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

These notes condense 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. Statements a to j are all here, and with no tiers in this qualification every line is examinable. The full teaching, with seven worked examples, is in the chemical bonds study guide.

When you have finished, test yourself with the chemical bonds practice questions. Course hub: OxfordAQA International GCSE Chemistry. Track progress on the printable checklist, and try a free 10-minute diagnostic to see which topics need work.

Key definitions (3.2.1 a to c)

• Compound: a substance in which atoms of two or more elements are chemically combined.
• Chemical bonding: transferring or sharing electrons in the highest occupied energy levels (outer shells) so that atoms reach a noble gas electron arrangement.
• Ion: an atom that has gained or lost electrons, so it carries an overall charge.
• Positive ion: formed by losing electrons. Metals do this.
• Negative ion: formed by gaining electrons. Non-metals do this.
• Ionic bonding: the strong electrostatic pull between oppositely charged ions, which works in every direction through a giant lattice.
• Covalent bond: a pair of electrons shared by two atoms; it is a strong bond.
• Delocalised electrons: outer electrons in a metal that no longer belong to one atom and can roam throughout the metal.

Ion charges from the group number (c, d, e)

Group	Change to outer shell	Ion charge	Ion arrangement example
1	loses 1 electron	1+	K⁺ is [2,8,8]⁺
2	loses 2 electrons	2+	Mg²⁺ is [2,8]²⁺
3	loses 3 electrons	3+	Al³⁺ is [2,8]³⁺
5	gains 3 electrons	3−	N³⁻ is [2,8]³⁻
6	gains 2 electrons	2−	O²⁻ is [2,8]²⁻
7	gains 1 electron	1−	Cl⁻ is [2,8,8]⁻

Every ion in the table has a noble gas arrangement: [2,8] is neon’s and [2,8,8] is argon’s.

• Group 1 (alkali metals) react with non-metals to form ionic compounds containing a 1+ metal ion.
• Group 7 (halogens) react with metals to form ionic compounds containing a 1− halide ion.
• Compounds of a metal with a non-metal consist of ions.

Method: showing the electron arrangement of an ion

1. Write the atom’s electron arrangement from its atomic number.
2. Metal: remove the outer-shell electrons. Non-metal: add electrons until the outer shell holds eight.
3. Charge = protons − electrons.
4. Draw the new shells (or write the arrangement) inside square brackets with the charge at the top right. The specification draws Na⁺ as a shell diagram with 2 and 8 electrons in brackets, charge +; written out, that is [2,8]⁺.

Ionic structures (f)

• An ionic compound is a giant structure of ions, not molecules.
• Sodium chloride: a regular cubic lattice. Each Na⁺ is surrounded by six Cl⁻; each Cl⁻ by six Na⁺. Na⁺ and Cl⁻ alternate in every direction. This is the only ionic structure you need to know.
• The attractions act in all directions, so each ion is held by all its oppositely charged neighbours.

Method: dot and cross diagram for an ionic compound

1. Find each atom’s outer electrons from its group.
2. Work out the electrons lost by one metal atom and gained by one non-metal atom.
3. Use enough atoms of each so that electrons lost = electrons gained.
4. Draw the metal’s electrons as dots and the non-metal’s as crosses (or the other way round).
5. Enclose every ion in square brackets and write its charge just outside, at the upper right.
6. Check that the charges add up to zero, then write the formula.

Worked reminder: barium oxide.

Ba (Group 2): loses 2 dots   -> [Ba]2+
O  (Group 6): gains 2 dots   -> [O: 6 x + 2 •]2-
+2 and -2 balance with one of each  ->  BaO

Covalent structures (g, h)

• Covalent bonds form between non-metal atoms by sharing pairs of electrons.
• Compounds formed from non-metals consist of molecules held together by covalent bonds.
• Simple molecules: H₂, Cl₂, O₂, N₂, HCl, H₂O, NH₃, CH₄.
• Giant covalent structures (macromolecules): diamond, silicon dioxide.

Molecule	Bonds	Non-bonding pairs	Notes
H₂	1 single	0	each H has 2 outer electrons
HCl	1 single	3 (on Cl)	Cl has 8 outer electrons
H₂O	2 single	2 (on O)	O has 8 outer electrons
NH₃	3 single	1 (on N)
CH₄	4 single	0
O₂	1 double	2 on each O	4 shared electrons

Method: dot and cross diagram for a molecule

1. Count each atom’s outer electrons from its group (hydrogen has 1).
2. Work out how many more each atom needs: 2 for hydrogen, 8 for the rest.
3. Share one pair for each electron needed, one electron from each atom per pair.
4. Add the remaining outer electrons as non-bonding pairs.
5. Check: total outer electrons = 2 × (shared pairs + non-bonding pairs).

You also need to recognise the other two forms the specification shows: the dot and cross without circles, and the displayed formula, where each shared pair becomes a line (H–Cl, H–O–H).

Worked reminder: tetrachloromethane, CCl₄. Carbon (Group 4) has 4 outer electrons and needs 4 more, so it shares one pair with each of four chlorine atoms. Each chlorine keeps three non-bonding pairs. That is 4 single bonds and 12 non-bonding pairs. Check: 4 + (4 × 7) = 32 = 2 × (4 + 12).

Telling simple from giant in a diagram

• Simple molecular: small, separate units; bonds stop at the edge of each unit.
• Giant covalent: bonds link every atom to its neighbours, and the network runs off the edges of the diagram.
• Diamond: each carbon bonded to four other carbons.
• Silicon dioxide: each silicon bonded to four oxygens, each oxygen to two silicons, giving the 1:2 ratio.

Metallic structures (i, j)

• A metal is a giant structure, its atoms packed in a regular pattern.
• The outer electrons are delocalised: they can move anywhere in the metal.
• Think of it as positive ions in a regular lattice with electrons between them.
• Strong electrostatic attractions between the positive ions and the delocalised electrons hold the structure together.
• The specification’s diagram: rows of circles labelled +, in a shaded background labelled “delocalised electrons” (e⁻ symbols drawn between the ions show the same idea).
• A Group 1 metal gives one delocalised electron per atom and forms 1+ ions in the lattice; a Group 2 metal gives two and forms 2+ ions.

Must-know distinctions

Ionic	Covalent	Metallic
Metal + non-metal	Non-metal + non-metal	Metal atoms only
Electrons transferred	Electron pairs shared	Outer electrons delocalised
Oppositely charged ions attract	Shared pair holds two atoms together	Positive ions attract delocalised electrons
Always giant lattice	Simple molecules or giant covalent	Always giant structure

Two more:

• Molecule vs giant structure: a molecule has a fixed, small number of atoms. A giant structure has no fixed size; its formula is only a ratio.
• Ion vs atom: an ion has the same protons as its atom but a different number of electrons.

Quick self-test

1. Is chlorine gas, Cl₂, a compound? Give a reason.
2. Give the electron arrangement of a beryllium ion, Be²⁺, written in square brackets with its charge.
3. Phosphorus has atomic number 15. Give the charge and electron arrangement of a phosphide ion.
4. Strontium is in Group 2 and fluorine in Group 7. Give the formula of strontium fluoride.
5. Why do all Group 1 metals form ions with a single positive charge?
6. Name the type of bonding in solid magnesium bromide, MgBr₂.
7. How many Na⁺ ions surround each Cl⁻ ion in sodium chloride?
8. How many shared pairs and non-bonding pairs are there in a hydrogen chloride molecule?
9. Name the two giant covalent structures listed in the specification.
10. In solid potassium, what particles are in the lattice, and what holds them together?
11. Disulfur dichloride, S₂Cl₂, contains only non-metals. What does that tell you about its structure?

Answers

1. No. Both atoms are chlorine, so it is an element (made of molecules), not a compound.
2. [2]²⁺ (beryllium 2,2 loses its two outer electrons).
3. 3−; [2,8,8]³⁻ (atom 2,8,5 gains three electrons; 15 protons, 18 electrons).
4. SrF₂ (one 2+ ion balances two 1− ions).
5. Each atom has one electron in its outer shell, and losing it leaves a noble gas arrangement.
6. Ionic: magnesium is a metal and bromine a non-metal; Mg²⁺ and Br⁻ ions.
7. Six.
8. One shared pair; three non-bonding pairs, all on the chlorine.
9. Diamond and silicon dioxide.
10. K⁺ ions in a regular pattern, with delocalised electrons between them; strong electrostatic attraction between the ions and the electrons holds them together.
11. It is made of S₂Cl₂ molecules, with the atoms in each molecule joined by covalent bonds.

Where marks are usually lost

• Describing an ionic compound as made of molecules, or as one ion bonded to one other ion.
• Writing “ionic bonds are the attraction between atoms”. The attraction is between oppositely charged ions.
• Leaving out “in all directions” when describing the forces in an ionic lattice.
• Missing square brackets or charges on ions in a dot and cross diagram.
• Drawing a metal ion that still has its outer electrons, or a non-metal ion with seven rather than eight outer electrons.
• Using too few metal atoms, so electrons lost and gained do not balance (for example drawing one lithium with one sulfur).
• Forgetting the non-bonding pairs on nitrogen, oxygen or chlorine in covalent diagrams.
• Showing a double bond in O₂ as a single shared pair.
• Saying metals contain “free ions” or “a sea of protons”. The delocalised particles are electrons.
• Calling silicon dioxide a simple molecular substance because it is a compound of two non-metals.

Official syllabus

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