What you'll learn
Chemical bonding explains how atoms join together, and structure explains how those bonded particles are arranged in a substance. Together, they account for why salt dissolves in water, why diamond is hard, why metals conduct and why some substances melt at room temperature while others survive a furnace. This guide covers ionic bonding, covalent bonding including dative (coordinate) bonds, metallic bonding, bond polarity and electronegativity, sigma (σ) and pi (π) bonds and hybridisation, and the four main types of structure — giant ionic, simple molecular, giant covalent (macromolecular) and giant metallic — with the physical properties each produces. Caribbean examples include the aluminium oxide refined from Jamaican bauxite and the silicon dioxide of beach sand.
In CAPE, bonding is examined both directly (dot-and-cross diagrams, explaining properties) and indirectly in almost every other topic. You must be able to link the type of bonding and structure to properties in clear cause-and-effect explanations.
Key terms and definitions
Ionic bond — The electrostatic attraction between oppositely charged ions, formed by the transfer of electrons.
Covalent bond — A shared pair of electrons between two atoms.
Dative (coordinate) bond — A covalent bond in which both shared electrons come from the same atom.
Metallic bond — The electrostatic attraction between a lattice of positive metal ions and a sea of delocalised electrons.
Electronegativity — The ability of an atom to attract the bonding pair of electrons in a covalent bond.
Polar bond — A covalent bond with an unequal distribution of electron density because of an electronegativity difference.
Sigma (σ) bond — A covalent bond formed by end-on (head-on) overlap of orbitals, with electron density along the line between nuclei.
Pi (π) bond — A covalent bond formed by sideways overlap of p orbitals, with electron density above and below the line between nuclei.
Hybridisation — The mixing of atomic orbitals to form new, equivalent hybrid orbitals.
Lattice — A regular, repeating three-dimensional arrangement of particles.
Polarisation — The distortion of an anion's electron cloud by a cation.
Core concepts
Ionic bonding
Ionic bonds form between metals and non-metals when electrons are transferred so that both atoms achieve stable (often noble-gas) configurations. Sodium loses one electron to form Na⁺; chlorine gains it to form Cl⁻. Dot-and-cross diagrams show the transfer and the resulting ions, with charges and square brackets.
The strength of ionic bonding increases with higher ionic charge and smaller ionic radius, because both increase the electrostatic attraction. Magnesium oxide (Mg²⁺ O²⁻) therefore has a much higher melting point than sodium chloride (Na⁺ Cl⁻).
Covalent bonding
Covalent bonds form between non-metal atoms that share pairs of electrons. Atoms may share one pair (single bond), two pairs (double bond) or three pairs (triple bond). In methane, carbon shares four pairs with four hydrogen atoms; in carbon dioxide, carbon forms two double bonds; in nitrogen, a triple bond holds the two atoms together, explaining nitrogen's low reactivity.
Some atoms can have expanded octets using d orbitals — sulfur in SF₆ has twelve electrons around it. Others have incomplete octets — boron in BF₃ has only six.
Dative bonding
In a dative bond, one atom provides both electrons. Examples:
- the ammonium ion, NH₄⁺, where nitrogen's lone pair forms a dative bond to H⁺;
- the hydroxonium ion, H₃O⁺;
- aluminium chloride dimer, Al₂Cl₆, where chlorine lone pairs donate to electron-deficient aluminium;
- complex ions in transition metal chemistry, where ligands donate lone pairs to a central metal ion.
Once formed, a dative bond is indistinguishable from other covalent bonds in the same species.
Bond polarity and electronegativity
When two atoms of different electronegativity are bonded, the bonding electrons are pulled towards the more electronegative atom, which gains a partial negative charge (δ−), leaving the other with a partial positive charge (δ+). The larger the difference, the more polar the bond. A very large difference leads to ionic bonding, so bonding can be seen as a continuum from pure covalent through polar covalent to ionic.
Ionic compounds can also show covalent character when a small, highly charged cation polarises a large anion. This is why aluminium chloride behaves partly covalently.
Sigma and pi bonds; hybridisation
A single bond is one σ bond. A double bond is one σ and one π bond. A triple bond is one σ and two π bonds. π bonds are weaker than σ bonds and restrict rotation around the bond, leading to geometric isomerism in alkenes.
Hybridisation explains shapes in organic molecules:
- sp³ — carbon in methane forms four equivalent hybrid orbitals, giving a tetrahedral shape (bond angle 109.5°).
- sp² — carbon in ethene forms three hybrid orbitals in a plane (120°), with an unhybridised p orbital forming the π bond.
- sp — carbon in ethyne forms two hybrid orbitals in a line (180°), with two π bonds.
Metallic bonding
Metal atoms release their outer electrons into a delocalised "sea", leaving positive ions in a lattice. The attraction between the ions and the delocalised electrons is the metallic bond. Strength increases with more delocalised electrons per atom and smaller ionic radius, so aluminium has a higher melting point than magnesium, which is higher than sodium.
Structure and properties
| Structure | Examples | Melting point | Electrical conductivity | Solubility |
|---|---|---|---|---|
| Giant ionic | NaCl, MgO, Al₂O₃ | High — strong ionic bonds throughout the lattice | Only when molten or dissolved — ions free to move | Often soluble in water |
| Simple molecular | H₂O, CO₂, I₂ | Low — weak intermolecular forces between molecules | Do not conduct — no free charged particles | Varies; many dissolve in non-polar solvents |
| Giant covalent | diamond, graphite, SiO₂ | Very high — many strong covalent bonds to break | Do not conduct (except graphite) | Insoluble |
| Giant metallic | Na, Cu, Fe | Usually high | Conduct when solid and liquid — delocalised electrons | Insoluble (some react) |
Diamond: each carbon forms four covalent bonds in a tetrahedral network — extremely hard, very high melting point, non-conductor.
Graphite: each carbon forms three bonds in hexagonal layers; the fourth electron is delocalised, so graphite conducts along the layers. Weak forces between layers let them slide, making graphite soft and useful as a lubricant and in pencils.
Silicon dioxide: each silicon is bonded to four oxygens and each oxygen to two silicons in a giant network. Beach sand and quartz are SiO₂, which explains their hardness and high melting point; the glass industry melts sand at very high temperatures.
A Caribbean example: alumina
Jamaica's bauxite is refined into aluminium oxide (alumina), a giant ionic structure with a very high melting point (above 2000 °C). This is why aluminium cannot be extracted economically by electrolysing molten pure alumina; it is first dissolved in molten cryolite, which lowers the operating temperature.
Worked examples
Example 1: Explaining melting points
Explain why magnesium oxide has a much higher melting point than sodium chloride. (3 marks)
Both are giant ionic lattices (1). Mg²⁺ and O²⁻ have double the charges of Na⁺ and Cl⁻, and Mg²⁺ is smaller than Na⁺ (1), so the electrostatic attraction between ions is much stronger and more energy is needed to break the lattice (1).
Example 2: Dative bonding
Draw and describe the bonding in the ammonium ion.
Nitrogen in NH₃ has three N–H covalent bonds and one lone pair. The lone pair is donated to an H⁺ ion, forming a dative bond. NH₄⁺ has four N–H bonds, all identical once formed, and a tetrahedral shape.
Example 3: Graphite conductivity
Explain why graphite conducts electricity but diamond does not.
In graphite, each carbon atom forms three covalent bonds, leaving one electron per atom delocalised between the layers, which can move and carry charge. In diamond, all four outer electrons are used in covalent bonds, so there are no free electrons.
Example 4: σ and π bonds
How many σ and π bonds are present in a molecule of ethene, C₂H₄?
There are four C–H single bonds (4 σ) and one C=C double bond (1 σ + 1 π). Total: 5 σ and 1 π.
Example 5: Comparing metallic bonding
Explain why the melting points increase from sodium (98 °C) to magnesium (650 °C) to aluminium (660 °C).
All three have giant metallic structures. From Na to Mg to Al, the number of delocalised electrons per atom increases from one to two to three, and the ionic radius decreases while the ionic charge increases (Na⁺, Mg²⁺, Al³⁺). The electrostatic attraction between the positive ions and the sea of delocalised electrons therefore becomes stronger, so more energy is needed to overcome it.
Example 6: Covalent character in ionic compounds
Explain why aluminium chloride shows significant covalent character but sodium chloride does not.
Al³⁺ is a small ion with a high charge, so it has a high charge density. It strongly polarises the large, easily distorted chloride ion, pulling its electron density towards the aluminium so that electrons are shared — giving covalent character. Na⁺ is larger with a single charge, so it polarises chloride much less, and NaCl remains essentially ionic.
Example 7: Describing a dot-and-cross diagram
Describe the bonding in carbon dioxide and in nitrogen.
In CO₂, the carbon atom shares two pairs of electrons with each oxygen atom, forming two C=O double bonds; each oxygen also has two lone pairs. Each double bond consists of one σ and one π bond. In N₂, the two nitrogen atoms share three pairs of electrons (one σ and two π bonds), and each nitrogen has one lone pair. The strong triple bond (bond energy about 945 kJ mol⁻¹) explains why nitrogen is very unreactive — and why the Haber process needs a catalyst and high temperature.
Example 8: Predicting properties from structure
Substance X melts at 1710 °C, does not conduct electricity when solid or molten, and is insoluble in water. Suggest its structure and an example.
The very high melting point and lack of conductivity in any state suggest a giant covalent structure, such as silicon dioxide (quartz or sand).
Properties by structure type: identifying an unknown
| Observation | Likely structure |
|---|---|
| High melting point, conducts only when molten or dissolved | giant ionic |
| High melting point, conducts when solid and liquid, malleable | giant metallic |
| Very high melting point, non-conductor (or conductor if graphite) | giant covalent |
| Low melting/boiling point, non-conductor | simple molecular |
Use this table to work backwards from data to structure in exam questions.
Common mistakes and how to avoid them
- Saying ionic compounds conduct when solid. Ions must be free to move.
- Saying simple molecular substances have low melting points because covalent bonds are weak. The covalent bonds are strong; it is the intermolecular forces that are weak.
- Confusing graphite's delocalised electrons with metallic bonding.
- Drawing dot-and-cross diagrams without charges on ions.
- Forgetting that a double bond is σ + π, not two σ bonds.
How this links to your Internal Assessment
You can investigate structure and bonding practically by testing the melting points, solubility and electrical conductivity of different substances — for example, comparing solid and molten salts, sugar solutions and metals. A well-designed SBA controls variables such as sample size and voltage, records observations systematically, and explains results using bonding and structure. Safety matters: molten salts and electrical tests must be handled carefully.
Exam technique for Chemical bonding and structure
- Always link property → structure → bonding → explanation.
- In dot-and-cross diagrams, show outer electrons only unless asked, with ion charges and brackets.
- When comparing melting points, name what must be broken (ionic bonds, covalent bonds, intermolecular forces, metallic bonds).
- Use precise terms: delocalised electrons, electrostatic attraction, giant lattice.
Quick revision summary
- Ionic: electron transfer, electrostatic attraction; stronger with higher charge and smaller ions.
- Covalent: shared pairs; single, double and triple bonds; expanded and incomplete octets.
- Dative: both electrons from one atom (NH₄⁺, H₃O⁺, Al₂Cl₆, complexes).
- Polarity from electronegativity differences; polarisation gives ionic compounds covalent character.
- σ bonds (end-on overlap) and π bonds (sideways overlap); sp³, sp², sp hybridisation.
- Metallic: positive ions in a sea of delocalised electrons.
- Giant ionic, simple molecular, giant covalent and giant metallic structures explain melting point, conductivity and solubility.