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Cambridge International AS Level Chemistry

Topic 3
Chemical Bonding

Topic 3 covers electronegativity, ionic, metallic and covalent bonding, the shapes of molecules, intermolecular forces and dot-and-cross diagrams. These pages cover electronegativity, ionic bonding and giant ionic lattices, metallic bonding, covalent bonding with sigma and pi bonds, dative bonds and giant covalent structures, the shapes and bond angles of molecules, polar molecules, and intermolecular forces with hydrogen bonding and the anomalous properties of water.

Exam Papers
Papers 1 & 2
9701/1 and 9701/2 (AS Level)
Learning Outcomes
3.1.1 – 3.7.1
7 parts
Topic Parts
6 of 7 live
3.1 to 3.6
Exam Board
Cambridge
CIE 9701 (2025-2027)

Revision Notes

Work through each part of Cambridge Topic 3 in a structured sequence.

Electronegativity
Electro-
negativity
1

3.1 Electronegativity

What electronegativity is, the factors and trends, and using Pauling values to predict bond type.

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Ionic
Ionic
Bonding
2

3.2 Ionic Bonding

Ionic bonding in sodium chloride, magnesium oxide and calcium fluoride, giant ionic lattices, polarisation and the properties of ionic compounds.

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Metallic
Metallic
Bonding
3

3.3 Metallic Bonding

Metal ions in a sea of delocalised electrons, the giant metallic lattice and how it explains the properties of metals.

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Covalent
Covalent
Bonding
4

3.4 Covalent Bonding

The covalent bond, sigma and pi bonds, dative bonds, expanded octets, giant covalent structures and polar molecules.

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VSEPR
Shapes of
Molecules
5

3.5 Shapes of Molecules

Shapes and bond angles of BF3, CO2, CH4, NH3, H2O, SF6 and PF5, and predicting analogous molecules and ions.

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Forces
Intermolecular
Forces
6
Available

3.6 Intermolecular Forces

Van der Waals’ forces, id-id and pd-pd forces, hydrogen bonding, the anomalous properties of water including surface tension, boiling point trends and solvents.

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Specification Coverage

Topic 3 Chemical Bonding – Cambridge International AS Level Chemistry

The following Cambridge learning outcomes state what candidates should be able to do for Topic 3. Wording is taken from the Cambridge International AS & A Level Chemistry 9701 syllabus for 2025 to 2027.

3.1 Electronegativity and bonding

3.1.1
define electronegativity as the power of an atom to attract electrons to itself
3.1.2
explain the factors influencing the electronegativities of the elements in terms of nuclear charge, atomic radius and shielding by inner shells and sub-shells
3.1.3
state and explain the trends in electronegativity across a period and down a group of the Periodic Table
3.1.4
use the differences in Pauling electronegativity values to predict the formation of ionic and covalent bonds (the presence of covalent character in some ionic compounds will not be assessed) (Pauling electronegativity values will be given where necessary)

3.2 Ionic bonding

3.2.1
define ionic bonding as the electrostatic attraction between oppositely charged ions (positively charged cations and negatively charged anions)
3.2.2
describe ionic bonding including the examples of sodium chloride, magnesium oxide and calcium fluoride

3.3 Metallic bonding

3.3.1
define metallic bonding as the electrostatic attraction between positive metal ions and delocalised electrons

3.4 Covalent bonding and coordinate (dative covalent) bonding

3.4.1
define covalent bonding as electrostatic attraction between the nuclei of two atoms and a shared pair of electrons:
(a)
describe covalent bonding in molecules including: • hydrogen, H2 • oxygen, O2 • nitrogen, N2 • chlorine, Cl2 • hydrogen chloride, HCl • carbon dioxide, CO2 • ammonia, NH3 • methane, CH4 • ethane, C2H6 • ethene, C2H4
(b)
understand that elements in period 3 can expand their octet including in the compounds sulfur dioxide, SO2, phosphorus pentachloride, PCl5, and sulfur hexafluoride, SF6
(c)
describe coordinate (dative covalent) bonding, including in the reaction between ammonia and hydrogen chloride gases to form the ammonium ion, NH4+, and in the Al2Cl6 molecule
3.4.2
(a)
describe covalent bonds in terms of orbital overlap giving σ and π bonds: σ bonds are formed by direct overlap of orbitals between the bonding atoms; π bonds are formed by the sideways overlap of adjacent p orbitals above and below the σ bond
(b)
describe how the σ and π bonds form in molecules including H2, C2H6, C2H4, HCN and N2
(c)
use the concept of hybridisation to describe sp, sp2 and sp3 orbitals
3.4.3
(a)
define the terms: • bond energy as the energy required to break one mole of a particular covalent bond in the gaseous state • bond length as the internuclear distance of two covalently bonded atoms
(b)
use bond energy values and the concept of bond length to compare the reactivity of covalent molecules

3.5 Shapes of molecules

3.5.1
state and explain the shapes of, and bond angles in, molecules by using VSEPR theory, including as simple examples: • BF3 (trigonal planar, 120°) • CO2 (linear, 180°) • CH4 (tetrahedral, 109.5°) • NH3 (pyramidal, 107°) • H2O (non-linear, 104.5°) • SF6 (octahedral, 90°) • PF5 (trigonal bipyramidal, 120° and 90°)
3.5.2
predict the shapes of, and bond angles in, molecules and ions analogous to those specified in 3.5.1

3.6 Intermolecular forces, electronegativity and bond properties

3.6.1
(a)
describe hydrogen bonding, limited to molecules containing N–H and O–H groups, including ammonia and water as simple examples
(b)
use the concept of hydrogen bonding to explain the anomalous properties of H2O (ice and water): its relatively high melting and boiling points; its relatively high surface tension; the density of the solid ice compared with the liquid water
3.6.2
use the concept of electronegativity to explain bond polarity and dipole moments of molecules
3.6.3
(a)
describe van der Waals’ forces as the intermolecular forces between molecular entities other than those due to bond formation, and use the term van der Waals’ forces as a generic term to describe all intermolecular forces
(b)
describe the types of van der Waals’ forces: instantaneous dipole–induced dipole (id–id) forces, also called London dispersion forces; permanent dipole–permanent dipole (pd–pd) forces, including hydrogen bonding
(c)
describe hydrogen bonding and understand that hydrogen bonding is a special case of permanent dipole–permanent dipole forces between molecules where hydrogen is bonded to a highly electronegative atom
3.6.4
state that, in general, ionic, covalent and metallic bonding are stronger than intermolecular forces

3.7 Dot-and-cross diagrams

3.7.1
use dot-and-cross diagrams to illustrate ionic, covalent and coordinate bonding including the representation of any compounds stated in 3.4 and 3.5 (dot-and-cross diagrams may include species with atoms which have an expanded octet or species with an odd number of electrons)