Reactivity series

Not all metals are equally worth their salt when it comes to reacting. Drop a lump of potassium into water and it skates about furiously, bursting into flame; leave a gold ring in the sea for a thousand years and it comes out unchanged. Chemists capture that contrast in a single ranked list — the reactivity-series — and this chapter builds it, uses it and then explains the reactions that hold it together. The series does more than rank metals: it introduces the idea of a redox-reaction, in which one substance takes oxygen (or electrons) and another gives it up, and it powers one of the most useful reaction types in chemistry, the displacement-reaction, where a more reactive metal steals a less reactive metal's place in a compound.

The payoff for learning the series is enormous, and it is not abstract. When aluminium reacts with iron(III) oxide, the reaction produces molten iron hot enough to weld railway rails together on site — a displacement reaction doing industrial work. When zinc is bolted onto a ship's steel hull, the zinc corrodes so the iron does not — the series predicting which metal sacrifices itself. Gold stays bright forever precisely because it sits at the very bottom of the series: in contact with air or water it stays chemically unchanged, essentially forever.

Here is the route through the chapter. First you will learn how reactions with water and with dilute acids arrange metals in the series and how displacement reactions — with metal oxides and with salt solutions — refine the ordering, meeting oxidation, reduction and the agent vocabulary on the way. Then comes the order itself, learned straight, followed by the two flagship applications: why iron rusts and how rusting is beaten with barriers, galvanising and sacrificial protection. A listed practical ties the whole idea down: investigating reactions between dilute hydrochloric and sulfuric acids and metals such as magnesium, zinc and iron.