The Applications of Electrolysis

This section explains the applications of electrolysis covering, what is electrolysis, electrolysis of molten ionic compounds, using electrolysis to extract metals and electrolysis of aqueous solutions. 

What is Electrolysis?

Electrolysis is the process of using electricity to cause a chemical reaction. It involves the decomposition of an ionic compound into its elements or simpler compounds when an electric current is passed through it. This process occurs in an electrolyte, which is a substance that conducts electricity when molten or dissolved in water.

During electrolysis, positive ions move towards the negative electrode (cathode), where they gain electrons (reduction), and negative ions move towards the positive electrode (anode), where they lose electrons (oxidation). The electric current provides the energy needed to drive these reactions.

Electrolysis of Molten Ionic Compounds

When an ionic compound is molten (i.e., it is heated until it melts), its ions are free to move. Electrolysis can occur in molten ionic compounds, as the ions are mobile enough to carry the electric current.

Example: Electrolysis of Molten Lead(II) Bromide (PbBr₂)

When lead(II) bromide (PbBr₂) is melted, it dissociates into lead ions (Pb²⁺) and bromide ions (Br⁻). When an electric current is passed through the molten lead(II) bromide, the lead ions are reduced at the cathode, and the bromide ions are oxidised at the anode.

At the Cathode (Reduction): Lead ions (Pb²⁺) gain electrons to form solid lead (Pb).

$$Pb^{2+} (l) + 2e^- \rightarrow Pb (s)$$

At the Anode (Oxidation): Bromide ions (Br⁻) lose electrons to form bromine gas (Br₂).

$$2Br^- (l) \rightarrow Br_2 (g) + 2e^-$$ 

Overall, the reaction produces lead metal at the cathode and bromine gas at the anode.

Using Electrolysis to Extract Metals

Electrolysis is commonly used to extract metals from their ores, particularly metals that are too reactive to be extracted by traditional methods such as smelting. Metals like aluminium and sodium are extracted using electrolysis.

Example: Electrolysis of Aluminium Oxide (Al₂O₃)

Aluminium is extracted from aluminium oxide (Al₂O₃) using electrolysis. Aluminium oxide is dissolved in molten cryolite (Na₃AlF₆) to lower its melting point. The electrolysis process involves the following reactions:

At the Cathode (Reduction): Aluminium ions (Al³⁺) gain electrons to form solid aluminium.

$$Al^{3+} (l) + 3e^- \rightarrow Al (l)$$ 

At the Anode (Oxidation): Oxide ions (O²⁻) lose electrons to form oxygen gas.

$$2O^{2-} (l) \rightarrow O_2 (g) + 4e^-$$ 

In this process, aluminium metal is deposited at the cathode, and oxygen gas is released at the anode. This is an important industrial process because aluminium is a lightweight, versatile metal used in many industries.

Electrolysis of Aqueous Solutions

When electrolysis is carried out on aqueous solutions, the ions from both the electrolyte and water (H₂O) are involved. Water itself ionises to form hydrogen ions (H⁺) and hydroxide ions (OH⁻), which also participate in the reactions at the electrodes.

For example, in the electrolysis of sodium chloride solution (NaCl), water’s ions can also be involved in the reaction:

Example: Electrolysis of Sodium Chloride Solution (Brine)

At the Cathode (Reduction): Hydrogen ions (H⁺) from water are reduced to form hydrogen gas (H₂).

$$2H^+ (aq) + 2e^- \rightarrow H_2 (g)$$

At the Anode (Oxidation): Chloride ions (Cl⁻) are oxidised to form chlorine gas (Cl₂).

$$2Cl^- (aq) \rightarrow Cl_2 (g) + 2e^-$$ 

In this process, chlorine gas is produced at the anode, and hydrogen gas is produced at the cathode. The reaction that takes place depends on the relative ease of oxidation and reduction for the ions involved. In the case of brine, chloride ions are easier to oxidise than hydroxide ions from water, so chlorine gas is produced at the anode.

Half Equations

A half equation shows either the oxidation or reduction half of the electrolysis reaction. It represents the change in oxidation state of a substance at a specific electrode.

At the Cathode (Reduction): The substance gains electrons. For example, in the electrolysis of molten lead(II) bromide:

$$Pb^{2+} (l) + 2e^- \rightarrow Pb (s)$$ 

This is the reduction half-equation because lead ions (Pb²⁺) gain electrons to form solid lead.

At the Anode (Oxidation): The substance loses electrons. For example, in the electrolysis of molten lead(II) bromide:

$$2Br^- (l) \rightarrow Br_2 (g) + 2e^-$$ 

This is the oxidation half-equation because bromide ions (Br⁻) lose electrons to form bromine gas.

Half equations are useful for understanding how the substances change at each electrode and help in balancing redox reactions.

Summary of Key Points:

• Electrolysis is a process where an electric current is used to drive a non-spontaneous chemical reaction.
• Electrolysis of molten ionic compounds involves the breakdown of the compound into its elements.
• Electrolysis is used to extract metals such as aluminium, especially for metals too reactive for extraction by heating.
• Electrolysis of aqueous solutions involves reactions at both electrodes, with water’s ions also being involved.
• Half equations show the oxidation and reduction processes that occur at the electrodes during electrolysis.

This process is key for understanding the practical applications of electrolysis in industries, particularly in metal extraction and the production of chemicals like chlorine and hydrogen.