S2.2Structure 2 — Models of bonding and structure

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

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S2.2 The covalent model — IB CHEMISTRY SL notes: Draw Lewis formulas, deduce shapes and polarity, describe giant covalent structures, link intermolecular forces to properties and interpret chromatography.Covalent bond = the electrostatic attraction between a shared pair of electr…

Occurrence
Orbital overlap
Octet rule
Exceptions
Multiple bonds
Representations

Part 1 of 10Covalent bonding

Covalent bond

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

Occurrence

Occurrence between two non-metals, with electrons shared rather than transferred.

Orbital overlap

Orbital overlap two atomic orbitals overlap to form a molecular orbital, and sharing is more stable than holding electrons on one atom.

Octet rule

Octet rule atoms share electrons to gain a noble-gas valence shell of eight electrons.
- (a duet for hydrogen)

Exceptions

Be in BeCl₂ has 4 electrons around it and B in BF₃ has 6, and these molecules are stable.

Multiple bonds

Multiple bonds single, double and triple bonds share 2, 4 and 6 electrons, and quadruple bonds are impossible because repulsion is too great.

Part 2 of 10Drawing Lewis formulas

Count the total valence electrons, adding one per negative charge and subtracting one per positive charge.

Draw the skeleton with the least electronegative atom central (never hydrogen), joined by single bonds.

Complete the octets of outer atoms with lone pairs, then place any remaining electrons on the central atom.

If the central atom lacks an octet, convert lone pairs on outer atoms into double or triple bonds.

Check that all valence electrons are used and every atom has an octet (a duet for H).

Lewis formula

Lewis formula shows all valence electrons as bonding pairs and lone pairs, so Cl–Cl alone is not a Lewis formula.

Representations

Representations a pair can be dots, crosses, a mixture of both, or a line.

Part 3 of 10Bond length and strength; coordinate bonds

Bond strength

Bond strength triple > double > single, because more shared electrons pull the nuclei closer.

Bond length

Bond length single > double > triple, e.g. C–C 154 pm, C=C 134 pm and C≡C 120 pm.

Bond energy

Bond energy the energy needed to break one mole of a bond in the gaseous state, in kJ mol⁻¹.

Coordinate bond

Coordinate bond a covalent bond in which both electrons come from the same atom.

Formation

Formation a lone pair is donated into an empty orbital, e.g. NH₃ + H⁺ → NH₄⁺.
Identical bonds — once formed it is the same as other covalent bonds, so all four N–H bonds in NH₄⁺ are identical.
Electron-deficient acceptor — BF₃ can accept a lone pair, e.g. forming F₃B–NH₃.

Part 4 of 10VSEPR shapes

Principle

Principle electron domains repel and spread as far apart as possible.
- (bonding pairs, lone pairs, each multiple bond counting as one)

Repulsion order

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

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

1Domain geometry versus molecular shape

Domain geometry versus molecular shape they can differ, e.g. NH₃ has tetrahedral domains but a trigonal pyramidal shape.

2Lone pairs

Lone pairs each reduces the bond angle by about 2–2.5°.

Part 5 of 10Bond polarity and molecular polarity

1Electronegativity

Electronegativity the ability of an atom to attract the shared pair in a bond, on the Pauling scale (F is highest at 4.0).

2Polar bond

Polar bond unequal sharing gives δ+ on the less electronegative atom and δ− on the other, and a larger difference means a more polar bond.

3Non-polar bond

Non-polar bond identical atoms such as H₂ and Cl₂ share equally.

4Non-polar molecule with polar bonds

Non-polar molecule with polar bonds symmetrical geometry cancels the dipoles, as in CO₂ and CCl₄.

5Polar molecule

Polar molecule asymmetrical geometry leaves a net dipole, as in H₂O, NH₃ and CHCl₃.

6Rule

Rule a molecule with only non-polar bonds cannot be polar.

Part 6 of 10Giant covalent structures

Substance	Structure	Conductivity	Properties
Diamond	Each C bonded tetrahedrally to 4 others	None	Very hard, very high melting point
Graphite	Each C bonded to 3 in hexagonal layers; layers held by London forces	Conducts along layers (one delocalised electron per atom)	Soft and slippery
Graphene	A single layer of graphite	Excellent	Strong, flexible and one atom thick
C₆₀ fullerene	Discrete spherical molecules with 20 hexagons and 12 pentagons	Poor (semiconductor)	Low melting point; not a giant structure
Silicon	Diamond-like network	Poor	High melting point
Silicon dioxide	Each Si bonded to 4 O and each O to 2 Si, formula SiO₂	None	High melting point, hard, insoluble

Melting points
- covalent bonds, not intermolecular forces, must break in network solids.

Part 7 of 10Intermolecular forces

1Intermolecular forces

Intermolecular forces act between molecules and are much weaker than covalent bonds, which are intramolecular.

2London (dispersion) forces

London (dispersion) forces arise from instantaneous dipoles inducing dipoles in neighbours, are present in all substances and grow with more electrons and greater surface area.

3Dipole-induced dipole

Dipole-induced dipole a permanent dipole distorts the electron cloud of a non-polar neighbour.

4Dipole-dipole

Dipole-dipole attraction between the δ+ and δ− ends of polar molecules.

5Hydrogen bonding

H bonded to N, O or F is attracted to a lone pair on N, O or F in a neighbouring molecule, the strongest of these forces.

6Strength order

London < dipole-dipole < hydrogen bonding for similar molar mass.

7Branching

Branching straight-chain alkanes have higher boiling points than branched isomers.

8Hydrogen bonding examples

Hydrogen bonding examples the anomalously high boiling points of H₂O, HF and NH₃, and ice being less dense than water.

Part 8 of 10Properties of molecular substances

Properties of molecular substances

Melting and boiling points

Melting and boiling points depend on the strength of intermolecular forces, and low boiling points mean volatile.

Solubility

"like dissolves like", so polar and hydrogen-bonding solutes dissolve in water and non-polar solutes in non-polar solvents such as hexane.

Conductivity

Conductivity molecular substances have no mobile charged particles, so they do not conduct, except graphite and graphene.

Polar molecules that ionise

Polar molecules that ionise such as HCl conduct in solution because ions form.

Small alcohols

Small alcohols ethanol dissolves in water, but solubility falls as the hydrocarbon chain lengthens.

Part 9 of 10Paper chromatography and R<sub>f</sub>

Draw a pencil baseline about 1 cm from the bottom of the paper.

Spot the samples on the baseline and let them dry.

Stand the paper in solvent with the level below the baseline and cover to saturate the atmosphere.

Remove the paper when the solvent is near the top and mark the solvent front immediately.

Locate colourless spots with a UV lamp or a locating agent such as ninhydrin.

Measure distances from the baseline and calculate R_f.

Rf distance moved by the spot ÷ distance moved by the solvent front, both from the baseline.

Properties of Rf

Properties of Rf it has no unit and lies between 0 and 1.

Interpretation

Interpretation components more soluble in the solvent and less attracted to the stationary phase travel further and have a higher R_f;
- compare with references run under identical conditions.

Part 10 of 10More from the specification

Key terms

Term	Meaning
Covalent bond	The attraction between a shared pair of electrons and the two nuclei.
Coordinate bond	A covalent bond in which both electrons come from the same atom.
VSEPR theory	Electron domains around a central atom repel to be as far apart as possible, with lone pairs repelling most.
Electronegativity	The ability of an atom to attract the shared pair in a covalent bond.
Allotropes	Different structural forms of the same element in the same state.
London (dispersion) forces	Weak attractions between instantaneous and induced dipoles.
Hydrogen bond	Attraction between H bonded to N, O or F and a lone pair on N, O or F of another molecule.

Avoid these mistakes

Saying covalent molecules melt by breaking covalent bonds — only intermolecular forces are overcome (except in network solids).

Assuming polar bonds always make a polar molecule — symmetrical shapes such as CO₂ cancel the dipoles.

Forgetting lone pairs when predicting shape — draw the Lewis formula first.

Quoting R_f with units or values above 1 — it is a ratio less than 1.

Last updated: 14 October 2026