AQA A-Level Chemistry: Bonding (7405)

This guide covers 3.1.3 Bonding, part of the Physical Chemistry content in AQA A-level Chemistry (7405), examined at both AS and A-level. It builds on Section 3.1.1’s atomic structure and underpins almost every later topic in the specification, since bonding type and molecular shape explain physical properties, reactivity, and much of organic mechanism. These notes complement the site’s guides to Atomic Structure and Relative Mass, the Mole and the Avogadro Constant.

Where this fits in 7405

Sections 3.1.1 to 3.1.7 of Physical Chemistry are designed to be covered in the first year of the A-level and also form the AS content, so bonding is examined from the very first AS papers onwards. Understanding why substances have the melting points, boiling points and electrical conductivities they do depends directly on the bonding and structure covered here.

Syllabus coverage

AQA A-LEVEL CHEMISTRY (7405) – 3.1.3 BONDING

• Ionic bonding (3.1.3.1): electrostatic attraction between oppositely charged ions in a lattice; the formulas of compound ions (e.g. sulfate, hydroxide, nitrate, carbonate and ammonium); predicting the charge on a simple ion from the element’s position in the Periodic Table, and constructing formulas for ionic compounds
• Covalent and dative covalent bonds (3.1.3.2): a single covalent bond as a shared pair of electrons; multiple bonds as multiple shared pairs; a co-ordinate (dative covalent) bond as a shared pair with both electrons supplied by one atom; representing a covalent bond with a line and a co-ordinate bond with an arrow
• Metallic bonding (3.1.3.3): attraction between delocalised electrons and positive ions arranged in a lattice
• Bonding and physical properties (3.1.3.4): the four types of crystal structure – ionic, metallic, macromolecular (giant covalent) and molecular – with diamond, graphite, ice, iodine, magnesium and sodium chloride as the named examples; relating melting point and conductivity to the type of structure and the bonding present; the energy changes associated with changes of state; drawing diagrams of these structures with specified numbers of particles
• The shapes of simple molecules and ions (3.1.3.5): bonding and lone pairs as charge clouds that repel each other and arrange themselves as far apart as possible, with lone pair-lone pair repulsion greater than lone pair-bond pair repulsion, which is greater than bond pair-bond pair repulsion; explaining the shapes of, and bond angles in, simple molecules and ions with up to six electron pairs around the central atom
• Bond polarity (3.1.3.6): electronegativity as the power of an atom to attract the pair of electrons in a covalent bond; an unsymmetrical electron distribution between atoms of different electronegativity producing a polar covalent bond, which may give a molecule a permanent dipole; using partial charges to show a bond is polar; explaining why some molecules with polar bonds have no permanent dipole
• Forces between molecules (3.1.3.7): permanent dipole-dipole forces, induced dipole-dipole (van der Waals, dispersion, London) forces and hydrogen bonding; how the strength of these forces influences the melting and boiling points of molecular substances; the importance of hydrogen bonding in the low density of ice and the anomalous boiling points of compounds

The specification’s skills-development column also suggests identifying the structure of an unknown substance by experiment (for example by testing its solubility, conductivity and ease of melting); that is a practical opportunity, not a separate content statement.

How to approach it

This section is best learned as a single connected argument rather than four separate bonding types: bonding type determines structure, structure determines the forces that must be overcome to melt or boil a substance, and those forces explain the physical properties examiners actually ask about. Practise relating a substance’s melting point and electrical conductivity to its type of crystal structure and the bonding present – that explanatory skill, not just definitions, is what higher-mark questions test.

Official syllabus

AQA A-level Chemistry (7405) specification, first teaching 2015 – aqa.org.uk.

Three bonding types, four types of crystal structure

Ionic bonding produces a giant lattice of oppositely charged ions, held together by strong electrostatic attraction in every direction – this is why ionic compounds have high melting points and conduct electricity only when molten or dissolved (the ions must be free to move). Covalent bonding shares electron pairs between specific atoms, producing either small discrete molecules (low melting points, no conductivity) or, in a few cases, giant covalent structures like diamond or graphite (very high melting points, and conductivity only for graphite, due to its delocalised electrons). Metallic bonding produces a lattice of positive ions in a “sea” of delocalised electrons, explaining both the electrical conductivity of metals (electrons are free to move) and their malleability (the lattice can deform without breaking fixed bonds).

Together these give the specification’s four types of crystal structure, each with named examples: ionic (sodium chloride), metallic (magnesium), macromolecular or giant covalent (diamond and graphite) and molecular (iodine and ice). Melting a molecular crystal such as iodine or ice overcomes only the relatively weak forces between molecules, not the covalent bonds inside them, which is why molecular substances have much lower melting points than macromolecular ones.

Shapes of molecules: electron-pair repulsion in three steps

Step 1 - count the total number of electron pairs around the
central atom:
Both bonding pairs and lone pairs count.

Step 2 - arrange the pairs to minimise repulsion:
The greater the repulsion type, the further apart those pairs
push each other -- lone pair-lone pair repulsion is the
strongest.

Step 3 - read off the shape and bond angle:
A lone pair takes up more space than a bonding pair, so each
lone pair present in a molecule compresses the bond angles
between the remaining bonding pairs.

This is why methane (4 bonding pairs, no lone pairs) is a perfect tetrahedron at 109.5 degrees, while ammonia (3 bonding pairs, 1 lone pair) is a trigonal pyramid with a compressed 107 degree angle, and water (2 bonding pairs, 2 lone pairs) is bent at approximately 104.5 degrees – the same starting electron-pair count, systematically compressed by each additional lone pair.

Worked example: predicting intermolecular forces

Given a molecule, ask three questions in order to identify every relevant intermolecular force:

Q1 - Does it contain O-H, N-H or H-F bonds?
     Yes -> hydrogen bonding is present (the strongest of the
     three intermolecular forces).

Q2 - Is the molecule polar overall (asymmetric, with polar bonds)?
     Yes -> permanent dipole-dipole forces are present.

Q3 - Does it contain any electrons at all?
     Yes -> induced dipole (van der Waals) forces are ALWAYS
     present, in every molecule, regardless of the answers above.

Van der Waals forces are the one force present in literally every substance; hydrogen bonding and dipole-dipole forces are additional to, not instead of, van der Waals forces.

Polar bonds, but no dipole. Q2 asks about the molecule overall for a reason. In a symmetrical molecule such as carbon dioxide (linear) or tetrachloromethane, CCl4 (tetrahedral), each bond is polar but the bond dipoles point in opposite or symmetrical directions and cancel, so the molecule has no permanent dipole and only van der Waals forces act between its molecules.

Hydrogen bonding and ice. In ice, each water molecule is hydrogen bonded to four others in an open, regular lattice that holds the molecules further apart than in liquid water, so ice is less dense than water and floats. Hydrogen bonding also explains the anomalously high boiling points of H2O, NH3 and HF compared with the hydrides of the other elements in their groups.

Common mistakes

Describing metallic bonding as “electrons shared between atoms” (that is covalent bonding language, not metallic). Forgetting that lone pairs, not just bonding pairs, must be counted before predicting a molecular shape. Stating a bond angle without explaining the lone-pair compression that produces it. Treating van der Waals forces as absent in polar molecules – they are always present alongside any stronger forces. Confusing electronegativity (an atom’s pull on shared electrons) with polarity (a property of the bond or molecule that results from an electronegativity difference).

Quick revision checklist

• Explain how ionic, covalent (including dative) and metallic bonding each produce a different kind of structure.
• Name the four types of crystal structure and the six example crystals, and relate melting point and conductivity to the type of structure and the bonding present.
• Write the formulas of the compound ions (sulfate, hydroxide, nitrate, carbonate, ammonium) and construct formulas for ionic compounds.
• Apply electron-pair repulsion theory to predict the shape and bond angle of a given molecule or ion.
• Identify which intermolecular forces are present in a given molecule, remembering van der Waals forces are always present.
• Distinguish electronegativity from bond/molecular polarity, and explain how one produces the other.
• Explain why a molecule such as CO2 has polar bonds but no permanent dipole.
• Explain how hydrogen bonding makes ice less dense than water.