Products of Electrolysis

Electrolysis is the process of using electrical energy to drive a non-spontaneous chemical reaction, forcing decomposition of an ionic compound. When electrodes are connected to a power supply and immersed in a molten salt or aqueous ionic solution, ions migrate toward oppositely charged electrodes: cations travel to the cathode (negative electrode) where they gain electrons (reduction), and anions travel to the anode (positive electrode) where they lose electrons (oxidation). The products formed depend critically on the state of the electrolyte — whether it is molten or dissolved in water. In a molten salt, only the ions of that salt are present, so the outcome is straightforward. However, in aqueous solution, water molecules introduce competing species: H⁺ and OH⁻ ions. This competition means that more than one species may be capable of being discharged at each electrode, and the actual product depends on which ion is preferentially discharged. The preferential discharge theory (also called the selective discharge theory) provides the framework chemists use to predict which ion will be oxidised or reduced when multiple candidates are present. Sodium chloride (NaCl) is the classic electrolyte used to illustrate these principles because its electrolysis products differ dramatically between the molten and aqueous states. Electrolysis of molten NaCl produces sodium metal and chlorine gas — the basis of the industrial Downs process — whereas electrolysis of aqueous NaCl (brine) produces hydrogen gas, chlorine gas, and sodium hydroxide solution, which is the basis of the chlor-alkali industry. Understanding why these products form requires a careful application of preferential discharge theory.

Key takeaways

Identify All Ions Present
Apply Preferential Discharge Rules at the Cathode
Apply Preferential Discharge Rules at the Anode
Effect of Electrode Material
Summary — Molten vs Aqueous NaCl

Identify All Ions Present

Before predicting products, list every ionic species in the electrolyte. For molten NaCl the only species are Na⁺ and Cl⁻. For aqueous NaCl (brine), four species are present: Na⁺(aq), Cl⁻(aq), H⁺(aq) (from water autoionisation, concentration ~10⁻⁷ mol dm⁻³ in neutral solution), and OH⁻(aq). Water itself can also be directly oxidised or reduced at the electrode surface, effectively acting as a source of H⁺ and OH⁻.

Apply Preferential Discharge Rules at the Cathode

At the cathode, reduction occurs. When competing cations or water are present, the species with the less negative (higher) standard electrode potential, E°, is preferentially discharged — it is easier to reduce. Key standard reduction potentials: Na⁺(aq) + e⁻ → Na(s), E° = −2.71 V; 2H⁺(aq) + 2e⁻ → H₂(g), E° = 0.00 V; 2H₂O(l) + 2e⁻ → H₂(g) + 2OH⁻(aq), E° = −0.83 V. In aqueous NaCl, H⁺/H₂O is far easier to reduce than Na⁺, so hydrogen gas is produced at the cathode. In molten NaCl, no water is present, so Na⁺ is the only option and sodium metal is deposited.

Apply Preferential Discharge Rules at the Anode

At the anode, oxidation occurs. The species with the less positive (lower) standard electrode potential for oxidation — equivalently, the more negative standard reduction potential — is preferentially discharged. Relevant half-equations: 2Cl⁻(aq) → Cl₂(g) + 2e⁻, E°(reduction for reverse) = +1.36 V; 4OH⁻(aq) → O₂(g) + 2H₂O(l) + 4e⁻, E°(reduction for reverse) = +0.40 V; 2H₂O(l) → O₂(g) + 4H⁺(aq) + 4e⁻, E°(reduction for reverse) = +1.23 V. In theory OH⁻/H₂O is easier to oxidise than Cl⁻, predicting oxygen. However, at high Cl⁻ concentration (concentrated brine), Cl⁻ is discharged preferentially due to a kinetic effect called overpotential: oxygen evolution has a higher overpotential at most electrode surfaces, raising its effective discharge potential above that of chlorine. Therefore concentrated brine gives Cl₂ at the anode, while dilute brine gives O₂.

Effect of Electrode Material

Electrode material also influences products. Inert electrodes (platinum, graphite) do not participate in reactions. Active electrodes (copper) can dissolve at the anode: Cu(s) → Cu²⁺(aq) + 2e⁻ (E° = +0.34 V), which is easier than oxidising water or Cl⁻, so copper is preferentially oxidised. This is exploited in the electrorefining of copper.

Summary — Molten vs Aqueous NaCl

Molten NaCl: Cathode: Na⁺ + e⁻ → Na; Anode: 2Cl⁻ → Cl₂ + 2e⁻. Overall: 2NaCl(l) → 2Na(l) + Cl₂(g). Aqueous NaCl (concentrated): Cathode: 2H₂O + 2e⁻ → H₂ + 2OH⁻ (or 2H⁺ + 2e⁻ → H₂); Anode: 2Cl⁻ → Cl₂ + 2e⁻. The Na⁺ and OH⁻ ions remain in solution, forming sodium hydroxide. The chlor-alkali process produces three valuable industrial chemicals: Cl₂, H₂, and NaOH.

The Key Equation

\ce{2NaCl(l) -> 2Na(l) + Cl2(g)}

\ce{NaCl(l)} — Molten sodium chloride (—): The electrolyte; fully ionised as Na⁺ and Cl⁻ in the molten state, providing charge carriers.
\ce{Na(l)} — Liquid sodium (—): Produced at the cathode by reduction: Na⁺ + e⁻ → Na. Sodium is liquid at the operating temperature of ~600 °C.
\ce{Cl2(g)} — Chlorine gas (—): Produced at the anode by oxidation: 2Cl⁻ → Cl₂ + 2e⁻. A yellow-green toxic gas collected above the anode.

Worked Example

Problem. A student electrolyses concentrated aqueous sodium chloride solution using inert graphite electrodes. (a) State the product formed at each electrode. (b) Write the half-equation for each electrode reaction. (c) If the student then dilutes the solution significantly before repeating the experiment, which product at the anode would change, and why?

\ce{2Cl^{-}(aq) -> Cl2(g) + 2e^{-}} \quad \text{(concentrated)}\quad\longrightarrow\quad \ce{4OH^{-}(aq) -> O2(g) + 2H2O(l) + 4e^{-}} \quad \text{(dilute)}

Answer. Concentrated brine: cathode → hydrogen gas; anode → chlorine gas. On dilution, the anode product changes to oxygen gas because the low Cl⁻ concentration no longer overcomes the thermodynamic and kinetic preference for OH⁻ oxidation.

Figure

Diagram

Origins · 1 minMichael Faraday and how this was discovered

Michael Faraday

1791–1867 · British

Faraday established the quantitative laws of electrolysis in the 1830s. His First Law states that the mass of substance deposited at an electrode is proportional to the quantity of electrical charge passed (Q = It). His Second Law states that the masses of different substances deposited by the same charge are proportional to their equivalent masses (molar mass ÷ number of electrons transferred). He also introduced the fundamental terminology still used today: electrode, anode, cathode, electrolyte, ion, cation, and anion — coined with the help of philosopher William Whewell. Faraday had almost no formal education beyond basic reading and writing, having been apprenticed as a bookbinder at age 14. His entire scientific career grew from the books he read while binding them and attending Humphry Davy's public lectures at the Royal Institution — which he also eventually led as director.

Half-Equations and Oxidation States — Writing balanced half-equations for oxidation and reduction is the core skill needed to describe what happens at each electrode during electrolysis.
Standard Electrode Potentials (E°) — Preferential discharge theory relies on comparing E° values to determine which ion is more easily reduced at the cathode or oxidised at the anode.
Faraday's Laws of Electrolysis — Once you know which products form, Faraday's Laws allow you to calculate the mass or volume of product deposited from a measured quantity of electric charge.
Electrochemical Cells and EMF — Electrolytic cells are the reverse of galvanic cells; understanding spontaneous redox in galvanic cells deepens understanding of why electrical energy must be supplied in electrolysis.
The Chlor-Alkali Process — The industrial-scale electrolysis of brine is a direct application of the products of electrolysis, producing chlorine, hydrogen, and sodium hydroxide for global manufacturing.