What is a Single-replacement Reaction?

A single-replacement reaction, also called a single-displacement reaction, is a reaction in which one element is substituted for another element in a compound. In order to understand the chemistry of a single-replacement reaction, one has to know about the activity series of elements in the periodic table. Elements that are higher in the series have a greater tendency to gain or lose electrons and hence, can easily displace elements that are lower in the series. Both metals and nonmetals participate in replacement reactions ^[1-4].

General Equation for Single-replacement Reaction

The single-displacement reaction equation is:

A + BC → AC + B

This reaction will proceed if A is more reactive than B. The reaction must be balanced when writing an actual reaction ^[1].

What Happens in a Single-replacement Reaction?

A characteristic of a single-replacement reaction is that one cation or anion trades places with another to form a new product. Metals that are more reactive than hydrogen readily dissolve in water. They form hydroxides and release hydrogen gas.

Examples of Single-replacement Reaction

An example of a single-displacement reaction occurs when potassium (K) reacts with water (H₂O). A colorless solid compound named potassium hydroxide (KOH) forms and hydrogen gas (H₂) is set free. The equation for the reaction is:

2 K (s) + 2 H₂O (l) → 2 KOH + H₂ (g)

How to Balance Single-replacement Reaction?

Consider the following example. Magnesium (Mg) replaces hydrogen (H) in hydrochloric acid (HCl) and forms magnesium chloride (MgCl₂) and gaseous hydrogen (H₂).

Mg (s) + HCl (aq.) → MgCl₂ (s) + H₂ (g)

This equation is not balanced. We note that two hydrogen and chlorine atoms are on the right-hand side of the equation. So, we multiply the compound HCl on the left by 2 and balance the two atoms.

Mg (s) + 2 HCl (aq.) → MgCl₂ (s) + H₂ (g)

Rules for Prediction of Single-replacement Reaction

The periodic table or an activity series can help to predict whether single-replacement reactions occur. The relative reactivities of the two elements can determine whether one element will replace another element from a compound. A single-replacement reaction will occur when a less reactive element can be replaced by a more reactive element in a compound ^[2].

1. Halogens

Order of reactivity of the halogens:

More reactive     F₂  >  Cl₂  >  Br₂  >  I₂    Less reactive

From this order, one can see that chlorine (Cl₂) will replace bromine (Br₂) from a bromide compound but cannot replace fluorine (F₂) from a fluoride compound.

2. Metals

Order of reactivity of metals:

More reactive    Cu  >  Ag  >  Hg  >  Pt  >  Au    Less reactive

This order shows that copper (Cu) will replace silver (Ag) in an aqueous solution consisting of Ag⁺ ions but not vice-versa.

Types of Single-replacement Reaction

There are two types of single-replacement reactions ^[2-4].

1. Cation Replacement

A cation is a positively charged ion or a metal. In this type of reaction, one cation replaces another.

Examples

• Zinc (Zn) can safely dissolve in hydrochloric acid (HCl), releasing hydrogen (H₂) gas ^[1].

Zn (s) + 2 HCl (aq.) → ZnCl₂ (s) + H₂ (g)

• Copper (Cu) can displace silver (Ag) in an aqueous solution of silver nitrate (AgNO₃).

Cu (s) + 2 AgNO₃ (aq.) → Cu(NO₃)₂ (aq.) + 2 Ag (s/ppt.)

2. Anion Replacement

An anion is a negatively charged ion or a nonmetal. In this type of reaction, one anion replaces another.

Examples

• Chlorine (Cl₂) will displace the bromine (Br₂) in a sodium bromide (NaBr) compound because it is more reactive than bromine ^[1].

Cl₂ (aq.) + 2 NaBr (aq.) → 2 NaCl (aq.) + Br₂ (aq.)

• Bromine (Br₂) replaces iodine (I₂) in a reaction between bromine (Br₂) and potassium iodide (KI).

Br₂ (aq.) + 2 KI (aq.) → 2 KBr (aq.) + I₂ (aq.)

Single-replacement Reactions Examples in Everyday Life

A few examples of single-replacement reactions occur in nature and real life.

• The reaction of saltwater with concrete pillars containing iron forms iron (II) chloride
• Silver reacts with hydrogen sulfide gas produced by some industrial processes or from decaying animal or plant materials. The reaction results in silver sulfide and hydrogen gas.