Metallic bonding

Chapter 7 (u1:c07) explained compounds built by electron transfer — metals handing electrons to non-metals to make ionic lattices. Chapter 8 (u1:c08) explained compounds built by electron sharing — non-metals holding each other with covalent bonds. This chapter answers a different question entirely: what holds a metal together? A metal is an element on its own. There is no non-metal to hand electrons to and nothing to share with — yet copper keeps its shape, steel frames carry whole buildings, and aluminium survives being beaten into foil. The bonding inside a metal is the third and final bonding model of the course, and it reuses ideas you already own: ion formation from Chapter 7 plus electrons that are free to move.

All three objectives of record for this lesson are chemistry-only specification statements — assessed in Paper 2 (2C) only; they do not appear on Paper 1 (1C) and they are not shared with Science Double Award:

• know how to represent a metallic lattice by a 2D diagram (spec 1.52C);
• understand metallic bonding in terms of electrostatic attractions (spec 1.53C);
• explain typical physical properties of metals, including electrical conductivity and malleability (spec 1.54C).

The lesson runs in three movements after this map: the model itself — how sodium atoms become a lattice of positive ions in a sea of delocalised electrons, and exactly how you draw it; what the model predicts about melting points, by climbing the sodium-to-magnesium comparison; and the two properties examiners love — conductivity and malleability — ending with the close that ties every property back to the one structure sentence. Work through the sections in order; each one's model feeds the next one's explanations.

Watch sodium build a metal, because one atom tells the whole story. A sodium atom has the electronic configuration [2, 8, 1] — a single outer electron. When sodium atoms bond together to form the solid metal, that outer electron leaves its atom and becomes free to move throughout the whole structure. The electron is said to be delocalised-electron: no longer attached to any particular atom or pair of atoms — picture it flowing around through the entire metal. The atom it left behind is no longer neutral: 11 protons now pull on only 10 electrons, so a sodium ion, Na⁺, remains.

Multiply that event by the enormous number of atoms in even the smallest piece of metal and you reach the structure sentence of the whole chapter: a metallic structure consists of a lattice — a regular arrangement — of positive ions in a sea of delocalised electrons. The bonding holding it together, metallic-bonding, is the electrostatic forces of attraction between each positive ion and the delocalised electrons. Every ion pulls on the electrons around it; the electrons pull back on every ion; and that two-way attraction — the same electrostatic idea that held ionic lattices and covalent molecules together — locks the whole metal into shape.

Spec 1.52C asks you to put this structure on paper in two dimensions. Draw a grid of Na⁺ symbols in neat rows and columns, then scatter dots or small circles in the gaps between them for the delocalised electrons, and label one of each: "Na⁺ ion" at a grid point, "delocalised electron" in a gap. The feature the mark scheme insists on is regularity — the ions must sit in a regular repeating arrangement, not a random scatter. Different metals arrange their ions in slightly different patterns, but nobody expects you to reproduce a real crystal: a tidy grid is always accepted.

Two structural facts complete the model. First, a metal is a giant structure: there are no individual molecules anywhere in it, and every positive ion in the lattice attracts all the delocalised electrons — not just its nearest neighbours. Second, the charge on each ion is fixed by the atom's outer-shell count: every Group 1 metal forms 1+ ions and every Group 2 metal forms 2+ ions. Hold on to that charge rule — it becomes the engine of the melting-point story in the next section.

Most metals are hard, with high melting points — and the model explains both facts in a single move. The electrostatic forces of attraction between the positive ions and the delocalised electrons are strong. Pulling the structure apart means fighting that attraction everywhere at once, so it takes an enormous amount of energy, and the metal stays solid until it is heated strongly. Whenever an exam asks you to justify hardness or a high melting point, the answer is that attraction, and the energy needed to overcome it.

The comparison of sodium with magnesium turns the model into a prediction machine. Both atoms release their outer electrons into the shared pool, but they do not release the same number. Sodium atoms have one outer electron each, so sodium's lattice is held together by 1+ ions and one delocalised electron per atom — the bonding, as metals go, is quite weak, which is why sodium is fairly soft with a low melting point for a metal. Magnesium atoms have two outer electrons, both of which are delocalised, leaving ions that carry a 2+ charge — and there is a much stronger electrostatic attraction between the 2+ ions and the denser sea of delocalised electrons. Stronger attraction means more energy needed to break it, so magnesium's melting point sits well above sodium's.

Attach a piece of metal to a power source and the delocalised electrons move through the metal. Electrons are attracted towards the positive terminal of the source and flow away along the wire; electrons flow in from the negative terminal to replace them. A steady stream of electrons passes through the lattice — and because the electrons were already free to move everywhere in the structure, the metal conducts in the solid state, with no melting required. Nothing in the lattice itself needs to move for a metal to carry a current.

A metal is malleable: hammer it and it flattens into a new shape instead of shattering. The model explains this with a single motion. Applying a force makes the layers of positive ions slide over each other — and nothing about the bonding breaks when they do. After the slide, every ion is still surrounded by the same sea of delocalised electrons attracting it exactly as before, because every ion attracts all the electrons, wherever the ions end up. The structure rearranges; the attraction never switches off.

Set that beside the ionic crystals of Chapter 7 (u1:c07) and the difference becomes unforgettable. Slide a layer of an ionic crystal and like charges are dragged face to face — attraction flips to repulsion, and the crystal shatters, which is why ionic crystals are brittle. In a metal no such disaster is possible: the electron sea surrounds every ion in its new position just as it did in the old one, so the layers glide and the lattice re-seals itself. Same lattice idea, same electrostatic attraction — but the mobile electrons make one material bend and the other break.

Metals are also ductile: they can be drawn out into wires. The explanation is the same sliding-layers story told lengthways. Pull the metal and its layers slip past each other while the delocalised-electrons keep every ion bonded to its surroundings, so the wire stretches and thins instead of snapping. Copper's job inside a cable uses the pair of properties at once: ductile enough to be drawn into fine wire, conductive enough to carry the current through it.

This chapter completes the bonding picture of Unit 1. Transfer builds ionic lattices; sharing builds covalent molecules and giant covalent structures; release builds metallic lattices. Three ways of arranging electrostatic attraction, three families of structure, three sets of properties. Every structure-and-bonding question from here to the end of the course is one of those three stories — or, in the best questions, a comparison that asks you to tell them apart.