Electrolytic Cells and Electrolysis

Full Notes

Introduction to Electrolytic Cells

Electrolytic cells use electricity to bring about chemical change. They are distinct from galvanic cells because the reactions occurring in electrolytic cells are non-spontaneous, requiring a source of DC power to proceed. The goal is to convert electrical energy into chemical energy.

In an electrolytic cell:

• The anode is connected to the positive terminal of the external battery and attracts anions. Anions get oxidised and release electrons that flow through an external wire to the power source.
• The cathode is connected to the negative terminal, attracting cations. Cations get reduced and gain electrons that flow from the power source to the electrode.
• Oxidation occurs at the anode, reduction at the cathode.

Faraday’s Contribution

Michael Faraday was the first to study electrolysis quantitatively. He formulated two important laws known as Faraday’s Laws of Electrolysis.

Faraday’s First Law of Electrolysis:
The amount of a substance deposited or liberated at an electrode is directly proportional to the quantity of electricity (charge) passed through the electrolyte.

Faraday’s Second Law of Electrolysis:
When the same quantity of electricity is passed through different electrolytes, the masses of substances deposited are proportional to their chemical equivalents.

Charge and the Faraday Constant

Charge (Q) is the product of current (I) in amperes and time (t) in seconds:

• Unit: Coulombs (C)

Faraday Constant (F):

• The charge of one electron is approximately 1.6021 × 10⁻¹⁹ C.
• Avogadro’s number (N_a) is 6.022 × 10²³ mol⁻¹.
• Therefore, the charge on one mole of electrons is:
  F = N_a × charge on one electron
  F ≈ 96487 C mol⁻¹ (called Faraday)
• For practical purposes, 1 F ≈ 96500 C mol⁻¹

Products of Electrolysis

The discharge of ions during electrolysis is not arbitrary – it's influenced by electrochemical principles.

Based on the Electrochemical Series

Ions lower in the electrochemical series (having higher reduction potentials) are preferentially discharged.

Example Na⁺ vs H⁺

Between Na⁺ and H⁺, H⁺ is reduced at the cathode due to higher E°.

Based on Ion Concentration

Sometimes, even if an ion is lower in the series, higher concentration can cause it to discharge preferentially.

Example Cl⁻ vs OH⁻

At high Cl⁻ concentration, Cl⁻ gets oxidised in preference to OH⁻, even though OH⁻ has a lower oxidation potential.

Based on Electrode Nature

Inert electrodes (Pt, graphite): don’t participate in reactions; only facilitate electron transfer.

Reactive electrodes (Cu, Ag): can participate by getting oxidised or depositing metal.

Molten Electrolytes

When an electrolyte is molten (melted to liquid state), we can predict the products formed at each electrode based on the type of electrolyte being used. This is because only the cations and anions of the compound are present:

• Cation is reduced at the cathode.
• Anion is oxidised at the anode.

Example Electrolysis of molten NaCl

• Cathode: Na⁺ + e⁻ → Na
• Anode: 2Cl⁻ → Cl₂ + 2e⁻

Aqueous Electrolytes

When an electrolyte is aqueous (aq), the ionic compound is dissolved in water. Now, H⁺(aq) and OH⁻(aq) ions from water are also present due to the natural ionisation of water:

2H₂O + 2e⁻ → H₂ + 2OH⁻ and 2H₂O → O₂ + 4H⁺ + 4e⁻.

As a result, water may compete with the ions from the compound at the electrodes for oxidation and reduction. We can use standard electrode potentials (E° values) or reactivity trends to predict which species is discharged.

• At the cathode:
  • H⁺(aq) ions are reduced to H₂(g) if the metal is more reactive than hydrogen (lower E° value).
  • The metal ion is reduced if it is less reactive than hydrogen (higher E° value).
• At the anode:
  • Generally, OH⁻(aq) is oxidised to form O₂ gas.
  • If halide ions (e.g. Cl⁻, Br⁻) are present, they are oxidised instead.
  • Ions like SO₄²⁻ and NO₃⁻ do not get oxidised.

Example Electrolysis of aqueous NaCl

This system contains Na⁺ and H₂O (possible reduction at the cathode) and Cl⁻ and H₂O (possible oxidation at the anode).

• At the Cathode (Reduction)
  • Competing species:
  • Na⁺ + e⁻ → Na (E° = −2.71 V)
  • 2H₂O + 2e⁻ → H₂ + 2OH⁻ (E° = −0.83 V)
  • Water is reduced, not sodium (because −0.83 V is more positive):
    Cathode reaction: 2H₂O + 2e⁻ → H₂ + 2OH⁻
• At the Anode (Oxidation)
  • Competing species:
  • 2Cl⁻ → Cl₂ + 2e⁻ (E° = +1.36 V)
  • 2H₂O → O₂ + 4H⁺ + 4e⁻ (E° = +1.23 V)
  • Even though water has a slightly lower reduction potential, Cl⁻ is preferentially oxidised in concentrated solutions of NaCl (this is due to other factors such as kinetics):
    Anode reaction: 2Cl⁻ → Cl₂ + 2e⁻

Another Example: Electrolysis of Aqueous CuSO₄

Concentration matters:

If halide concentration is very low (e.g., dilute NaCl), OH⁻ from water may be oxidised instead.

If the metal ion concentration is very low, hydrogen gas may form instead of metal at the cathode.

Summary

• Electrolytic cells require external DC power to drive non-spontaneous redox reactions.
• Anode is positive and cathode is negative; oxidation at anode and reduction at cathode.
• Products depend on E° values, concentrations, and electrode nature.
• Molten salts give products from their own ions; aqueous systems must also consider water discharge.
• Faraday’s laws connect deposited mass to total charge passed enabling quantitative electrolysis.