Electrochemistry

Overview

Electrolysis is the process of using electrical energy to break down ionic compounds into their elements. It occurs when an electric current passes through a molten ionic compound or aqueous solution containing ions. Electrolysis has important industrial applications including metal extraction, purification, and electroplating. Understanding electrode reactions and ion movement is essential for predicting products.

Principle of Electrolysis

Electrolysis requires:

• **Electrolyte**: A molten ionic compound or aqueous solution containing free-moving ions
• **Electrodes**: Two conductors (often inert materials like graphite or platinum) connected to a power supply
• **Direct current (DC)**: Electrical energy source

When current flows:

• **Positive ions (cations)** move toward the **cathode (negative electrode)**
• **Negative ions (anions)** move toward the **anode (positive electrode)**
• Chemical reactions occur at the electrodes, breaking down the compound

The electrolyte must contain free-moving ions - solid ionic compounds cannot conduct electricity because ions are locked in the lattice.

Cathode and Anode Reactions

At the CATHODE (negative electrode):

• Positive ions (cations) gain electrons
• This is **reduction** (gain of electrons)
• Metals or hydrogen are typically produced
• Example: Cu²⁺ + 2e⁻ → Cu

At the ANODE (positive electrode):

• Negative ions (anions) lose electrons
• This is **oxidation** (loss of electrons)
• Non-metals are typically produced
• Example: 2Cl⁻ → Cl₂ + 2e⁻

Memory aid: "An Ox" - Anode Oxidation

Electrolysis of Molten Ionic Compounds

When a molten (liquid) ionic compound undergoes electrolysis, only the metal and non-metal from the compound are present.

Example 1: Molten lead(II) bromide (PbBr₂)

Ions present: Pb²⁺ and Br⁻

Cathode reaction: Pb²⁺ + 2e⁻ → Pb

Product: Lead metal (grey liquid)

Anode reaction: 2Br⁻ → Br₂ + 2e⁻

Product: Bromine gas (brown vapour)

Example 2: Molten sodium chloride (NaCl)

Cathode: Na⁺ + e⁻ → Na (sodium metal)

Anode: 2Cl⁻ → Cl₂ + 2e⁻ (chlorine gas)

This is used industrially to extract sodium metal and chlorine gas.

Electrolysis of Aqueous Solutions

Aqueous solutions contain:

• Ions from the dissolved compound
• H⁺ and OH⁻ ions from water (H₂O ⇌ H⁺ + OH⁻)

Predicting products:

At the CATHODE: Choose between metal ions and H⁺ ions

• If the metal is **more reactive than hydrogen** (e.g., Na, Mg, Al), then **hydrogen gas** is produced: 2H⁺ + 2e⁻ → H₂
• If the metal is **less reactive than hydrogen** (e.g., Cu, Ag), then the **metal** is produced: Cu²⁺ + 2e⁻ → Cu

At the ANODE: Choose between non-metal ions and OH⁻ ions

• If the solution contains **halide ions** (Cl⁻, Br⁻, I⁻), the **halogen** is produced: 2Cl⁻ → Cl₂ + 2e⁻
• If **no halide ions** are present, or if sulfate/nitrate, then **oxygen gas** is produced: 4OH⁻ → O₂ + 2H₂O + 4e⁻

Example 1: Copper(II) sulfate solution (CuSO₄)

Ions present: Cu²⁺, SO₄²⁻, H⁺, OH⁻

Cathode: Cu²⁺ + 2e⁻ → Cu (copper is less reactive than hydrogen)

Product: Copper metal coating the cathode

Anode: 4OH⁻ → O₂ + 2H₂O + 4e⁻ (no halide ions present)

Product: Oxygen gas

Example 2: Sodium chloride solution (NaCl)

Ions present: Na⁺, Cl⁻, H⁺, OH⁻

Cathode: 2H⁺ + 2e⁻ → H₂ (sodium is more reactive than hydrogen)

Product: Hydrogen gas

Anode: 2Cl⁻ → Cl₂ + 2e⁻ (chloride ions present)

Product: Chlorine gas

This process produces hydrogen, chlorine, and leaves sodium hydroxide in solution - important industrial process called chlor-alkali industry.

Industrial Applications

1. Extraction of Metals

• **Aluminum extraction**: Electrolysis of molten aluminum oxide (Al₂O₃) dissolved in cryolite
• Cathode: Al³⁺ + 3e⁻ → Al
• Anode: 2O²⁻ → O₂ + 4e⁻
• Requires huge amounts of electricity, so aluminum plants are near cheap energy sources

2. Purification of Copper

• Impure copper anode, pure copper cathode, copper(II) sulfate electrolyte
• Pure copper deposits on cathode
• Impurities fall to bottom as "anode sludge"

3. Chlor-Alkali Industry

• Electrolysis of brine (concentrated NaCl solution)
• Products: chlorine (water treatment, PVC), hydrogen (fuel, ammonia), sodium hydroxide (cleaning products)

4. Electroplating

• See separate Electroplating topic for details

Energy Considerations

Electrolysis requires continuous electrical energy input because:

• Breaking ionic bonds requires energy (endothermic)
• Chemical energy is being stored in the products
• The process reverses processes that naturally release energy

This is why electrolysis is expensive, especially for reactive metals like aluminum.

Exam Tips

• Always identify which electrode is the anode (positive) and cathode (negative)
• For molten compounds, only two ions present - predict products directly
• For aqueous solutions, use reactivity series to decide cathode product; check for halides for anode product
• Write half equations showing electrons (e⁻) and balance for atoms and charge
• Remember: reduction at cathode (electrons gained), oxidation at anode (electrons lost)
• State symbols are important: (l) for molten metals, (g) for gases, (s) for solid metals deposited
• Describe observations: color changes, bubbles of gas, metal coating electrodes
• Link industrial processes to products: aluminum extraction, copper purification, chlor-alkali