Electrolysis

Electrolysis: The decomposition of a chemical compound (in its molten state or aqueous solution) by passing a direct electric current through it.

Electrolyte: A substance that conducts electricity in its molten or aqueous state and is simultaneously decomposed. Examples: Molten NaCl, Aqueous H₂SO₄.

Non-electrolyte: A substance that does not conduct electricity in any state. Examples: Sugar (C₁₂H₂₂O₁₁), Alcohol (C₂H₅OH).

Electrodes: Conductors (usually metallic or graphite) through which electric current enters and leaves the electrolyte.

Anode: Connected to the positive terminal of the battery; site of oxidation (anions move here). Electrons are lost here.

Cathode: Connected to the negative terminal of the battery; site of reduction (cations move here). Electrons are gained here.

Ions: Charged particles formed when atoms gain or lose electrons. Cations (positive ions) move to cathode, Anions (negative ions) move to anode.

The Redox Nature: Electrolysis is fundamentally a redox (reduction-oxidation) reaction driven by electrical energy.

Oxidation at Anode: Loss of electrons occurs at the anode. Anions discharge by losing electrons. Example: 2Cl⁻ → Cl₂ + 2e⁻

Reduction at Cathode: Gain of electrons occurs at the cathode. Cations discharge by gaining electrons. Example: Cu²⁺ + 2e⁻ → Cu

Ion Movement: In the electrolyte, cations migrate towards the cathode (negative electrode) and anions migrate towards the anode (positive electrode).

Discharge of Ions: The process where ions gain or lose electrons at the electrodes to form neutral atoms or molecules.

Electrochemical Series for Cations (Decreasing order of discharge preference):

Ag⁺ > Hg²⁺ > Cu²⁺ > H⁺ > Pb²⁺ > Sn²⁺ > Ni²⁺ > Zn²⁺ > Al³⁺ > Mg²⁺ > Na⁺ > Ca²⁺ > K⁺

Electrochemical Series for Anions (Decreasing order of discharge preference):

SO₄²⁻ < NO₃⁻ < OH⁻ < Cl⁻ < Br⁻ < I⁻

Factors Affecting Selective Discharge:

1. Position in Electrochemical Series: Lower in series = easier to discharge

2. Concentration: Higher concentration favors discharge (e.g., concentrated Cl⁻ over OH⁻)

3. Nature of Electrode: Active electrodes participate in reactions; Inert electrodes (Pt, C) don't react

Electrolysis of Molten Lead Bromide (PbBr₂):

• At Cathode: Pb²⁺ + 2e⁻ → Pb (lead metal deposited)

• At Anode: 2Br⁻ → Br₂ + 2e⁻ (bromine gas evolved)

• Overall: PbBr₂ → Pb + Br₂

Electrolysis of Acidified Water (H₂SO₄ added):

• At Cathode: 2H⁺ + 2e⁻ → H₂ (hydrogen gas)

• At Anode: 4OH⁻ → 2H₂O + O₂ + 4e⁻ (oxygen gas)

• Volume ratio H₂ : O₂ = 2 : 1

• Acid is added to increase conductivity and provide H⁺ ions

Electrolysis of Aqueous Copper Sulphate (CuSO₄):

• With Inert Electrodes (Pt/Graphite):

- Cathode: Cu²⁺ + 2e⁻ → Cu (copper deposited)

- Anode: 4OH⁻ → 2H₂O + O₂ + 4e⁻ (oxygen evolved)

• With Active Copper Electrodes:

- Cathode: Cu²⁺ + 2e⁻ → Cu (pure copper deposited)

- Anode: Cu → Cu²⁺ + 2e⁻ (impure copper dissolves)

- Used for electro-refining of copper

Electroplating: Coating one metal with another using electrolysis.

• Object to be plated = Cathode

• Plating metal = Anode

• Electrolyte contains ions of plating metal

• Uses: Silver plating, Gold plating, Chromium plating

• Advantages: Corrosion resistance, improved appearance, durability

Electro-refining: Purification of metals using electrolysis.

• Impure metal = Anode (dissolves)

• Pure metal = Cathode (deposits)

• Example: Copper refining using CuSO₄ solution

• Impurities settle as anode mud

Industrial Applications:

• Production of hydrogen and oxygen from water

• Extraction of reactive metals (Na, Mg, Al)

• Production of chlorine and sodium hydroxide

• Electroplating industry