Displacement Reactions

Displacement reactions are a type of chemical reaction in which a more reactive metal replaces a less reactive metal in a compound. These reactions are known as redox (reduction-oxidation) reactions.

During a displacement reaction:

  • The more reactive metal undergoes oxidation, which means it loses electrons and becomes a cation
  • The less reactive metal undergoes reduction, which means it gains electrons and becomes an anion

An example of a displacement reaction is when iron is added to copper sulfate solution. Since iron is more reactive than copper, it replaces the copper in the compound, forming iron(II) sulfate and copper. The equation for this is:

Iron + Copper sulfate → Iron(II) sulfate + Copper

The image illustrates a chemical reaction using two test tubes and symbolic representations. On the left, a test tube contains a blue solution labelled "CuSO₄" with a iron nail labelled "Fe" immersed in it. On the right, the iron nail is coated in a brown layer, indicating the formation of copper, "Cu". The resulting solution in the test tube is labelled "FeSO₄". Above the test tubes, there are icons representing the reactants and products: a yellow circle for iron, and combined green and purple circles for copper sulphate. This results in a separate green circle for copper and yellow and purple circles combined to represent iron sulphate. At the bottom, the chemical equation "Fe + CuSO₄ → Cu + FeSO₄" summarises the reaction.

During the reaction, the blue colour of the copper sulfate solution changes as iron(II) sulfate forms. Also, copper coats the surface of the iron nail.

Deducing a Reactivity Series

We can use displacement reactions to determine the relative positions of metals in the reactivity series. To do this, we add a piece of metal to a solution of a metal salt and observe the reaction. Then we can predict the order of reactivity of the metals based on the results.

For example, let’s consider what happens when we add metals to three different metal salt solutions: silver nitrate, magnesium chloride, and zinc sulfate. The results are recorded in the table below:

MetalSilver nitrateMagnesium chlorideZinc sulphate
SilverNo reactionNo reactionNo reaction
MagnesiumMagnesium nitrate and silverNo reactionMagnesium sulfate and zinc
ZincZinc nitrate and silverNo reactionNo reaction

Based on these results, we can determine the relative order of reactivity of these metals.

  • Magnesium displaces both silver and zinc, indicating that it is the most reactive of the three
  • Silver is displaced by both magnesium and zinc, indicating that it is the least reactive
  • Zinc can displace silver but not magnesium, indicating that it is in between these two metals in terms of reactivity

Therefore, the order of reactivity, from most reactive to least, is magnesium, zinc and silver.

Mg > Zn > Ag

Displacement Reactions are Redox Reactions

Displacement reactions are redox (reduction-oxidation) reactions, meaning they involve the transfer of electrons between the reactants. Half equations can be used to represent the redox reactions that occur during displacement reactions.

For example, let’s look at the reaction between zinc (Zn) and copper sulfate (CuSO4):

Zn (s) + CuSO4 (aq) → Cu(s) + ZnSO4 (aq)

During this reaction, zinc displaces copper from the copper sulfate solution.

Zn + Cu²⁺ → Cu + Zn²⁺

Zinc atoms lose electrons to form zinc ions.

The equations for the oxidation and reduction that occur during this reaction are:

Oxidation: ZnZn²⁺ + 2e 

Reduction: Cu²⁺ + 2eCu

The oxidation half equation shows that zinc atoms lose electrons to form zinc ions, while the reduction half equation shows that copper ions (Cu²⁺) gain electrons to form copper atoms (Cu). These half equations show the transfer of electrons between the reactants.

The sulfate ion (SO4²⁻) is unchanged on both the product and the reactant side of the reaction. Therefore, it is considered a spectator ion, which means it does not participate in the redox reaction and does not need to be included in the half equations.

Oxidising Agents and Reducing Agents

Oxidation is the loss of electrons (or the gain of oxygen), while reduction is the gain of electrons (or the loss of oxygen.

When a substance undergoes oxidation, it is called the reducing agent, because it causes another substance to be reduced. In contrast, when a substance undergoes reduction, it is called the oxidising agent, because it causes another substance to be oxidised.

It’s important to remember that the oxidising and reducing agents are not always the same in different reactions. The identity of the oxidising and reducing agents depends on the specific reaction and the substances involved.

Let’s look at some examples to see how this works:

Example

Calcium + Zinc sulphate Calcium sulfate + Zinc

In this reaction, zinc sulfate is reduced to form zinc metal, while calcium metal is oxidised to form calcium sulfate.

Therefore, zinc sulfate is the oxidising agent, because it is causing the oxidation of calcium metal. Meanwhile, calcium metal is the reducing agent, because it is causing the reduction of zinc sulfate.

Example

Aluminium + Copper oxide Aluminium oxide + Copper

In this reaction, copper oxide is reduced to form copper metal, while aluminium metal is oxidised to form aluminium oxide.

Therefore, copper oxide is the oxidising agent, because it is causing the oxidation of the aluminium metal. Meanwhile, aluminium metal is the reducing agent, because it is causing the reduction of copper oxide.

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