Bonding, structure and the properties of matter.

HookSame atom, three different materials

ModelWhy atoms bond, and the three ways they do it

Atoms bond because a full outer shell of electrons is a stable arrangement, and most atoms do not start with one. There are three ways to reach it, and which one operates depends on the elements involved. When a metal reacts with a non-metal, electrons transfer from the metal to the non-metal, producing charged ions held together by ionic bonding. When non-metals bond with each other, neither will give electrons up, so they share pairs of electrons in covalent bonding. When metal atoms bond among themselves, they pool their outer electrons into a shared 'sea', giving metallic bonding.

That gives you a fast prediction rule straight from the periodic table: metal plus non-metal is ionic, non-metal plus non-metal is covalent, and metal with metal (or a pure metal) is metallic. Sodium chloride is a metal with a non-metal, so it is ionic; methane is carbon with hydrogen, both non-metals, so it is covalent; magnesium is a metal, so it is metallic. Getting the bond type right is the first move in almost every C2 answer, because the bonding decides the structure and the structure decides the properties. Everything downstream in this topic — melting point, conductivity, hardness — is read off from this first decision.

MechanismIonic bonding and giant lattices

In ionic bonding, a metal atom loses its outer electron or electrons to become a positive ion, and a non-metal atom gains them to become a negative ion. Sodium loses one electron to become Na⁺; chlorine gains it to become Cl⁻; the oppositely charged ions attract by strong electrostatic forces. Dot-and-cross diagrams show this transfer, drawing one atom's electrons as dots and the other's as crosses so you can follow where each electron goes, usually showing only the outer shells and the resulting charges in brackets.

The ions do not exist as isolated pairs. They pack into a giant ionic lattice — a regular, repeating three-dimensional arrangement in which every positive ion is surrounded by negative ions and vice versa, held by strong electrostatic forces acting in all directions. A single grain of salt is one enormous lattice of billions of alternating Na⁺ and Cl⁻ ions. The formula, NaCl, gives the simplest whole-number ratio of ions, not the number in a molecule — there are no ionic molecules.

MechanismCovalent and metallic bonding

Covalent bonding is the sharing of a pair of electrons between two non-metal atoms, and the shared pair — counting towards both atoms' outer shells — holds them together strongly. A hydrogen molecule shares one pair; oxygen shares two pairs (a double bond); methane's carbon shares a pair with each of four hydrogens. Covalent bonding builds three kinds of substance: small molecules such as water, carbon dioxide and chlorine; giant covalent structures such as diamond and silicon dioxide, where covalent bonds extend unbroken through the whole solid; and polymers, very long molecules made of repeating units. Dot-and-cross diagrams for molecules show only the shared and outer electrons.

Metallic bonding is different again. In a metal, the atoms lose their outer electrons into a shared pool, leaving a regular lattice of positive ions surrounded by a 'sea' of delocalised electrons that are free to move throughout the structure. The bond is the strong electrostatic attraction between the positive ions and this sea of negative electrons. Because the electrons are delocalised — no longer tied to one atom — they can drift, and that single feature, as later sections show, is what makes metals conduct electricity and heat so well. Three bond types, three pictures: transfer, share, and pool.

ModelStates of matter, the particle model and state symbols

The simple particle model represents every substance as small spheres. In a solid the particles are packed close in a regular pattern, vibrating on the spot, which is why solids hold their shape and volume. In a liquid they still touch but can move past one another, so a liquid flows and takes the shape of its container. In a gas they are far apart and move quickly in random directions, so a gas spreads to fill any space. Changes of state — melting, freezing, boiling, condensing — are physical changes: no new substance forms and the change can be reversed. Whether a substance is solid, liquid or gas at a given temperature depends on how strong the forces between its particles are; the stronger the forces, the more energy is needed and the higher the melting and boiling points.

The model is useful but limited, and examiners reward knowing its limitations: it treats particles as solid, inelastic spheres with nothing between them, it ignores the forces between particles, and it ignores that real particles are not simple spheres and vary in size. Finally, state symbols record the state in equations: (s) for solid, (l) for liquid, (g) for gas and (aq) for aqueous — dissolved in water. Writing NaCl(s) dissolving to give Na⁺(aq) + Cl⁻(aq) tells a reader the salt has gone from a solid lattice to free ions in solution.

CaseReading properties off structure — ionic, molecular, polymer

Now read properties off structure. Ionic compounds have high melting and boiling points because the giant lattice is held by strong electrostatic forces between oppositely charged ions, and a large amount of energy is needed to overcome them all — sodium chloride melts at 801 °C. They do not conduct electricity when solid, because the ions are locked in place, but they do conduct when molten or dissolved in water, because the ions are then free to move and carry charge. That 'free to move' phrase is the whole explanation.

Small molecules are the opposite. They have low melting and boiling points — many are liquids or gases at room temperature — because although the covalent bonds inside each molecule are strong, the intermolecular forces between separate molecules are weak, and it is only these weak forces that break when the substance melts or boils. The larger the molecules, the stronger the intermolecular forces and the higher the melting point, which is why longer hydrocarbons are more viscous. Small molecules do not conduct electricity, because they have no overall charge and no free electrons or ions. Polymers sit in between: very large molecules whose atoms are joined by strong covalent bonds into long chains, with intermolecular forces between the chains large enough that polymers are solid at room temperature, though they soften more easily than giant structures.

CaseGiant covalent structures and the carbon allotropes

Giant covalent structures, also called macromolecules, are single lattices in which every atom is joined to its neighbours by strong covalent bonds, so melting one means breaking a vast number of strong bonds. That gives them very high melting points. Diamond and silicon dioxide are examples, and carbon's several forms — its allotropes — show how much structure alone can do. In diamond, each carbon atom forms four covalent bonds to other carbons in a rigid three-dimensional network, making it extremely hard with a very high melting point; because all four outer electrons are used in bonding, none are free, so diamond does not conduct electricity.

Graphite is the same element arranged completely differently. Each carbon bonds to only three others, forming flat layers of hexagons, and the fourth outer electron of every atom is delocalised. Those delocalised electrons let graphite conduct electricity and heat, like a metal, while the weak forces between layers let the layers slide, making graphite soft and a good lubricant. Peel graphite down to a single layer and you have graphene — one atom thick, extremely strong and highly conductive, useful in electronics and composites. Fullerenes are hollow carbon structures shaped as spheres or tubes: buckminsterfullerene (C₆₀) is a football-shaped molecule, and carbon nanotubes are cylinders with very high tensile strength and useful electrical properties, used in nanotechnology, as catalysts and to deliver drugs into the body.

MechanismMetals and alloys — malleability and conduction

Metals and alloys close the topic, and both come straight from metallic bonding. A pure metal is a regular arrangement of layers of atoms, and because the layers are even they can slide over one another when a force is applied — which is why pure metals are malleable, able to be hammered into shape, and often too soft for use. An alloy is a mixture of a metal with one or more other elements whose atoms are a different size; those different-sized atoms distort the regular layers so they can no longer slide easily, making the alloy harder than the pure metal. This is why we use steel rather than pure iron, and bronze or brass rather than pure copper.

The delocalised electrons explain why metals conduct. Because these electrons are free to move through the whole structure, they carry electrical charge through the metal when a voltage is applied, making metals good electrical conductors; the same mobile electrons also transfer thermal energy quickly, so metals are good thermal conductors. Metallic bonding — the strong attraction between the positive ions and the sea of electrons — also gives most metals high melting and boiling points. So one feature, delocalised electrons free to move, accounts for conduction, while the arrangement in sliding layers accounts for malleability: structure and bonding, once again, doing all the explaining.

VocabularyKey terms the mark scheme pays for

TrapsMisconceptions that cost marks

ExamWhat examiners want