OCR A Level Chemistry A: Bonding and structure (H432) – Revision Notes

These notes condense section 2.2.2, Bonding and structure, from OCR’s A Level Chemistry A specification (H432, Version 3.1 dated May 2026; first assessment 2017), statements (a) to (o). H432 is the full A level with no tiers, and Module 2 is assessed in all three written components. For full explanations and worked examples, read the bonding and structure study guide first.

When you are ready, use the practice questions. To see where this unit sits, open the course hub for OCR A Level Chemistry or the printable checklist; the diagnostics page flags gaps quickly.

Definitions to learn word for word

Term	Definition
Ionic bond (a)	Electrostatic attraction between positive and negative ions
Covalent bond (d)	Strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms
Dative covalent (coordinate) bond (e)	A covalent bond whose shared pair is supplied entirely by one of the two atoms
Average bond enthalpy (f)	Used as a measure of covalent bond strength: larger value, stronger bond (definition and calculations belong to 3.2.1)
Electronegativity (i)	An atom’s ability to attract the bonding electrons in a covalent bond
Permanent dipole (j)	Charge separation (δ+, δ−) in a bond between atoms of different electronegativity
Polar molecule (j)	A molecule with an overall dipole: polar bonds whose dipoles do not cancel because of the shape
Van der Waals’ forces (k)	Covers both permanent dipole–dipole and induced dipole–dipole interactions
London (dispersion) forces (k)	Another name for induced dipole–dipole interactions
Simple molecular lattice (n)	Covalently bonded molecules held in a regular arrangement by intermolecular forces

Ionic compounds: (a) to (c)

• Dot-and-cross: ions in square brackets, charge outside top right, outer shells only unless asked.
• Reminder: in calcium hydride, CaH₂, each Ca gives one electron to each of two H atoms. Each [H]⁻ has two electrons (one dot, one cross), filling its first shell.
• Structure (b): giant ionic lattice. In NaCl, each ion is surrounded by six of the opposite charge; attraction acts in all directions.

Property (c)	Result	Reason in one line
Melting/boiling point	High	Strong attractions between oppositely charged ions throughout the lattice need lots of energy to overcome
Solubility	Many dissolve in water	Polar water molecules attract and surround the ions
Solid conductivity	None	Ions fixed in place
Molten or aqueous	Conducts	Ions mobile

Covalent dot-and-cross: (d) and (e)

Method in steps:

1. Write the outer electron count for each atom.
2. Pair electrons into bonds so each atom reaches a stable arrangement (often eight; two for H; up to twelve round a Period 3 or later central atom).
3. Draw shared pairs in the overlap. Draw every lone pair.
4. For a dative bond, the donor’s lone pair becomes the shared pair.
5. For an ion, add or remove electrons for the charge and put the whole ion in brackets with its charge.

Small worked reminder: carbon monoxide, CO, has a triple bond. Two shared pairs have one electron from each atom; the third pair comes entirely from oxygen, so it is a dative bond, O→C. Each atom also keeps one lone pair.

Shapes and bond angles: (g) and (h)

Repulsion order: lone–lone > lone–bonding > bonding–bonding. Each lone pair cuts about 2.5° from the angle.

Shape	Bonding pairs, lone pairs	Angle
Linear	two, none	180°
Trigonal planar	three, none	120°
Tetrahedral	four, none	109.5° (CH₄)
Pyramidal	three, one	107° (NH₃)
Non-linear	two, two	104.5° (H₂O)
Octahedral	six, none	90°

Counting electrons round the central atom: outer electrons of the central atom, plus one per singly bonded atom, plus one per negative charge (minus one per positive charge). Halve for pairs. A double or triple bond counts as one region.

Small worked reminder: NCl₃ gives 5 + 3 = 8 electrons, 4 pairs, 3 bonding and 1 lone, so pyramidal, 107°.

3-D diagrams: line in plane, solid wedge towards you, dashed wedge away.

Electronegativity and polarity: (i) and (j)

• Increases towards F: across a period, up a group.
• Pauling scale: difference 0 means non-polar; small means polar covalent; large means mostly ionic.
• The more electronegative atom is δ−.
• Polar molecule test: (1) are the bonds polar? (2) does the shape leave the dipoles uncancelled?

Molecule	Polar bonds?	Shape	Overall dipole?
H₂O	Yes	Non-linear	Yes
CO₂	Yes	Linear	No (dipoles cancel)
BCl₃	Yes	Trigonal planar	No
CH₃Cl	Yes	Tetrahedral, unsymmetrical	Yes

Intermolecular forces: (k) to (o)

Force	Between	Origin	Strength depends on
Induced dipole–dipole (London)	All molecules	Instantaneous dipole induces a dipole in a neighbour	Number of electrons
Permanent dipole–dipole	Polar molecules	δ+ of one attracts δ− of another	Size of the dipole
Hydrogen bond	Molecules with H on N, O or F	Lone pair on N, O or F attracted to strongly δ+ H	Strongest of the three

Hydrogen-bond diagram checklist: δ+ on H and δ− on N/O/F, the lone pair drawn, a dashed line from the lone pair to the H, and a straight O–H···O line (about 180°).

Water (m)

• Ice: each H₂O hydrogen bonds to four others in a tetrahedral, open lattice, so ice is less dense than water; molecules move closer when it melts.
• Melting and boiling points higher than the trend in the other Group 16 hydrides predicts, because hydrogen bonds must be broken as well as London forces.

Simple molecular substances (n) and (o)

• Lattice examples: I₂ (London forces), ice (hydrogen bonds).
• Low melting and boiling points: only intermolecular forces break, not covalent bonds.
• Solubility: like dissolves like. Non-polar in non-polar solvents; hydrogen-bonding molecules in water.
• No conductivity in any state: no mobile charged particles.

Must-know distinctions

• Ionic lattice vs simple molecular lattice. Melting NaCl breaks ionic bonds; melting I₂ overcomes London forces only.
• Polar bond vs polar molecule. CO₂ and BCl₃ have polar bonds but are non-polar molecules.
• Electron-pair arrangement vs shape. NH₃ has four pairs arranged tetrahedrally, but its shape is pyramidal.
• Intermolecular vs intramolecular. Hydrogen bonds hold molecules together; O–H covalent bonds hold atoms together inside each molecule.
• More electrons vs heavier. London forces depend on electrons; mass only tracks the electron count.

Quick self-test

1. How many electrons are shown around each hydride ion in a dot-and-cross diagram of CaH₂?

2. State the shape and bond angle of BeF₂.

3. Which atom donates the dative pair in carbon monoxide?

4. XeF₄ has six pairs round Xe. How many are lone pairs, and what is the shape?

5. Why is SiCl₄ non-polar even though Si–Cl bonds are polar?

6. OF₂ and H₂O have the same shape. State it and the bond angle you would predict.

7. Br₂ and ICl each have 70 electrons. Use these boiling points: Br₂ 59 °C, ICl 97 °C. Explain the difference.

8. Why can HF molecules hydrogen bond to each other but HCl molecules cannot?

9. Which is stronger, a C≡C or a C–C bond, and what evidence would show it?

10. Explain why magnesium oxide conducts when molten but not when solid.

11. Why does iodine have a low melting point even though the I–I bond is covalent?

12. Describe how to draw the 3-D shape of CH₄.

Answers

1. Two: one dot and one cross, giving H⁻ a full first shell.

2. Linear, 180°: two bonding pairs, no lone pairs (2 + 2 = 4 electrons).

3. Oxygen. It supplies both electrons of the third shared pair (O→C).

4. 8 + 4 = 12 electrons, 6 pairs. Four bonding, so two lone pairs; square planar.

5. The molecule is tetrahedral and symmetrical, so the four identical dipoles cancel.

6. Non-linear, about 104.5°: two bonding and two lone pairs (6 + 2 = 8 electrons).

7. London forces are similar because the electron count is the same. ICl is polar (Cl is more electronegative than I), so it also has permanent dipole–dipole interactions; more energy is needed, so its boiling point is 38 °C higher.

8. F is electronegative enough to make the H strongly δ+, and F has lone pairs that the H can approach closely. Cl is less electronegative and larger, so hydrogen bonds do not form.

9. C≡C. It has the larger average bond enthalpy.

10. Solid: Mg²⁺ and O²⁻ ions are locked in position. Molten: the lattice has broken down and the ions can move to carry charge.

11. Melting only overcomes the weak London forces between I₂ molecules. The covalent bonds stay intact.

12. Two C–H bonds as lines in the plane of the paper, one as a solid wedge (towards you) and one as a dashed wedge (away), angle 109.5°.

Where marks are usually lost

• Leaving out “and the nuclei of the bonded atoms” from the covalent bond definition.
• Missing lone pairs on the central atom, especially on O, N, S and halogens.
• Forgetting square brackets and charges on ions, or drawing a covalent overlap for an ionic compound.
• Counting a double bond as two bonding regions and getting the wrong shape.
• Giving 109.5° for a pyramidal or non-linear molecule, or not explaining that lone pairs repel more.
• Calling a molecule polar from its bonds alone, without checking the shape.
• Writing “breaking bonds” for boiling a simple molecular substance without saying they are intermolecular forces.
• Drawing a hydrogen bond without the lone pair, or from a C–H hydrogen.
• Explaining ice floating by “molecules expand” instead of the open, hydrogen-bonded lattice.

Next steps

Try the practice questions, revisit the study guide wherever an answer slipped, then check off 2.2.2 on the checklist. The diagnostics page has short tests for other topics.

Official syllabus

OCR, A Level Chemistry A specification (H432), Version 3.1 dated May 2026, first assessment 2017: section 2.2.2 Bonding and structure.