the relative reactivity of metals as shown in the activity series, the prediction of metal displacement reactions in aqueous solution, and the relationship between metal reactivity and reactions with water, acids and oxygen

What this dot point is asking

VCAA wants you to use the metal activity series to predict and justify three things: (i) whether a metal will displace another metal from a solution of its ions, (ii) how vigorously a metal reacts with water, dilute acids and oxygen, and (iii) the role of each metal as the reductant in any reaction that does occur.

The answer

The activity series

Metals can be ranked by how readily they lose electrons (how readily they are oxidised). The more readily they lose electrons, the more reactive they are. A workable VCE-level order, most reactive first:

Hydrogen is included as a reference even though it is not a metal: any metal above hydrogen will react with a dilute acid to produce hydrogen gas; any metal below hydrogen will not.

The order matches the standard reduction potentials in the electrochemical series, read in the opposite direction. The most reactive metal (potassium) has the most negative reduction potential for its $K^+ / K$ couple; the least reactive (gold) has the most positive.

Predicting a metal displacement reaction

A metal displacement reaction is one where a metal in elemental form reduces the cation of a less reactive metal, taking its place in solution:

The rule: the displacing metal must be more reactive (further up the activity series) than the metal in the salt.

Example: $Zn$ is more reactive than $Cu$, so:

This is a redox reaction. Half-equations:

Oxidation: $Zn(s) \to Zn^{2+}(aq) + 2e^-$

Reduction: $Cu^{2+}(aq) + 2e^- \to Cu(s)$

The zinc is the reductant (it donates electrons; it is oxidised). The copper ion is the oxidant (it accepts electrons; it is reduced).

If you reverse the direction and try $Cu(s) + Zn^{2+}(aq)$: no reaction, because copper is less reactive than zinc and cannot give up its electrons to a zinc ion.

Reaction with water

The most reactive metals (group 1 and the more reactive group 2) react directly with cold water to produce hydrogen gas and a metal hydroxide:

Magnesium reacts slowly with cold water and more rapidly with hot water or steam. Aluminium and iron react with steam at high temperature. Copper, silver and gold do not react with water at any practical temperature.

Reaction with dilute acid

Any metal above hydrogen in the activity series reacts with a dilute acid (such as $HCl$ or dilute $H_2SO_4$) to give hydrogen gas and a soluble salt:

Net ionic: $Mg(s) + 2H^+(aq) \to Mg^{2+}(aq) + H_2(g)$.

The rate is set by the metal's reactivity: $K$, $Na$ and $Ca$ react explosively (do not perform); $Mg$ reacts vigorously; $Zn$ and $Fe$ react steadily; $Sn$ and $Pb$ react slowly. Copper, silver and gold do not react with dilute non-oxidising acids. (Note: copper does react with concentrated or hot nitric acid, but that is an oxidising acid; that chemistry is outside VCE Unit 2.)

Reaction with oxygen

All but the noble metals react with oxygen, though the rate and the temperature required differ enormously.

• The most reactive metals burn brightly when ignited and form a metal oxide: $2Mg(s) + O_2(g) \to 2MgO(s)$ (the classic VCE prac).
• Less reactive metals tarnish slowly at room temperature (iron rusts; copper develops a green patina).
• Gold and platinum do not corrode in air. This stability is why gold is used for electrical contacts and jewellery.

Putting it together

A metal that is high in the activity series is also reactive with water, with dilute acid and with oxygen, and is a strong displacer of less reactive metals. A metal at the bottom is unreactive with all three and cannot displace any other metal in solution.

This is why the metal activity series and the predictive rules for displacement, water and acid reactivity are taught together: they all reflect the same underlying property, the metal's tendency to lose electrons.

Common traps

Predicting a reaction that the activity series rules out. Always check: is the elemental metal above or below the metal in the salt? If below, the answer is "no reaction" and that is the full answer (no equation needed).

Forgetting that a metal below hydrogen will not react with dilute acid. Adding copper to dilute hydrochloric acid gives no visible change. Many students write a reaction anyway out of habit.

Writing the wrong charge on the metal ion. Iron forms both $Fe^{2+}$ and $Fe^{3+}$; the typical product of a simple displacement is $Fe^{2+}$. Copper forms $Cu^+$ and $Cu^{2+}$; the standard form in solution is $Cu^{2+}$.

Including the spectator anion in the net ionic equation. The anion of the salt ($SO_4^{2-}$, $NO_3^-$, $Cl^-$) is always a spectator and is not shown.

Confusing reactivity with electronegativity. A reactive metal is a strong reductant (gives up electrons easily). Electronegativity describes a covalent bonding partner's attraction for shared electrons. Sodium is very reactive (low electronegativity 0.93); fluorine is highly electronegative (3.98) and is itself reactive but on the oxidising side, not the metal side.

Saying group 1 metals "explode" because they are unstable. They are reactive, not unstable. They are stable in their unreactive state under oil; the explosive behaviour is the rapid reaction with water releasing hydrogen which ignites.

In one sentence

The metal activity series ranks metals by how readily they donate electrons, so a higher metal will displace a lower one from its salt solution, will react with water or dilute acid (only if above hydrogen) to give hydrogen gas, and will be the reductant in every redox reaction it is part of.