Comprehensive Guide to Electrolysis: Melts, Solutions, and Industrial Extractions

Introduction to Electrolysis

• Definition: Electrolysis is the electrical decomposition of ionic compounds.

• Primary Applications:

  • Metal extraction: Used to obtain reactive metals from their ores.

  • Electroplating: The process of coating the surface of a metal with another metal using electrical current.

  • Chlor-alkali process: A major industrial process that consumes approximately $1\%$ of all electricity in the United Kingdom.

The Nature of Ionic Substances

• State of Matter and Ion Mobility:

  • Solid: In a solid ionic lattice, ions are fixed in position and cannot move; therefore, solid ionic compounds do not conduct electricity for electrolysis.

  • Liquid (Melted): When an ionic compound is molten, the lattice breaks down, and the ions ($+$ and $-$) are free to move and carry charge.

  • Gas: Ions are also free in a gaseous state, though electrolysis typically involves liquids or solutions.

• Ion Separation in Solution: When an ionic substance dissolves in water ($aq$), the ions separate and become free to move throughout the solution.

Electrolysis of Molten Compounds (Melts)

Electrolysis of Molten Lead Bromide ($PbBr_2(l)$)

• Composition: Contains Lead ions ($Pb^{2+}+$) and Bromide ions ($Br^-$).

• Negative Electrode (Cathode):

  • Attraction: $Pb^{2+}$ ions are attracted to the negative electrode.

  • Process: Reduction (gain of electrons).

  • Equation: $Pb^{2+} + 2e^- \rightarrow Pb$.

  • Result: Lead metal is formed.

• Positive Electrode (Anode):

  • Attraction: $Br^-$ ions are attracted to the positive electrode.

  • Process: Oxidation (loss of electrons).

  • Equation: $2Br^- - 2e^- \rightarrow Br_2$ (alternatively written as $2Br^- \rightarrow Br_2 + 2e^-$).

  • Result: Bromine is produced.

Electrolysis of Molten Sodium Chloride ($NaCl(l)$)

• Composition: Contains Sodium ions ($Na^+$) and Chloride ions ($Cl^-$).

• Negative Electrode (Cathode):

  • Attraction: $Na^+$ ions are attracted.

  • Process: Reduction.

  • Equation: $Na^+ + e^- \rightarrow Na$.

  • Result: Sodium metal is made.

• Positive Electrode (Anode):

  • Attraction: $Cl^-$ ions are attracted.

  • Process: Oxidation.

  • Equation: $2Cl^- \rightarrow Cl_2 + 2e^-$.

  • Result: Chlorine gas is made.

Extraction of Aluminium

• Source: Aluminium is extracted from the ore Bauxite, which contains Aluminium Oxide ($Al_2O_3$).

• Element Profile: Aluminium ($Al$), Atomic Number $13$, Relative Atomic Mass $26.9815386$.

• Electrolysis of Molten Aluminium Oxide ($Al_2O_3(l)$):

  • Composition: Contains Aluminium ions ($Al^{3+}$) and Oxide ions ($O^{2-}$).

  • Negative Electrode (Cathode):

  • $Al^{3+}$ ions are attracted and gain electrons.

  • Equation: $Al^{3+} + 3e^- \rightarrow Al$.

  • Result: Aluminium metal is produced.

  • Positive Electrode (Anode):

  • $O^{2-}$ ions are attracted and lose electrons.

  • Equation: $2O^{2-} \rightarrow O_2 + 4e^-$ (Note: Transcript page 12 contains a typo stating $2O^{2-} \rightarrow Cl_2$, while page 14 provides the correct reaction producing $O_2$).

  • Result: Oxygen gas is made.

  • Graphite Anode Consumption: The anodes used are made of graphite (carbon). Because the process occurs at high temperatures, the oxygen produced reacts with the carbon anodes.

  • Reaction: $C + O_2 \rightarrow CO_2$.

  • Practical Implication: The graphite anodes burn away as Carbon Dioxide and must be replaced regularly.

Electrolysis of Aqueous Solutions

In aqueous solutions ($aq$), the competition between ions from the salt and ions from the water ($H^+$ and $OH^-$) determines which substances are discharged.

Positive Ion Discharge (Cathode Rules)

• Competition: $H^+$ vs. Metal Ions.

• Case 1: Low Reactivity Metals (e.g., $Ag^+$, $Cu^{2+}$):

  • These metals are easier to discharge than Hydrogen.

  • Result: The metal ions are discharged and the metal forms on the electrode.

• Case 2: High Reactivity Metals (e.g., $Na^+$, $Mg^{2+}+$ , $Zn^{2+}$):

  • These metals are more reactive (harder to discharge) than Hydrogen.

  • Result: Hydrogen ions ($H^+$) are discharged, forming Hydrogen gas ($H_2$).

Negative Ion Discharge (Anode Rules)

• Competition: $OH^-$ vs. Other Negative Ions.

• Case 1: Halide Ions (e.g., $Cl^-$, $Br^-$, $I^-$):

  • Halide ions are discharged.

  • Result: Halogens are formed ($Cl_2$, $Br_2$, or $I_2$).

• Case 2: Other Negative Ions (e.g., $SO_4^{2-}$, $NO_3^-$):

  • Hydroxide ions ($OH^-$) are easier to discharge than these complex ions.

  • Result: Hydroxide ions are discharged, forming Oxygen gas ($O_2$).

Aqueous Electrolysis Case Studies

1. Sodium Chloride Solution ($NaCl(aq)$)

• Ions Present: $Na^+$, $H^+$, $Cl^-$, $OH^-$.

• Positive Electrode (Anode):

  • Reaction: $2Cl^- \rightarrow Cl_2 + 2e^-$.

  • Product: Chlorine gas.

• Negative Electrode (Cathode):

  • Reaction: $2H^+ + 2e^- \rightarrow H_2$.

  • Product: Hydrogen gas.

2. Silver Nitrate Solution ($AgNO_3(aq)$)

• Ions Present: $Ag^+$, $H^+$, $NO_3^-$, $OH^-$.

• Positive Electrode (Anode):

  • Reaction: $4OH^- \rightarrow 2H_2O + O_2 + 4e^-$ (Note: Page 21 excludes electrons in its summary line).

  • Product: Oxygen gas.

• Negative Electrode (Cathode):

  • Reaction: $Ag^+ + e^- \rightarrow Ag$.

  • Product: Silver metal.

3. Potassium Sulfate Solution ($K_2SO_4(aq)$)

• Ions Present: $K^+$, $H^+$, $SO_4^{2-}$, $OH^-$.

• Positive Electrode (Anode):

  • Reaction: $4OH^- \rightarrow 2H_2O + O_2 + 4e^-$.

  • Product: Oxygen gas.

• Negative Electrode (Cathode):

  • Reaction: $2H^+ + 2e^- \rightarrow H_2$.

  • Product: Hydrogen gas.