
N	2.5	N³⁻	2.8
O	2.6	O²⁻	2.8
Cl	2.8.7	Cl⁻	2.8.8

Ion charges (1.38)

1+	2+	3+	1−	2−	3−
Group 1, Ag⁺, H⁺, NH₄⁺	Group 2, Cu²⁺, Fe²⁺, Pb²⁺, Zn²⁺	Group 3, Fe³⁺	Group 7, OH⁻, NO₃⁻	Group 6, CO₃²⁻, SO₄²⁻	Group 5 (N³⁻)

Group 1 → 1+, Group 2 → 2+, Group 3 → 3+. Group 5 → 3−, Group 6 → 2−, Group 7 → 1−. Roman numerals give the metal ion charge: iron(II) = Fe²⁺, iron(III) = Fe³⁺.

Method: writing a formula (1.39)

1. Write both ions with charges.
2. Find the lowest common multiple (LCM) of the charges.
3. Number of each ion = LCM / its charge.
4. Bracket any group ion used more than once.
5. Check: total + charge = total - charge.

Small reminder: Ca²⁺ and NO₃⁻ → LCM 2 → one Ca²⁺, two NO₃⁻ → Ca(NO₃)₂.

Ionic dot-and-cross (1.40)

• Only Groups 1, 2, 3 combined with Groups 5, 6, 7 are required.
• Show only outer electrons; dots for one element, crosses for the other.
• Show transfer: square brackets round each ion, charge outside the bracket at the top right.
• Negative ion: 8 outer electrons, with the gained ones drawn as the other symbol.
• Number of ions drawn must match the formula (e.g. two Na⁺ for one O²⁻).

Structure and properties (1.41-1.43)

Property	Reason
High melting and boiling points	Many strong electrostatic attractions between oppositely charged ions throughout the lattice need a lot of energy to overcome
Solid does not conduct	Ions held in fixed positions
Molten or aqueous does conduct	Ions free to move and carry charge

1(g) Covalent bonding

Covalent dot-and-cross (1.46)

• Overlap the outer-shell circles; each shared pair goes in the overlap.
• Each atom (except H) ends with 8 outer electrons; H ends with 2.
• Draw every lone pair.
• Check: total outer electrons = 2 × shared pairs + 2 × lone pairs.

Molecule	Shared pairs	Lone pairs	Outer electrons
H₂	1	0	2
O₂	2 (double)	4	12
N₂	3 (triple)	2	10
Cl₂	1	6	14
HCl	1	3	8
H₂O	2	2	8
NH₃	3	1	8
CO₂	4 (two double)	4	16
CH₄	4	0	8
C₂H₆	7	0	14
C₂H₄	6 (one double)	0	12
CH₃Cl	4	3	14

The list in the specification covers: H₂, O₂, N₂, the halogens, the hydrogen halides, water, ammonia, carbon dioxide, and organic molecules with up to two carbon atoms (including methane, ethane, ethene and molecules containing halogen atoms).

Simple molecular substances (1.47-1.48)

• Strong covalent bonds within molecules; weak intermolecular forces between molecules.
• Melting or boiling overcomes only the intermolecular forces → little energy → gases, liquids, or low-melting solids.
• Higher relative molecular mass → stronger intermolecular forces → higher melting and boiling points (in general).

Small reminder (halogens, Mr from A_r F = 19, Cl = 35.5, Br = 80, I = 127):

	F₂	Cl₂	Br₂	I₂
Mr	38	71	160	254
State at 25 °C	gas	gas	liquid	solid

Giant covalent structures (1.49-1.50)

• Network of strong covalent bonds throughout → melting breaks covalent bonds → high melting and boiling points; solids.

	Diamond	Graphite	C60 fullerene
Structure type	Giant covalent	Giant covalent (layers)	Simple molecular
Bonds per carbon	4	3	3
Hardness	Very hard	Soft, slippery	Soft
Why	Rigid 3-D network	Weak forces between layers; layers slide	Weak intermolecular forces
Electrical conductivity	None	Conducts	Poor
Why	No free electrons	One delocalised electron per C, moves along layers	Electrons cannot move from molecule to molecule

Conductivity (1.51)

Covalent compounds do not usually conduct electricity. Graphite is the exception to learn.

1(h) Metallic bonding – Paper 2 only

2-D lattice diagram (1.52C)

• Regular, closely packed rows of circles, each labelled + (or 2+ for magnesium): positive ions.
• Electrons (e⁻) drawn between the ions: delocalised electrons.
• Both labels are needed.

Bonding and properties (1.53C-1.54C)

Property	Explanation
Bonding	Electrostatic attraction between positive metal ions and delocalised electrons
Conducts electricity (solid and liquid)	Delocalised electrons free to move through the structure and carry charge
Malleable	Layers of positive ions slide over each other; attraction to the delocalised electrons is kept
Usually high melting point	Many strong attractions to overcome

Must-know distinctions

• Ionic diagram vs covalent diagram: brackets and charges (transfer) vs overlapping circles (sharing). Never mix them.
• Simple molecular vs giant covalent: melting overcomes intermolecular forces vs breaks covalent bonds.
• Graphite vs C60: both have three bonds per carbon, but graphite is a giant layered structure that conducts; C60 is molecular and a poor conductor.
• Ionic vs metallic conduction: ions must be free to move (molten or aqueous only) vs delocalised electrons move (solid too).
• Lose vs gain: metals lose electrons; nothing gains protons.

Quick self-test

1. Give the electronic configuration of the sulfide ion.
2. Write the formula of zinc hydroxide.
3. Write the formula of copper(II) nitrate.
4. Write the formula of potassium sulfate.
5. In forming potassium sulfide, how many electrons does each potassium atom lose, how many does the sulfur atom gain, and what is the formula?
6. How many shared pairs and lone pairs are in a molecule of hydrogen bromide?
7. How many shared pairs and lone pairs are in a molecule of dichloromethane, CH₂Cl₂?
8. How many shared pairs of electrons are in a molecule of ethane?
9. Explain why chlorine is a gas at room temperature but iodine is a solid.
10. Explain why C60 is soft although diamond is very hard.
11. Explain why solid potassium bromide does not conduct but its aqueous solution does.
12. (Paper 2 only) State what is meant by metallic bonding.

Answers

1. 2.8.8 (sulfur 2.8.6 gains two electrons).
2. Zn(OH)₂.
3. Cu(NO₃)₂.
4. K₂SO₄.
5. Each K loses 1; S gains 2; formula K₂S.
6. 1 shared pair, 3 lone pairs (on Br).
7. 4 shared pairs, 6 lone pairs (three on each Cl). Check: 4 + 2 + 14 = 20 = 2 × 4 + 2 × 6.
8. 7 (one C–C and six C–H).
9. Both are simple molecular. Iodine has a higher Mr (254 vs 71), so its intermolecular forces are stronger and need more energy to overcome.
10. C60 is simple molecular: only weak intermolecular forces hold the molecules together. Diamond is giant covalent: every carbon is held by four strong covalent bonds in a rigid network.
11. In the solid the ions are fixed in the lattice. In solution the ions are free to move and carry charge.
12. The electrostatic attraction between positive metal ions and delocalised electrons.

Where marks are usually lost

• Writing “electrostatic attraction between atoms” in an ionic bonding answer. It is between oppositely charged ions.
• Missing the “shared pair” wording: “atoms share electrons” alone is vaguer than the definition, which needs a pair of electrons.
• Forgetting brackets: Fe₂(SO₄)₃ written as Fe₂SO₄₃, or (NH₄)₂CO₃ as NH₄2CO₃.
• Putting the charge inside the bracket, or omitting it, on an ionic dot-and-cross diagram.
• Drawing too few ions, e.g. one Na⁺ with O²⁻.
• Leaving lone pairs off N, O or Cl in a covalent diagram.
• “Covalent bonds break when it boils” for a simple molecular substance.
• Explaining graphite’s softness with weak covalent bonds rather than weak forces between layers.
• In metals, saying “free electrons and positive atoms”: write positive ions and delocalised electrons.
• Saying ionic compounds conduct because “electrons move”. It is the ions.

Ready to practise? Go to the bonding practice questions, the Principles of chemistry practice set, or the free 4CH1 all-topics diagnostic.

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.