S2.2Structure 2 — Models of bonding and structure

The covalent model — IB Chemistry HL Revision Notes & Topical Questions

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S2.2 The covalent model — IB CHEMISTRY HL notes: Describe covalent bonding, Lewis formulas, shapes, polarity, giant structures, intermolecular forces and properties, and the HL ideas of resonance, expanded octets, formal charge and hybridisation.Covalent bond = the electrostatic …

Between non-metals
Normal bond
Octet rule
Electron deficient
Not a Lewis formula
Single, double, triple

Part 1 of 10Covalent bonds and the octet rule

Covalent bond

Covalent bond the electrostatic attraction between a shared pair of electrons and the positively charged nuclei of two atoms.

Between non-metals

Between non-metals no electrons are transferred; two atomic orbitals overlap to form a molecular orbital.

Normal bond

Normal bond each atom contributes one electron to the shared pair, and both reach a noble gas configuration.

Stability

Stability electrons are more stable shared between two nuclei than held by one atom.

Octet rule

Octet rule atoms tend to gain a valence shell of 8 electrons.

Electron deficient

Be in BeCl₂ has 4 electrons, and B in BF₃ and Al have 6.

Expanded octet

Expanded octet some period 3 central atoms hold more than 8 electrons.

Part 2 of 10Lewis formulas

Lewis formula

Lewis formula shows all valence electrons, as bonding pairs and lone pairs, using dots, crosses or lines.

Not a Lewis formula

Cl–Cl, because it does not show the lone pairs.

Count the total valence electrons of all atoms, adjusting for any charge.

Draw the skeleton with the least electronegative atom in the centre.

Add a bonding pair between each pair of bonded atoms.

Add lone pairs to complete octets (duets for hydrogen), starting with the outer atoms.

Check that all electrons are used and every atom has a full shell.

Part 3 of 10Multiple and coordinate bonds

Single, double, triple

Single, double, triple they share 2, 4 and 6 electrons.

No quadruple bonds

No quadruple bonds repulsion between 8 electrons between the nuclei is too great.

Bond energy

Bond energy the energy to break one mole of a covalent bond in the gas phase, in kJ mol⁻¹;
- a higher value means a stronger bond.

Bond length

Bond length the distance between the nuclei; more shared electrons pull the nuclei closer.

Trend

Trend triple bonds are shortest and strongest, single bonds longest and weakest.

Coordinate bond

Coordinate bond both electrons of the shared pair come from the same atom.

It forms when one atom has a lone pair and the other is electron deficient.
NH₃ donates its lone pair to H⁺ to form NH₄⁺.
At HL, [Al(H₂O)₆]³⁺ has six coordinate bonds from water oxygen atoms.

Part 4 of 10Shapes of molecules (VSEPR)

1Principle

Principle electron domains repel and spread as far apart as possible.
- (bonding pairs, lone pairs, or a multiple bond counted once)

2Repulsion order

Repulsion order lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair.

3Lone pairs

Lone pairs they are closer to the nucleus, so they repel more and squeeze bond angles.

4Lewis first

Lewis first draw the Lewis formula first so hidden lone pairs are not missed.

5Domain geometry versus molecular shape

Domain geometry versus molecular shape they are the same only when there are no lone pairs.

Bonding	Lone	Molecular shape	Angle	Example
2	0	Linear	180°	BeCl₂, CO₂
3	0	Trigonal planar	120°	BF₃
2	1	Bent	about 118°	SO₂
4	0	Tetrahedral	109.5°	CH₄
3	1	Trigonal pyramidal	107°	NH₃
2	2	Bent	104.5°	H₂O

Part 5 of 10Bond and molecular polarity

Electronegativity

Electronegativity the ability of an atom to attract the shared pair;
- values are on the Pauling scale in the data booklet.

Nonpolar bond

Nonpolar bond identical atoms share electrons equally, as in H₂ and Cl₂.

Polar bond

Polar bond the more electronegative atom is δ− and the other δ+;
- a larger electronegativity difference gives a more polar bond.

Dipole arrow

Dipole arrow the crossed tail is at the δ+ atom and the head points to the δ− atom.

Nonpolar molecule

Nonpolar molecule only nonpolar bonds, or polar bonds in a symmetrical arrangement so dipoles cancel, as in CO₂ and CCl₄.

Polar molecule

Polar molecule polar bonds arranged asymmetrically so dipoles do not cancel, as in H₂O, NH₃ and CH₃Cl.

Part 6 of 10Giant covalent structures

Giant covalent structures

Giant covalent structure

Giant covalent structure strong covalent bonds extend throughout the lattice, so there are no discrete molecules.

Silicon and SiO₂

Silicon and SiO₂ tetrahedral networks; SiO₂ has each Si bonded to four O and each O to two Si.

Diamond

Diamond each C bonded to four others, very hard, and non-conducting because all outer electrons are bonded.

Graphite

Graphite each C bonded to three others in flat hexagonal layers;
- delocalised electrons conduct and weak forces between layers make it soft.

Graphene

Graphene a single layer of graphite, one atom thick, strong, flexible and conducting.

C60

C60 a molecule, not a giant structure; a semiconductor.

Properties

Properties very high melting points, mostly insoluble, and usually non-conducting except graphite and graphene.

Part 7 of 10Intermolecular forces

Intermolecular forces

Intermolecular versus covalent

Intermolecular versus covalent forces between molecules are much weaker than covalent bonds within molecules.

London (dispersion) forces

London (dispersion) forces present in all substances;
- strength rises with the number of electrons and with surface area.

Dipole–dipole

Dipole–dipole between permanent dipoles, in addition to London forces.

Dipole–induced dipole

Dipole–induced dipole a polar molecule distorts the electron cloud of a nonpolar one.

Hydrogen bonding

Hydrogen bonding the strongest;
- needs H bonded to O, N or F and a lone pair on O, N or F of a neighbouring molecule.

Order

London < dipole–dipole < hydrogen bonding.

Van der Waals

Van der Waals collective name for London, dipole–dipole and dipole–induced dipole forces.

Part 8 of 10Physical properties and chromatography

Boiling point

Boiling point stronger intermolecular forces give higher melting and boiling points;
- weaker forces give volatile substances.

Solubility

Solubility like dissolves like: polar in polar solvents, nonpolar in nonpolar solvents;
- large hydrocarbon parts reduce solubility in water.

Conductivity

Conductivity molecular substances do not conduct, though some polar ones ionise in water, such as HCl.

Chromatography

Chromatography separates by different solubility in the mobile phase and attraction to the stationary phase;
- paper and thin-layer chromatography are required.

Rf distance moved by the component divided by the distance moved by the solvent front, with no units and always less than 1.

Part 9 of 10HL: resonance, expanded octets and formal charge

Resonance

Resonance more than one valid Lewis structure exists, and the real structure is a hybrid with delocalised electrons.

Examples

NO₃⁻, CO₃²⁻, O₃, RCOO⁻ and benzene all have equal bond lengths between single and double.

Benzene evidence

Benzene evidence hydrogenation enthalpy of −208 instead of −360 kJ mol⁻¹, all C–C bonds 140 pm, no bromine water decolourisation, no C=C infrared peak.

Expanded octet

Expanded octet period 3 and below can hold 10 or 12 electrons, as in PCl₅, SF₆, SF₄ and XeF₄.

Formal charge

FC = V − ½B − N;
- the preferred structure has formal charges closest to zero, with any negative on the most electronegative atom.

Part 10 of 10HL: sigma bonds, pi bonds and hybridisation

Sigma bond

Sigma bond head-on overlap, present in all single bonds and in every multiple bond.

Pi bond

Pi bond sideways overlap of p orbitals, present only in double and triple bonds.

Double bond

Double bond one σ and one π; triple bond — one σ and two π.

sp3

sp3 four domains, tetrahedral, 109.5°.

sp2

sp2 three domains, trigonal planar, 120°, with one π bond.

sp two domains, linear, 180°, with two π bonds.

Hybridisation

Hybridisation mixing of s and p orbitals to give equal hybrid orbitals, explaining equal bonds in methane.

Key terms

Term	Meaning
Covalent bond	Attraction between a shared electron pair and two nuclei.
Electron domain	A bonding pair, lone pair or multiple bond region around an atom.
Electronegativity	Ability of an atom to attract the shared pair in a bond.
Hydrogen bond	Strong intermolecular attraction involving H bonded to O, N or F.
Resonance hybrid	The real structure formed by delocalised electrons.
Sigma bond	Head-on overlap bond.
Pi bond	Sideways overlap bond.
Hybridisation	Mixing of atomic orbitals to form equal hybrid orbitals.

Avoid these mistakes

Confusing electron domain geometry with molecular shape — lone pairs change the shape.

Saying a molecule with polar bonds must be polar — symmetrical molecules such as CO₂ and CCl₄ are nonpolar.

Breaking covalent bonds when a molecular substance melts — only intermolecular forces are overcome.

Counting a double bond as two domains in VSEPR — a multiple bond is one domain.

Giving an R_f value with units or a value above 1.

Last updated: 18 October 2026