What are equal repulsion for all electron pairs?

Lone pairs repel more, so each lone pair lowers the bond angle by about 2.5^.

EnglandChemistrySyllabus dot point

How does electron configuration explain bonding, molecular shape and the properties of materials?

Electron configuration in shells, sub-shells and orbitals, ionic, covalent (including dative) and metallic bonding, electronegativity and bond polarity, electron-pair repulsion and molecular shapes, and the properties of the four crystal structures.

An OCR H432 module 2 answer covering electron configuration in sub-shells and orbitals, ionic, covalent, dative and metallic bonding, electronegativity and polarity, electron-pair repulsion shapes, and crystal structures.

Generated by Claude Opus 4.811 min answerUpdated 2026-06-02

Reviewed by: AI editorial process; not yet individually human-reviewed

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What this topic is asking
Electron configuration
Bonding types
Electronegativity and polarity
Molecular shapes
Intermolecular forces and crystal structures
Examples in context
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What this topic is asking

OCR specification points 2.2.1 and 2.2.2 want you to write electron configurations using shells, sub-shells and orbitals, classify bonding (ionic, covalent including dative, and metallic), use electronegativity to judge bond polarity, predict molecular shapes from electron-pair repulsion, describe intermolecular forces, and link the four crystal structures to physical properties.

Electron configuration

Electrons occupy shells (principal energy levels), divided into sub-shells (s, p, d) made of orbitals, each holding up to two electrons of opposite spin. Sub-shells fill in order of increasing energy: $1\text{s}, 2\text{s}, 2\text{p}, 3\text{s}, 3\text{p}, 4\text{s}, 3\text{d}, 4\text{p}$ .

Bonding types

Electronegativity and polarity

Electronegativity is the ability of an atom to attract the bonding electrons in a covalent bond; it increases across a period and decreases down a group. A difference in electronegativity makes a bond polar, giving partial charges $\delta+$ and $\delta-$ .

Molecular shapes

Electron pairs around a central atom repel and move as far apart as possible (electron-pair repulsion theory). Lone pairs repel more strongly than bonding pairs, so they reduce bond angles.

Electron pairs (lone)	Shape	Bond angle	Example
2 (0)	linear	$180^{\circ}$	$\text{CO}_2$
3 (0)	trigonal planar	$120^{\circ}$	$\text{BF}_3$
4 (0)	tetrahedral	$109.5^{\circ}$	$\text{CH}_4$
4 (1)	pyramidal	$107^{\circ}$	$\text{NH}_3$
4 (2)	bent	$104.5^{\circ}$	$\text{H}_2\text{O}$
6 (0)	octahedral	$90^{\circ}$	$\text{SF}_6$

Intermolecular forces and crystal structures

The three intermolecular forces, weakest first, are London (induced dipole) forces, permanent dipole-dipole forces, and hydrogen bonding (when H is bonded to N, O or F). The four crystal structures explain bulk properties:

Ionic (e.g. $\text{NaCl}$ ): high melting point, conducts when molten or aqueous, often soluble in water.
Simple molecular (e.g. $\text{I}_2$ ): low melting point (weak intermolecular forces), does not conduct.
Giant covalent (e.g. diamond, graphite, $\text{SiO}_2$ ): very high melting point; graphite conducts because of delocalised electrons, diamond does not.
Metallic (e.g. $\text{Mg}$ ): high melting point, conducts in all states, malleable.

Explaining the anomalous boiling points of the hydrides

The boiling points of $\text{H}_2\text{O}$ , $\text{HF}$ and $\text{NH}_3$ are much higher than the trend of their groups predicts. Explain why.

step 1: Identify the dominant force

Each contains hydrogen bonded directly to a small, highly electronegative atom (O, F or N), so hydrogen bonding occurs.

step 2: Compare with neighbours

The heavier hydrides in the same groups (such as $\text{H}_2\text{S}$ and $\text{HCl}$ ) have only London and dipole-dipole forces, which are weaker.

step 3: Link energy to boiling point

Hydrogen bonds need more energy to break, so molecules that form them boil at a higher temperature than expected.

step 4: Conclude

The anomalously high boiling points of water, hydrogen fluoride and ammonia are direct evidence for hydrogen bonding, and explain why water is liquid at room temperature.

Examples in context

Example 1. Ice floating on water. Hydrogen bonds hold water molecules in an open hexagonal lattice when frozen, making ice less dense than liquid water so it floats. This unusual behaviour, vital for aquatic life in winter, is a direct consequence of the hydrogen bonding predicted from oxygen's high electronegativity.

Example 2. Graphite as a lubricant and electrode. Graphite's giant covalent layers, held together by weak London forces, slide over one another (lubrication) while delocalised electrons within each layer carry current (electrodes). The same bonding model explains two very different industrial uses, connecting structure directly to properties.

Try this

Q1. State the shape and bond angle of a methane molecule. [2 marks]

Cue. Tetrahedral, $109.5^{\circ}$ (four bonding pairs, no lone pairs).

Q2. Explain why magnesium has a higher melting point than sodium. [2 marks]

Cue. Magnesium ions are $2+$ and release two delocalised electrons per atom, giving stronger metallic bonding than sodium's $1+$ ions.

Exam-style practice questions

Practice questions written in the style of OCR exam questions on this dot point, with worked answer explainers. The year tag is the paper they imitate, not the source.

OCR 20195 marks(a) Write the full electron configuration of a chromium atom (atomic number 24). (b) State and explain the shape and bond angle of a molecule of ammonia,

\text{NH}_3

Show worked answer →

(a) Chromium is an exception: $1\text{s}^2 2\text{s}^2 2\text{p}^6 3\text{s}^2 3\text{p}^6 3\text{d}^5 4\text{s}^1$ (one $4\text{s}$ electron promoted to give a stable half-filled $3\text{d}$ ) (1).

(b) Ammonia is pyramidal with a bond angle of about $107^{\circ}$ (1 for shape, 1 for angle). There are three bonding pairs and one lone pair around nitrogen (1); the lone pair repels more strongly than bonding pairs, pushing the bonds closer together below the tetrahedral $109.5^{\circ}$ (1).

OCR 20214 marks(a) Explain why the boiling point of

\text{HF}

is higher than that of

\text{HCl}

, even though

\text{HCl}

has more electrons. (b) State whether

\text{CCl}_4

is a polar molecule and justify your answer.

Show worked answer →

(a) $\text{HF}$ molecules form hydrogen bonds because F is highly electronegative and bonded to H (1); these are stronger than the permanent dipole-dipole and London forces in $\text{HCl}$ , so more energy is needed to separate the molecules (1).

(b) $\text{CCl}_4$ is non-polar (1): although each C-Cl bond is polar, the molecule is tetrahedral and symmetrical, so the four bond dipoles cancel and there is no overall dipole (1).

Related dot points

Sources & how we know this

OCR A-Level Chemistry A (H432) specification — OCR (2015)