Unit 6: Common Redox Reactions
Exploring practical examples of redox reactions in chemistry.
6.7 Reaction of Metals with Acids
The reaction between a reactive metal and an acid is a classic example of a redox reaction. In this single displacement reaction, the metal is oxidised and the hydrogen from the acid is reduced.
General Equation: Metal + Acid → Salt + Hydrogen Gas
Example: Zinc reacting with hydrochloric acid.
$Zn(s) + 2HCl(aq) \rightarrow ZnCl_2(aq) + H_2(g)$
- Oxidation: The zinc metal loses two electrons and its oxidation number increases from 0 to +2.
$Zn \rightarrow Zn^{2+} + 2e^-$ - Reduction: The hydrogen ions from the acid gain electrons and their oxidation number decreases from +1 to 0.
$2H^+ + 2e^- \rightarrow H_2$
This reaction is also a useful method for preparing salts of reactive metals (see Unit 5).
Solved Examples:
- What are the products when magnesium reacts with sulphuric acid?
Solution: Magnesium sulphate ($MgSO_4$) and hydrogen gas ($H_2$). - Write the oxidation half-equation for the reaction of aluminium with hydrochloric acid.
Solution: $Al \rightarrow Al^{3+} + 3e^-$. - Identify the oxidising agent in the reaction $Fe(s) + 2H^+(aq) \rightarrow Fe^{2+}(aq) + H_2(g)$.
Solution: The hydrogen ion ($H^+$) is the oxidising agent because it is reduced (gains electrons). - Will copper react with dilute hydrochloric acid? Why or why not?
Solution: No. Copper is less reactive than hydrogen and cannot displace it from an acid. - What gas is produced when a metal reacts with an acid?
Solution: Hydrogen gas ($H_2$). - Write a balanced chemical equation for the reaction of calcium with nitric acid.
Solution: $Ca(s) + 2HNO_3(aq) \rightarrow Ca(NO_3)_2(aq) + H_2(g)$. - In the reaction of zinc with acid, what happens to the oxidation number of zinc?
Solution: It increases from 0 to +2. - What is the reducing agent when iron reacts with sulphuric acid?
Solution: Iron metal ($Fe$) is the reducing agent because it is oxidised. - How could you test for the gas produced in these reactions?
Solution: Collect the gas and place a lit splint in it. Hydrogen gas will extinguish the splint with a 'squeaky pop'. - Why is this reaction considered a single displacement reaction?
Solution: Because the more reactive metal displaces the less reactive hydrogen from the acid compound.
6.8 Reaction of Metals with Water and Bases
Very reactive metals can also undergo redox reactions with water, and in some cases, with strong bases.
Reaction with Water:
Highly reactive metals (like those in Group 1) can displace hydrogen from water, even at room temperature. The metal is oxidised, and hydrogen in water (with an oxidation state of +1) is reduced to hydrogen gas.
Example: Sodium reacting with water.
$2Na(s) + 2H_2O(l) \rightarrow 2NaOH(aq) + H_2(g)$
- Oxidation: $Na \rightarrow Na^+ + e^-$
- Reduction: $2H_2O + 2e^- \rightarrow H_2 + 2OH^-$
Reaction with Bases:
Amphoteric metals (like zinc and aluminium) can react with strong alkalis. In these reactions, the metal is oxidised, and hydrogen is once again reduced from the water molecules present in the aqueous base.
Example: Aluminium reacting with sodium hydroxide solution.
$2Al(s) + 2NaOH(aq) + 6H_2O(l) \rightarrow 2Na[Al(OH)_4](aq) + 3H_2(g)$
Solved Examples:
- What are the products when potassium reacts with water?
Solution: Potassium hydroxide ($KOH$) and hydrogen gas ($H_2$). - In the reaction of sodium with water, what is the oxidising agent?
Solution: Water ($H_2O$) is the oxidising agent, as the hydrogen within it is reduced from +1 to 0. - Which metal would react more vigorously with water: lithium or caesium?
Solution: Caesium, as reactivity in Group 1 metals increases down the group. - Write the balanced equation for the reaction of calcium with water.
Solution: $Ca(s) + 2H_2O(l) \rightarrow Ca(OH)_2(aq) + H_2(g)$. - What type of metal can react with a strong base like NaOH?
Solution: An amphoteric metal, such as zinc or aluminium. - What is the oxidation number of hydrogen in $H_2$ gas?
Solution: 0. - What is the oxidation number of hydrogen in water?
Solution: +1. - What is the reducing agent when zinc reacts with sodium hydroxide solution?
Solution: Zinc metal ($Zn$) is the reducing agent as it gets oxidised. - Why doesn't a less reactive metal like iron react with cold water?
Solution: Iron is not reactive enough to displace hydrogen from cold water. It will, however, react with steam. - What is the name of the product $Na[Al(OH)_4]$?
Solution: Sodium tetrahydroxoaluminate(III).
6.9 Metal Displacement Reactions
A more reactive metal will displace a less reactive metal from a solution of its salt. This is a direct application of the reactivity series (or standard electrode potentials).
General Equation: More Reactive Metal + Salt of Less Reactive Metal → Salt of More Reactive Metal + Less Reactive Metal
Example: A piece of zinc placed in copper(II) sulphate solution.
$Zn(s) + CuSO_4(aq) \rightarrow ZnSO_4(aq) + Cu(s)$
- Oxidation: The more reactive metal, zinc, is oxidised.
$Zn \rightarrow Zn^{2+} + 2e^-$ - Reduction: The ion of the less reactive metal, copper(II), is reduced.
$Cu^{2+} + 2e^- \rightarrow Cu$
The sulphate ion ($SO_4^{2-}$) is a spectator ion as it does not participate in the reaction.
Solved Examples:
- What would you observe if you placed a piece of magnesium in a solution of zinc nitrate?
Solution: The magnesium strip would dissolve, and a layer of solid zinc metal would form on its surface. - Write the ionic equation for the reaction between magnesium and zinc nitrate.
Solution: $Mg(s) + Zn^{2+}(aq) \rightarrow Mg^{2+}(aq) + Zn(s)$. - Identify the reducing agent in the reaction from the previous question.
Solution: Magnesium ($Mg$) is the reducing agent. - Would a reaction occur if you placed a strip of copper in a solution of magnesium sulphate?
Solution: No, because copper is less reactive than magnesium and cannot displace it. - What is a spectator ion? Give an example from the reaction of zinc with copper(II) sulphate.
Solution: An ion that is present in the solution but does not take part in the chemical reaction. The sulphate ion ($SO_4^{2-}$) is the spectator ion. - A student places an iron nail into a blue solution of copper(II) sulphate. What changes would be observed?
Solution: The iron nail would become coated with a layer of pinkish-brown copper metal, and the blue colour of the solution would fade (as $Cu^{2+}$ ions are consumed). - Write the oxidation half-equation for the reaction of iron with copper(II) sulphate.
Solution: $Fe(s) \rightarrow Fe^{2+}(aq) + 2e^-$. - Which is the stronger reducing agent: zinc or copper?
Solution: Zinc, because it is more easily oxidised (it is more reactive). - Complete the reaction: $2AgNO_3(aq) + Cu(s) \rightarrow$
Solution: $2AgNO_3(aq) + Cu(s) \rightarrow Cu(NO_3)_2(aq) + 2Ag(s)$. - What is the oxidising agent in the reaction from the previous question?
Solution: The silver ion ($Ag^+$).
6.10 Formation and Reduction of Metallic Oxides
Formation of Oxides:
The reaction of a metal with oxygen to form a metal oxide is a common oxidation process.
Example: Formation of aluminium oxide.
$4Al(s) + 3O_2(g) \rightarrow 2Al_2O_3(s)$
- Oxidation: Aluminium is oxidised (oxidation number 0 → +3).
- Reduction: Oxygen is reduced (oxidation number 0 → -2).
Reduction of Oxides (Metallurgy):
The extraction of metals from their ores often involves the reduction of metal oxides. This is a crucial process in metallurgy. A more reactive element is used as a reducing agent to remove the oxygen.
Examples:
- Reduction with Carbon: Used in a blast furnace to extract iron.
$Fe_2O_3(s) + 3CO(g) \rightarrow 2Fe(l) + 3CO_2(g)$
In this case, carbon monoxide is the reducing agent. Iron is reduced (from +3 to 0) and carbon is oxidised (from +2 to +4). - Reduction with a More Reactive Metal: The Thermite reaction.
$Fe_2O_3(s) + 2Al(s) \rightarrow 2Fe(l) + Al_2O_3(s)$
Here, aluminium is the reducing agent.
Solved Examples:
- What happens to the metal during the formation of a metallic oxide?
Solution: It is oxidised (loses electrons). - What is the role of oxygen in the formation of a metallic oxide?
Solution: It is the oxidising agent (it is reduced). - In the extraction of iron, what is the reducing agent in the blast furnace?
Solution: Carbon monoxide ($CO$). - Write a balanced equation for the burning of calcium in air.
Solution: $2Ca(s) + O_2(g) \rightarrow 2CaO(s)$. - What is the change in oxidation number for iron in the reaction $Fe_2O_3 + 3CO \rightarrow 2Fe + 3CO_2$?
Solution: It decreases from +3 to 0. - Can copper be used to reduce zinc oxide ($ZnO$)?
Solution: No, because copper is less reactive than zinc. - What is the reducing agent in the Thermite reaction, $Fe_2O_3 + 2Al \rightarrow 2Fe + Al_2O_3$?
Solution: Aluminium ($Al$). - The extraction of a metal from its oxide ore is what type of process for the metal?
Solution: A reduction process. - Name the products when titanium(IV) oxide ($TiO_2$) is reduced by magnesium.
Solution: Titanium ($Ti$) and magnesium oxide ($MgO$). - Write the balanced equation for the reduction of lead(II) oxide ($PbO$) by carbon.
Solution: $2PbO(s) + C(s) \rightarrow 2Pb(l) + CO_2(g)$.