GCSE Chemistry Bonding revision

Everything in this chapter answers one question — what holds the particles together, and how strongly? The answer depends on which elements meet. When a metal meets a non-metal, the metal hands its outer electrons over: ionic bonding. When two non-metals meet, neither can afford to lose electrons, so they share: covalent bonding. When metal atoms gather together, they release their outer electrons into a shared pool: metallic bonding. Elements bond to reach stable outer shells (usually eight electrons, like a noble gas) — but the deeper truth is electrostatics: positive charges attract negative charges, and the three bond types are three ways of arranging that attraction.

A metal atom has few outer electrons and loses them readily; a non-metal atom is short of a full shell and grabs electrons. Bring them together and the transfer happens: sodium (2,8,1) loses its lone outer electron to become Na⁺ (2,8); chlorine (2,8,7) accepts it and becomes Cl⁻ (2,8,8). The compound formed contains ions — charged particles — not atoms. Magnesium gives up two electrons to become Mg²⁺; oxygen takes two on to become O²⁻.

The formula follows from the charges automatically: the compound must be electrically neutral, so positive and negative charges cancel. Two singly charged chloride ions balance one doubly charged calcium ion, giving CaCl₂. Two aluminium ions (+3 each, +6 total) balance three oxide ions (−2 each, −6 total), giving Al₂O₃.

Dot-and-cross diagrams show the transfer: draw the outer shells of both atoms, dots for one element’s electrons and crosses for the other’s, then close each resulting ion in square brackets with its charge written outside, top right. Every ion in the diagram should show a full outer shell — that is the whole point of the transfer.

Ions do not pair off into neat little molecules. Instead each ion surrounds itself with oppositely charged neighbours, and the pattern repeats in every direction: a giant ionic lattice. In sodium chloride, every Na⁺ touches six Cl⁻ and every Cl⁻ touches six Na⁺, each pair pulling electrostatically. Melting the crystal does not break bonds between pairs of ions — it has to overcome enormous numbers of strong attractions running through the whole lattice, which is why sodium chloride melts at 801 °C and magnesium oxide (charges of 2+ and 2− attracting) holds out until 2852 °C.

Conductivity reads the lattice like a meter. In the solid, ions are locked in place — vibrating, but unable to travel — so a solid ionic crystal does not conduct. Melt it, or dissolve it in water, and the ions break free: push one end of the liquid with a voltage and the ions march — cations to the cathode, anions to the anode — carrying current. Conduction happens only when the ions can move.

When two non-metals meet, neither will give electrons away, so they share. Two hydrogen atoms each supply an electron to a shared pair, and both atoms count the pair as their own — a single covalent bond. Oxygen atoms share two pairs (a double bond, O₂); nitrogen shares three (a triple bond, N₂). In methane, carbon shares one pair with each of four hydrogen atoms, completing carbon’s octet and each hydrogen’s pair. The molecules stay electrically neutral: nothing transferred, just pooled.

In dot-and-cross diagrams for molecules, overlap circles so the shared pair sits in the overlap, and show lone pairs — the outer electrons not used in bonding. Water’s two lone pairs, ammonia’s one and the halogens’ three are examined as often as the bonds themselves.

Inside a molecule of water or methane the covalent bonds are strong — but the molecules only touch each other through intermolecular forces, and those are feeble. Heat a molecular solid and the molecules need only a whisper of energy to pull apart from one another, so molecular substances have low melting and boiling points — water boils at 100 °C, methane at −162 °C. The vital subtlety, examined constantly: when a molecular substance melts or boils, the covalent bonds do not break. Steam is still H₂O — the O–H bonds inside each molecule are as strong as ever. Only the attractions between molecules are overcome.

Intermolecular forces grow with molecule size. Methane (CH₄) is a gas; octane (C₈H₁₈) is a liquid boiling at 126 °C; candle wax (chains of 20+ carbons) is a solid. Every extra atom adds surface for the weak attractions to act over, so bigger molecules cling together more and boil higher. This single trend powers a third of the crude oil chapter.

Because molecules are neutral and their electrons are locked into bonds, simple molecular substances do not conduct electricity in any state — there are no charged particles free to move.

A polymer is a simple molecular substance on a monstrous scale: chains of thousands of atoms, every neighbour-to-neighbour link a strong covalent bond. The forces between the chains are still intermolecular, but with so much chain in contact they add up to something solid: nearly all polymers are solid at room temperature and soften over a range rather than melting sharply. Stretch poly(ethene) and the chains uncoil; heat it and they slide — which is why polymer properties depend on chain length and side branches as much as on the monomer.

A giant covalent structure abandons molecules entirely. In diamond, every carbon atom bonds covalently to four neighbours, and the pattern fills the crystal: to melt it you must break covalent bonds themselves, and those are the strongest forces chemistry knows. Hence diamond’s astronomical melting behaviour and its place at the top of the hardness league. (The carbon allotropes — diamond, graphite, graphene, fullerenes — get their own chapter next.)

In a metal, every atom surrenders its outer electrons to a common sea of delocalised electrons that belongs to the whole crystal. What remains — positive metal ions — sits in orderly layers, electrostatically glued to the surrounding electron sea from every side. This arrangement explains the classic metal property list. The mobile electrons carry charge through the crystal, so metals conduct in solid and liquid states alike; they also carry heat, bouncing energy along the electron sea. The strong, all-directional attraction gives most metals high melting points (mercury’s −39 °C is the famous exception).

Malleability comes from the layers. Tap a metal and whole planes of ions slide one row over, into a new set of identical positions — the electron sea re-glues them instantly, so the metal dents rather than shatters. Now the alloy trick: add atoms of a different size — carbon into iron for steel, copper into gold for jewellery — and the orderly layers buckle. Rows can no longer slide smoothly, so alloys are harder and stiffer than the pure metals they improve.

Each bonding model trades truth for clarity. Dot-and-cross diagrams imply electrons are owned — dots belonging to one atom, crosses to another — when in reality no electron carries an owner’s stamp; they also show outer shells only, in flat circles, not real 3D orbitals. Ball-and-stick models make bonds look like solid rods and push atoms further apart than they truly are, while suggesting bond strength lives in the stick. 2D lattice drawings flatten a three-dimensional crystal; space-filling models hide the spaces entirely. Displayed formulae show every bond but nothing of the molecule’s shape. None of this makes the models useless — it means every conclusion drawn from them must respect what they exaggerate and omit.

• Dot-and-cross: electron “ownership” is imaginary; shells are drawn 2D.
• Ball-and-stick: sticks exaggerate bond length and pretend bonds are physical objects.
• Flat lattice drawings: hide the 3D repetition and the true packing.
• Space-filling: swallows the gaps, so conduction and diffusion become hard to picture.