Single-Replacement Reactions and the Activity Series

In a single-replacement reaction, one element replaces another element in a compound, A + BC → AC + B, and the activity series ranks metals (and halogens) by how readily they undergo this kind of reaction, so that a more active element will replace a less active one. In brief: zinc replaces copper in copper(II) sulfate (Zn + CuSO₄ → ZnSO₄ + Cu) because zinc is higher in the activity series, but copper does not replace zinc; metals above hydrogen in the series displace hydrogen from acids.

Single-replacement reactions are redox reactions in which one element is oxidized and another reduced. They explain the rusting of iron, the action of acids on metals, the extraction of some metals and the working of simple batteries. This article explains the three types of single replacement, presents the activity series with standard electrode potentials, shows how to use it to predict whether a reaction will happen, works through examples and calculations, and gives some history.

What is a single-replacement reaction?

There are three common types:

A metal replaces another metal in a compound: Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s). Zinc dissolves, and brown copper deposits on the zinc.
A metal replaces hydrogen from water or an acid: Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g), and 2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g).
A halogen replaces a less active halogen: Cl₂(aq) + 2KBr(aq) → 2KCl(aq) + Br₂(aq).

In the net ionic equation for the first reaction, Zn + Cu²⁺ → Zn²⁺ + Cu, zinc loses two electrons (oxidation) and copper(II) gains them (reduction).

What is the activity series?

The activity series is a list of elements arranged by their reactivity in these reactions, from the most reactive at the top to the least reactive at the bottom. An element will replace any element below it from a compound in solution, but not an element above it. The ranking corresponds closely to the standard reduction potentials: the more negative the potential, the more readily the metal gives up electrons.

The activity series of metals

Activity series with standard reduction potentials and typical reactions
Metal	Ion	E° (V)	Reaction with water and acids
Lithium	Li⁺	−3.04	Reacts with cold water
Potassium	K⁺	−2.93	Reacts violently with cold water
Calcium	Ca²⁺	−2.87	Reacts with cold water
Sodium	Na⁺	−2.71	Reacts vigorously with cold water
Magnesium	Mg²⁺	−2.37	Reacts slowly with cold water, quickly with steam and acids
Aluminum	Al³⁺	−1.66	Reacts with acids (an oxide layer slows the reaction)
Zinc	Zn²⁺	−0.76	Reacts with steam and acids
Iron	Fe²⁺	−0.44	Reacts with acids; slowly with steam
Tin	Sn²⁺	−0.14	Reacts slowly with acids
Lead	Pb²⁺	−0.13	Reacts very slowly with acids
Hydrogen	H⁺	0.00	Reference
Copper	Cu²⁺	+0.34	No reaction with water or HCl
Silver	Ag⁺	+0.80	No reaction
Gold	Au³⁺	+1.50	No reaction, even with most acids

How do you predict whether a reaction occurs?

Find the free element and the metal in the compound in the activity series.
If the free element is higher in the series than the metal in the compound, the reaction occurs. If it is lower, there is no reaction.
Write the products: the free element goes into the compound with the anion, and the displaced metal becomes the free element.
Balance the equation, using the charge of the new cation.
For reactions with acids, only metals above hydrogen displace hydrogen. For reactions with water, only the most active metals react with cold water, and the next ones react with steam to give the metal oxide and H₂.

What about the halogens?

The activity of the halogens decreases down the group: F₂ > Cl₂ > Br₂ > I₂. A more active halogen displaces a less active one from its salt, so Cl₂ + 2NaI → 2NaCl + I₂ occurs, but I₂ + 2NaCl → no reaction. This corresponds to the standard potentials for reduction: F₂ +2.87 V, Cl₂ +1.36 V, Br₂ +1.07 V and I₂ +0.54 V.

Worked examples

Will it react? Cu + ZnSO₄: copper is below zinc in the series, so there is no reaction. Zn + CuSO₄: zinc is above copper, so the reaction occurs.
Write the equation. Aluminum with copper(II) chloride: Al is above Cu. Al + CuCl₂ → AlCl₃ + Cu. Balance: 2Al + 3CuCl₂ → 2AlCl₃ + 3Cu.
Acid and metal. Magnesium ribbon of 0.243 g (0.0100 mol) reacts with excess hydrochloric acid. From Mg + 2HCl → MgCl₂ + H₂, 0.0100 mol of H₂ forms. At 25 °C and 1 atm, the volume is 0.0100 × 24.5 = 0.245 L = 245 mL.
Copper with acid. Cu + 2HCl: copper is below hydrogen, so no reaction occurs and no gas is produced.
Silver recovery. A copper wire is placed in silver nitrate solution: Cu + 2AgNO₃ → Cu(NO₃)₂ + 2Ag. How much silver forms if 0.635 g of copper dissolves? n(Cu) = 0.635 ÷ 63.55 = 0.01000 mol, so n(Ag) = 0.0200 mol, with a mass of 0.0200 × 107.87 = 2.16 g. The wire gains silver and loses copper, and the solution turns blue.
The thermite reaction. 2Al + Fe₂O₃ → Al₂O₃ + 2Fe is a single replacement in which aluminum displaces iron from its oxide. ΔH = −852 kJ per mole of Fe₂O₃. For 2.70 g of Al (0.100 mol): energy = 0.100 ÷ 2 × 852 = 42.6 kJ. The temperature reaches over 2,000 °C, enough to melt the iron.
Cell potential. For the zinc and copper reaction, E°_cell = E°(Cu²⁺/Cu) − E°(Zn²⁺/Zn) = 0.34 − (−0.76) = +1.10 V. A positive value shows that the reaction is spontaneous.

Where are these reactions used?

Batteries: the Daniell cell, with zinc and copper, produces a voltage of about 1.1 V.
Corrosion protection: galvanizing coats iron with zinc, which is more active and corrodes in place of the iron (sacrificial protection). Magnesium blocks are attached to ship hulls and pipelines.
Metal recovery: scrap iron added to copper-containing solutions displaces copper (cementation).
Hydrogen production: in the laboratory, zinc or magnesium with acid.
Welding rails: the thermite reaction.
Bromine and iodine production: chlorine displaces bromine from brine.

How was the activity series developed?

The relative reactivity of metals was known from early experience: gold and silver are found native, while sodium and potassium were unknown until electrolysis. Humphry Davy isolated several reactive metals in 1807 and 1808. Alessandro Volta ranked metals by their ability to generate a voltage when in contact, in 1800, and the voltaic series of the early 19th century is an ancestor of the activity series. Walther Nernst developed the theory of electrode potentials in 1889, and the potentials were then measured relative to the standard hydrogen electrode.

Which mistakes are common?

Mistake 1: Reversing the direction of replacement

A higher metal replaces a lower metal, not the other way round.

Mistake 2: Writing the wrong charge for the new cation

Aluminum gives Al³⁺ and the product is AlCl₃, not AlCl. Iron usually gives Fe²⁺ with acids and copper(II) salts.

Mistake 3: Expecting all metals to react with water

Only the most active metals react with cold water. Iron and copper do not.

Mistake 4: Forgetting that the reaction needs the metal to be in contact with the solution

Coatings of oxide, as on aluminum, can prevent a reaction until the film is removed.