The Activity Series of Metals and Halogens

The previous two chapters built the tools to describe how a redox reaction is balanced once it is known to happen. This chapter asks a different question, which logically comes first: given two possible reactants, will a redox reaction happen between them at all, and if so, in which direction? The answer, for now, is organised empirically into two ranked lists — the activity series of metals and the activity series (oxidising power) of halogens — built directly from the observed outcomes of displacement reactions. These lists are the qualitative ancestor of the quantitative standard electrode potentials met two chapters from now; everything developed here is the groundwork that makes that later, more powerful tool make physical sense.

By the end of this chapter you will be able to predict, from a single ranked list, whether a given metal will displace another from solution, whether a metal will react with dilute acid or with water, and whether a given halogen will displace another halide from solution — all without yet needing a single numerical electrode potential.

1. Displacement Reactions as Evidence of Relative Reactivity

A displacement reaction occurs when a more reactive element takes the place of a less reactive element already combined in a compound. For metals, the clearest evidence comes from placing a metal into a solution containing the ion of a different metal:

Using the oxidation-state and half-equation tools of the previous two chapters, this is recognised immediately as a redox reaction: zinc is oxidised ($0 \to +2$), losing two electrons, and copper(II) is reduced ($+2 \to 0$), gaining them. The reverse experiment — placing copper metal into zinc sulfate solution — produces no observable reaction at all. This asymmetry is the experimental fact the activity series exists to record: zinc gives up its electrons to copper(II) far more readily than copper gives up its electrons to zinc(II).

Key distinction. A metal displaces a second metal from solution only if the first metal is more reactive — that is, more willing to be oxidised (to lose electrons) — than the second. Reactivity, in this specific sense, is a statement about how readily a species acts as a reducing agent.

2. The Activity Series of Metals

Testing many such pairs systematically produces a single consistent ranking, independent of which pair is tested — exactly as a transitive ordering should. The resulting activity series (or reactivity series) lists common metals from most to least readily oxidised:

Hydrogen is conventionally included, in parentheses, as a reference point, since it is not a metal but behaves as the dividing line for reactions with dilute acids (Section 4). A metal higher in this list is a stronger reducing agent: it loses its own electrons more readily than a metal below it, and correspondingly its ions are harder to reduce (less willing to gain electrons back) than the ions of a metal below it.

Definition. The activity series of metals ranks metals by how readily they are oxidised (lose electrons) to form their cations. The most reactive metals sit at the top; the least reactive — the so-called "noble" metals such as silver and gold — sit at the bottom.

2.1 The displacement rule

Rule. A metal $\text{M}$ will displace a second metal $\text{N}$ from a solution of $\text{N}$'s ions if, and only if, $\text{M}$ lies above $\text{N}$ in the activity series.

Worked example. Predict the outcome, if any, of adding iron filings to silver nitrate solution, and of adding silver to iron(II) sulfate solution.

Iron lies above silver in the activity series, so iron displaces silver:

(Balanced using the half-equation method: $\text{Fe} \to \text{Fe}^{2+} + 2e^-$ combines with $2\times[\text{Ag}^+ + e^- \to \text{Ag}]$.) Silver, lying below iron, cannot displace it: adding silver metal to iron(II) sulfate solution produces no reaction at all.

2.2 Why the ordering is consistent

The activity series works as a single, universal ranking (rather than needing to be re-measured for every possible pair) because reactivity, in this sense, is a single underlying property of each metal — how strongly it "wants" to lose its valence electrons — set by the balance of its ionisation energy and the energy released on forming its solid lattice or hydrated ion. A full quantitative account of why the ordering comes out exactly as it does is given by the standard electrode potentials of two chapters ahead; for now, the ordering is taken as an empirical, directly testable fact.

3. Metals Reacting with Water

The activity series also predicts, qualitatively, how vigorously a metal reacts with water, since this reaction is itself a redox displacement: the metal reduces water's hydrogen (from $+1$ in $\text{H}_2\text{O}$ to $0$ in $\text{H}_2$), being itself oxidised.

The very top of the series reacts readily with cold water: potassium and sodium react so vigorously that the hydrogen produced often ignites, while calcium reacts more steadily, with bubbling and some warming. Magnesium reacts only extremely slowly with cold water but readily with steam. Zinc and iron react with steam but not with cold or even hot liquid water under ordinary conditions. Metals from lead downward show essentially no reaction with water at all under normal laboratory conditions.

Worked example. Potassium dropped into water: $2\text{K}(s) + 2\text{H}_2\text{O}(l) \longrightarrow 2\text{KOH}(aq) + \text{H}_2(g)$. Potassium's oxidation state rises from $0$ to $+1$; hydrogen's falls from $+1$ (in water) to $0$ (in $\text{H}_2$) — a redox reaction entirely consistent with potassium's position at the very top of the series.

4. Metals Reacting with Dilute Acids

A metal above hydrogen in the activity series reduces the $\text{H}^+$ ion of a dilute acid to hydrogen gas, being itself oxidised to its cation:

Worked example. Magnesium ribbon added to dilute hydrochloric acid fizzes vigorously:

Magnesium, well above hydrogen in the series, is readily oxidised ($0 \to +2$) while hydrogen is reduced ($+1 \to 0$). By contrast, copper — below hydrogen — shows no reaction with dilute hydrochloric acid or dilute sulfuric acid: copper is simply not a strong enough reducing agent to transfer electrons to $\text{H}^+$. This is precisely why hydrogen's position in the series, though it is not itself a metal, is such a useful reference point: it marks the boundary between metals that liberate hydrogen gas from a dilute acid and metals that do not.

Warning. This rule concerns dilute, non-oxidising acids such as dilute $\text{HCl}$ or dilute $\text{H}_2\text{SO}_4$, where $\text{H}^+$ itself is the oxidising agent. Concentrated nitric acid and concentrated sulfuric acid are themselves strong oxidising agents (as seen in the half-equations of the previous chapter) and can oxidise even metals below hydrogen, such as copper, by a different mechanism entirely; that behaviour is not predicted by the simple activity-series rule given here.

5. The Activity Series of Halogens

An entirely analogous ranking exists for the halogens, though in the opposite sense: since a halogen reacts by gaining an electron to form its halide ion, the halogens are ranked by oxidising power — how readily each one is itself reduced.

This ordering decreases down Group 17, tracking the trend of decreasing electronegativity and increasing atomic radius established in an earlier block: a smaller, more electronegative halogen atom attracts an incoming electron more strongly, making $\text{F}_2$ the most powerful halogen oxidising agent and $\text{I}_2$ the weakest.

Rule. A halogen $\text{X}_2$ will displace a second halogen $\text{Y}_2$ from a solution of $\text{Y}^-$ ions if, and only if, $\text{X}_2$ lies above $\text{Y}_2$ in the activity series (i.e. is the stronger oxidising agent).

Worked example. Chlorine water added to potassium bromide solution turns the solution orange-brown as bromine is liberated:

In ionic form, with potassium a spectator: $\text{Cl}_2(aq) + 2\text{Br}^{-}(aq) \longrightarrow 2\text{Cl}^{-}(aq) + \text{Br}_2(aq)$. Chlorine, the stronger oxidising agent, is reduced ($0 \to -1$) while bromide is oxidised ($-1 \to 0$). The reverse — adding bromine water to potassium chloride solution — produces no reaction, since bromine cannot out-compete chlorine for the electrons already held by chloride.

Worked example. Bromine water added to potassium iodide solution turns the solution dark brown/black (or purple in an organic solvent layer) as iodine is liberated:

Fluorine is excluded from these aqueous displacement experiments in practice: it is so powerfully oxidising that it reacts violently and directly with water itself, rather than cleanly displacing another halide, so the series $\text{F}_2 > \text{Cl}_2 > \text{Br}_2 > \text{I}_2$ is established by combining fluorine's known extreme reactivity with the aqueous displacement results among chlorine, bromine, and iodine.

Key distinction. The metal activity series ranks metals by how readily they are oxidised (lose electrons); the halogen activity series ranks halogens by how readily they are reduced (gain electrons). A metal high in its series and a halogen high in its series are therefore opposite in character: one is a powerful reducing agent, the other a powerful oxidising agent — and, usefully, a strong reducing agent reacts most readily with a strong oxidising agent, which is exactly why reactive metals and reactive halogens combine so vigorously.

6. Using the Activity Series to Predict Reaction Feasibility

Both series answer the same underlying question in their respective domains: will electron transfer actually happen, and in which direction? The general principle is that a reaction proceeds in the direction that lets the stronger reducing agent present hand its electrons to the stronger oxidising agent present, not the reverse.

Worked example. Will solid lead react with a solution of tin(II) chloride?

Checking the activity series: $\dots \text{Fe} > \text{Pb} > \dots$, and tin falls between iron and hydrogen, specifically above lead: $\text{Sn}$ sits above $\text{Pb}$ in an extended activity series. Since lead is below tin, lead cannot displace tin from solution, and no reaction occurs. (Were the question reversed — tin metal added to lead(II) chloride solution — a reaction would occur, since tin, the stronger reducing agent here, would displace lead.)

7. Common Exam Pitfalls

• Reversing the displacement rule. The higher metal (or halogen, for oxidising power) is the one that displaces the other; a common error is checking only "are these two different elements" without checking which one is actually higher in the ranking.
• Applying the dilute-acid rule to concentrated oxidising acids. Concentrated nitric or sulfuric acid can oxidise metals below hydrogen (such as copper) by a mechanism unrelated to simple $\text{H}^+$ reduction; the activity series alone does not predict this.
• Forgetting that halogen oxidising power decreases down the group, the opposite direction to the metal series' pattern of reactivity decreasing down the list — both lists run "most reactive at the top," but "reactive" means opposite things (oxidised easily vs reduced easily) in each case.
• Treating "no visible reaction" as meaning no redox chemistry is possible in principle. It specifically means the proposed direction of electron transfer is unfavourable; the reverse reaction (if it were tried) might well be favourable instead.
• Omitting fluorine's exceptional behaviour. Fluorine's position at the top of the halogen series is inferred from its known extreme reactivity, not from a clean aqueous displacement experiment, since it reacts destructively with water itself.
• Forgetting to balance the displacement equation, especially when the ionic charges of the two metals differ (e.g. $\text{Fe} + 2\text{Ag}^+ \to \text{Fe}^{2+} + 2\text{Ag}$, not a 1:1 ratio) — the half-equation method from the previous chapter still applies in full.

Summary

• A displacement reaction occurs when a more reactive element takes the place of a less reactive element already in a compound, and is itself a redox reaction analysable with the previous two chapters' tools.
• The activity series of metals ranks metals by how readily they are oxidised: $\text{K} > \text{Na} > \text{Ca} > \text{Mg} > \text{Al} > \text{Zn} > \text{Fe} > \text{Pb} > (\text{H}) > \text{Cu} > \text{Ag} > \text{Au}$.
• A metal displaces a second metal from solution only if it lies above that metal in the series; the same ranking predicts reaction with water (vigorous at the top, negligible below lead) and with dilute acid (vigorous above hydrogen, absent below it).
• The activity series of halogens, $\text{F}_2 > \text{Cl}_2 > \text{Br}_2 > \text{I}_2$, ranks halogens by oxidising power — how readily each is reduced — decreasing down the group with decreasing electronegativity.
• A halogen displaces a second halide from solution only if it lies above that halogen in its series; fluorine's extreme reactivity precludes its use in a clean aqueous displacement test.
• Both series are qualitative predictors of reaction feasibility and direction; the fully quantitative version of the same idea, standard electrode potentials, is developed two chapters from now, once galvanic cells have been introduced.