Chemical bonding is the single topic that underpins the largest proportion of A Level Chemistry. Structure and bonding explains melting points, solubility, conductivity, reactivity, and the physical state of every substance you will encounter in the course. Students who understand bonding at the level of electron behaviour — not just at the level of definition — can answer any properties question they encounter, including ones about substances they have never seen.
This guide covers the three primary bonding types from first principles: why each type forms, what the resulting structure looks like, and exactly how the structure explains every physical property examiners test. An interactive substance explorer and a bonding properties quiz follow the content sections.
"Every physical property of a substance — its melting point, conductivity, solubility, hardness — is a direct consequence of the forces holding its particles together. Understand the forces and you can predict the properties without memorising a single substance."
— Fahad Rafiq, Chemistry & Biology TutorBefore distinguishing between bond types, it is essential to understand the concept that determines which type forms: electronegativity. Electronegativity is the tendency of an atom to attract the shared electrons in a bond toward itself, and it increases across periods (more protons, same shielding) and decreases down groups (more shielding, increased atomic radius).
When two atoms bond, what matters is the difference in their electronegativities. A large difference (typically >1.7 on the Pauling scale) means electrons are transferred completely — ionic bonding. A small-to-moderate difference between non-metals means electrons are shared but unequally — polar covalent bonding. Effectively zero difference between identical atoms means perfectly equal sharing — non-polar covalent bonding. Metals, with low electronegativities and loosely held outer electrons, form metallic bonding — a fundamentally different arrangement.
Examiners frequently ask you to predict bond type from electronegativity values, or to explain why a bond is polar. The answer always begins with electronegativity difference. A student who can state "the larger electronegativity difference between Na and Cl (2.1 on the Pauling scale) means electrons are transferred from Na to Cl, forming ions" earns marks that a student who simply states "ionic bonding occurs between metals and non-metals" cannot — because the latter is a rule, and the former is a mechanism.
Ionic bonding forms when the electronegativity difference between two atoms is large enough that the less electronegative atom (always a metal) loses one or more electrons entirely to the more electronegative atom (a non-metal). The result is two ions: a cation (positively charged, having lost electrons) and an anion (negatively charged, having gained electrons). These opposite charges attract by Coulomb's law — the electrostatic force — and this attraction is the ionic bond.
Critically, ionic bonding does not produce discrete molecules. The ions arrange themselves into a three-dimensional giant ionic lattice — an extended, repeating pattern of alternating positive and negative ions, each surrounded by six ions of the opposite charge (in the case of NaCl). The total lattice energy — the energy released when one mole of ionic compound forms from its gaseous ions — is enormous, typically 600–3000 kJ mol⁻¹, which directly explains the high melting points of ionic compounds.
"NaCl conducts electricity because it has ionic bonds." This is incomplete and earns no marks. Solid NaCl does not conduct electricity — the ions are fixed in the lattice. Conductivity requires ions that are free to move.
"Solid NaCl does not conduct electricity because the ions are held in fixed positions within the lattice. When molten or dissolved in water, the ions are free to move toward the electrodes — so conduction occurs."
At A Level, ionic bonding connects directly to the Born-Haber cycle, which uses Hess's law to calculate lattice energies indirectly. The magnitude of lattice energy depends on two factors: the charge on the ions (higher charge = stronger attraction = higher lattice energy) and the ionic radii (smaller ions = shorter interionic distance = stronger attraction). This is why MgO has a much higher lattice energy than NaCl — Mg²⁺ and O²⁻ carry double the charge of Na⁺ and Cl⁻, and are smaller.
Covalent bonds form when atoms share electron pairs rather than transferring them. The shared electrons experience attraction from both nuclei simultaneously — this mutual attraction is what holds the atoms together. The bond forms because the bonded state is lower in energy than the separated atoms. Covalent bonding almost exclusively involves non-metal atoms, where the electronegativity difference is insufficient for complete electron transfer.
The critical distinction for exam purposes is between simple molecular covalent substances and giant covalent structures. This distinction explains apparently contradictory properties — CO₂ is a covalent compound with a very low boiling point (−78°C), while diamond is also entirely covalent but has one of the highest melting points of any substance (over 3500°C). The explanation lies not in the covalent bonds themselves but in what holds the structures together.
In simple molecular substances (H₂O, CO₂, NH₃, CH₄, HCl), the covalent bonds within each molecule are strong. But the forces between molecules — intermolecular forces — are weak. When these substances are heated, the intermolecular forces break, not the covalent bonds. The result is low melting and boiling points. These substances do not conduct electricity because no ions or free electrons are present.
In giant covalent structures (diamond, graphite, silicon dioxide, silicon), every atom is covalently bonded to its neighbours in an extended three-dimensional network. There are no discrete molecules — the entire crystal is effectively one giant molecule. Melting requires breaking covalent bonds throughout the structure, which demands enormous energy. Hence these substances have extremely high melting points.
Graphite is a giant covalent structure where each carbon atom forms three covalent bonds with neighbours in flat hexagonal layers. The fourth electron from each carbon is delocalised across the layer — it is not involved in bonding to any particular atom. These delocalised electrons can move freely parallel to the layers, giving graphite electrical conductivity. This is a favourite examiner topic because it is both counterintuitive (a non-metal conducting electricity) and fully explainable from first principles.
When two atoms of different electronegativities share electrons, the bond is polar — the electron density is shifted toward the more electronegative atom, creating a partial negative charge (δ−) on that atom and a partial positive charge (δ+) on the other. Whether the molecule has an overall dipole (and hence is polar) depends on molecular geometry. CO₂ has two polar C=O bonds, but they point in opposite directions and cancel — the molecule is non-polar. H₂O has two polar O–H bonds that are angled at 104.5°, producing a net dipole. This explains why CO₂ dissolves poorly in water while H₂O is miscible with polar solvents.
Metallic bonding is qualitatively different from ionic and covalent bonding. In metals, the outer electrons — one, two, or three depending on the metal — are not associated with any particular atom. Instead, they are delocalised: free to move throughout the entire metal lattice. What remains is an array of positively charged metal cations embedded in this mobile "sea" of electrons. The bonding is the electrostatic attraction between the positive cation lattice and the delocalised electrons.
This model — the sea of electrons or electron sea model — directly and elegantly explains all the characteristic properties of metals. It is not a metaphor; it is a mechanistic description, and examiners expect you to use it as such when answering properties questions.
"Metals are good conductors because they have free electrons." This is correct but earns only one mark — it states the fact without the mechanism.
"Metals have a lattice of positive ions surrounded by delocalised electrons. When a potential difference is applied, these delocalised electrons move through the lattice toward the positive terminal, carrying charge — this constitutes an electric current."
Both are metals with the same basic bonding type. But Mg²⁺ ions have a higher charge than Na⁺, and Mg²⁺ is smaller — so the charge density of the cation is higher. Each Mg atom contributes two delocalised electrons rather than Na's one. The electrostatic attraction between Mg²⁺ ions and the denser electron sea is therefore much stronger. More energy is needed to overcome the metallic bonding — Mg melts at 650°C, Na at 98°C. This argument applies across Group 1 and Group 2 and explains the general trend of Group 2 having higher melting points than Group 1.
For simple molecular substances, physical properties are determined not by the covalent bonds within molecules but by the forces between molecules. These intermolecular forces are one to two orders of magnitude weaker than covalent bonds, which is why simple molecular substances have low melting and boiling points. But within the category of intermolecular forces, significant variation exists — and examiners exploit it constantly.
| Force | What causes it | Relative strength | Which substances |
|---|---|---|---|
| London dispersion (van der Waals) | Temporary dipoles from random electron movement; more electrons = stronger force | Weakest | All molecules — the only force in non-polar molecules (e.g. noble gases, alkanes) |
| Permanent dipole–dipole | Attraction between δ+ of one polar molecule and δ− of another | Medium | Polar molecules without O–H, N–H, or F–H bonds (e.g. HCl, propanone) |
| Hydrogen bonding | Electrostatic attraction between δ+ H (bonded to O, N, or F) and lone pair on O, N, or F of adjacent molecule | Strongest IMF | Molecules with O–H, N–H, or H–F bonds (e.g. H₂O, NH₃, HF, ethanol, DNA) |
The anomalous boiling point of water (100°C, far higher than expected for its molecular mass compared to H₂S at −60°C) is entirely explained by hydrogen bonding. Each water molecule can form up to four hydrogen bonds — two as donor (via O–H) and two as acceptor (via the two lone pairs on oxygen). This extensive hydrogen bond network also explains water's high surface tension, high specific heat capacity, and the fact that ice is less dense than liquid water — all consequences of the same intermolecular interaction.
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Select a substance to explore its bonding type, structure, and how the structure explains every physical property. These are the substances most commonly used in exam questions.
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Every physical property of every substance you encounter in A Level Chemistry can be explained by two things: the forces holding particles together, and the energy needed to overcome them. Master the three bonding types and the three intermolecular forces — know what determines their strength, what their structures look like, and how those structures determine each observable property — and every bonding question becomes a reasoning exercise rather than a memory exercise.
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