OCR GCSE Chemistry A: Bonding (J248) – Revision Notes

For full explanations and worked examples, read the Bonding study guide first. These notes condense sub-topic C2.2 Bonding (C2.2a to C2.2i and CM2.2i to CM2.2iii) of the OCR GCSE (9–1) Chemistry A (Gateway Science) J248 specification, Version 4.0 (August 2026), for first assessment in 2018. Nothing in C2.2 is restricted to the Higher Tier: Foundation and Higher candidates need all of it. Within the written papers, C2 content is assessed on Paper 1 and Paper 3, and assumed on Papers 2 and 4.

When you are ready, try the Bonding practice questions. Course links: OCR GCSE Chemistry hub, J248 checklist, and the free 10-minute diagnostics.

Key definitions

• Ion: an atom or group of atoms with a charge, because electrons have been lost or gained. Protons never change.
• Ionic bond: the electrostatic pull holding positive and negative ions together, acting in all directions.
• Covalent bond: an electron pair held between two non-metal atoms, with both nuclei pulling on it.
• Metallic bond: positive metal ions in a regular lattice pulled towards the delocalised electrons that surround them.
• Lone pair: a pair of outer electrons on one atom that is not shared.
• Group number (1 to 7): the number of outer-shell electrons. Group 0: full outer shell.
• Period number: the number of occupied electron shells.

Key facts and numbers

Idea	What to remember
Shell capacities (first 20 elements)	2, then 8, then 8, then the rest
Atomic radius and bond length	Order of 10⁻¹⁰ m (0.1 nm)
Nanoparticle size	Roughly 1 to 100 nm
1 nm	1 × 10⁻⁹ m
Ion charge from group	Groups 1, 2, 3: +1, +2, +3. Groups 5, 6, 7: −3, −2, −1
Formula of an ionic compound	Choose ion numbers so the total charge is zero
Atoms across a particle	Particle width ÷ atom width (same units)

Metals vs non-metals (C2.2a, C2.2b)

	Metals	Non-metals
Position	Left and centre of the table	Right of the table
Outer electrons	Few (usually 1 to 3)	4 to 7, or a full shell in group 0
In reactions	Lose electrons, form positive ions	Gain electrons (negative ions) or share them
Physical properties	Shiny, good conductors, malleable, mostly high melting points	Dull, poor conductors, brittle, mostly low melting points
Oxides	Basic; neutralise acids; soluble ones are alkaline	Acidic; soluble ones give acidic solutions

Example equations: 4Li + O₂ → 2Li₂O (a basic oxide) and C + O₂ → CO₂ (an acidic oxide).

Method: electron arrangement to position (C2.2c)

1. Atomic number = number of electrons in the neutral atom.
2. Fill shells: 2, 8, 8, then 2.
3. Count shells for the period. Count outer electrons for the group.

Example: magnesium, Z = 12 → 2,8,2 → period 3, group 2.

Bonding types side by side (C2.2d, C2.2e)

	Ionic	Simple molecular	Giant covalent	Polymer	Metallic
Elements involved	Metal + non-metal	Non-metals	Non-metals (e.g. C, Si with O)	Mostly C and H	Metal(s)
Particles	Ions	Molecules	Atoms	Long-chain molecules	Ions + delocalised electrons
Bond	Ionic, all directions	Covalent within molecules	Covalent throughout	Covalent along chains	Metallic
Between units	Same ionic bonds, whole lattice	Weak intermolecular forces	None in diamond (one network); weak forces between graphite layers	Weaker forces between chains	Same metallic bonds

All chemical bonds are electrostatic. Ionic bonds come from transfer of electrons, and covalent bonds from sharing.

Method: dot and cross for a binary ionic compound (C2.2f)

1. Write both electron arrangements.
2. Work out the electron loss per metal atom and the electron gain per non-metal atom.
3. Work out the ratio of atoms so electrons lost = electrons gained.
4. Put brackets round every ion and write its charge outside, top right. Use crosses for the metal’s electrons and dots for the non-metal’s.
5. Check that the charges add to zero.

Reminder: in calcium nitride, each Ca gives up 2 electrons and each N takes 3. The lowest common number is 6, so 3 Ca²⁺ and 2 N³⁻ give Ca₃N₂.

Method: dot and cross for a simple molecule (C2.2f)

1. Write the outer electrons for each atom.
2. Work out how many each atom needs to reach a full shell (hydrogen needs 2 in total).
3. Share pairs so every atom reaches its full shell. Use double or triple bonds where single bonds are not enough.
4. Add lone pairs.
5. Check: (shared pairs × 2) + (lone pairs × 2) = total outer electrons.

Reminder: dichlorine oxide, Cl₂O, has 7 + 6 + 7 = 20 outer electrons. Oxygen sits in the middle with one shared pair to each chlorine. Oxygen keeps 2 lone pairs and each chlorine 3, so 2 shared + 8 lone = 10 pairs.

Method: formula from a model or diagram (CM2.2ii, CM2.2iii)

1. Count how many bonds each type of atom makes in the picture.
2. Total bonds from one side = total bonds from the other side.
3. Simplify the ratio of atoms.

For a ball and stick model of a simple molecule, just count each colour of ball. Example: a model with 3 black (C) and 8 white (H) balls is C₃H₈ (propane). It has 2 C–C and 8 C–H bonds.

Limitations of models (C2.2g)

Model	Shows	Does not show
Dot and cross	Which atom electrons came from; how they pair	3D shape; that electrons are identical; relative sizes
Ball and stick	3D shape; bond angles	True sizes; that bonds are not rods; for ionic lattices, that ions attract all neighbours
2D displayed formula	Which atoms are joined	Shape and bond angles
3D drawing (wedges)	An idea of depth	Still a flat projection; sizes not to scale

Reactions and the table (C2.2h, C2.2i)

• Same group, same number of outer electrons, so similar reactions.
• Group 1: reactivity increases down the group. Extra shells put the outer electron further out and screen it from the nucleus, so it is removed more easily.
• Group 7: reactivity decreases down the group. The electron to be gained is attracted less strongly as the atom gets bigger.
• Group 0: full outer shells, so very unreactive.
• Mendeleev ordered by atomic weight, left gaps and predicted missing elements, and swapped some pairs to match properties.
• Ordering by atomic number (proton number) explains those swaps. The modern table is in atomic number order.

Must-know distinctions

• Ion vs atom: same protons, different electrons.
• Ionic lattice vs molecule: sodium chloride is a lattice, not a molecule.
• Covalent bond vs intermolecular force: the first joins atoms within one molecule and is strong; the second acts between neighbouring molecules and is weak.
• Group vs period: group tells outer electrons; period tells shells.
• Atomic weight vs atomic number: Mendeleev used the first; the modern table uses the second.

Quick self-test

1. Give the electron arrangement, group and period of sulfur (Z = 16).
2. An element has the arrangement 2,8,8,1. Name it and say whether it is a metal.
3. How many protons and electrons are in a sulfide ion, S²⁻?
4. Deduce the formula of aluminium sulfide.
5. How many bonding pairs and lone pairs are on the sulfur atom in H₂S?
6. How many lone pairs are on the nitrogen atom in nitrogen trifluoride, NF₃?
7. An oxide dissolves in water and the solution turns universal indicator red. Was the element a metal or a non-metal?
8. What did Mendeleev do when atomic weight order would have put an element into a family with the wrong properties?
9. A nanoparticle is 30 nm across. Its atoms are 0.3 nm across. Estimate the number of atoms across it.
10. State something a ball and stick model gets wrong about an ionic lattice.
11. What type of bond joins the repeating units in a polymer chain?
12. Explain why potassium reacts more vigorously than sodium.

Answers

1. 2,8,6; group 6; period 3.
2. Potassium; yes, a metal (group 1).
3. 16 protons, 18 electrons.
4. Al³⁺ and S²⁻: 2 × (+3) = +6 and 3 × (−2) = −6, so Al₂S₃.
5. 2 bonding pairs and 2 lone pairs.
6. 1 lone pair (each F also has 3).
7. A non-metal: its oxide is acidic.
8. He swapped the pair so each element sat with elements of similar properties. Ordering by atomic number later explained why the swap was right.
9. 30 ÷ 0.3 = 100 atoms.
10. The sticks suggest each ion is bonded to particular ions; really each ion attracts all its neighbours in every direction. (Also: ions are not to scale, and gaps are exaggerated.)
11. Covalent bonds.
12. In potassium the single outer electron occupies the fourth shell, not the third, so it is further from the nucleus with more inner shells screening it. The weaker pull lets it leave more readily.

Where marks are usually lost

• Saying a positive ion has gained protons, or that the nucleus changes.
• Ionic dot and cross diagrams with no square brackets, no charges, or electrons still shown on the metal’s old outer shell.
• Charges that do not balance in the formula, for example writing AlS for aluminium sulfide.
• Placing lone pairs in the overlap region, or forgetting them entirely.
• Drawing hydrogen with 8 outer electrons instead of 2.
• Describing sodium chloride as made of molecules.
• Giving period number as the number of outer electrons.
• Explaining group 1 reactivity by “more electrons” without linking distance from the nucleus to weaker attraction.
• Saying Mendeleev ordered elements by atomic number.

Official syllabus

OCR GCSE (9–1) Chemistry A (Gateway Science) J248 specification, Version 4.0 (August 2026), for first assessment in 2018, published by OCR. Sub-topic C2.2 Bonding.