The Chloralkali Process

What the chloralkali process does

Take saturated table-salt solution, drop two electrodes in it, and push direct current through. Chlorine gas bubbles off one electrode, hydrogen off the other, and the liquid left behind turns into caustic soda. Three of the twenty most-produced industrial chemicals — chlorine, hydrogen, and sodium hydroxide — fall out of a single beaker of salt water. That is the chloralkali process, and it consumes roughly 2% of the electricity generated on Earth to do it.

The overall reaction is deceptively simple:

    2 NaCl(aq)  +  2 H₂O(l)  ──electricity──→  Cl₂(g)  +  H₂(g)  +  2 NaOH(aq)

The subtlety is entirely in the details: which species get oxidized and reduced, why chlorine wins over oxygen at the anode even though it "shouldn't," and how you keep three reactive products from immediately reacting back together. Every design decision in a modern cell exists to answer one of those three questions.

The two half-reactions

Electrolysis splits into an oxidation at the anode (electrons leave the solution into the wire) and a reduction at the cathode (electrons enter the solution from the wire). In brine, several candidates compete at each electrode.

At the anode (oxidation, +): chloride loses an electron and pairs up into chlorine gas.

    2 Cl⁻  →  Cl₂  +  2 e⁻            E° = +1.36 V (vs SHE)

The competing reaction is water oxidation to oxygen:

    2 H₂O  →  O₂  +  4 H⁺  +  4 e⁻     E° = +1.23 V (vs SHE)

Notice the paradox: oxygen has the lower potential (+1.23 V < +1.36 V), so on pure thermodynamics water should oxidize first and you'd make oxygen, not chlorine. The process only works because of kinetics — more on that below.

At the cathode (reduction, −): water is reduced to hydrogen gas, releasing hydroxide.

    2 H₂O  +  2 e⁻  →  H₂  +  2 OH⁻    E° = −0.83 V (at pH 14)

Sodium ion (Na⁺/Na, E° = −2.71 V) is far too hard to reduce in water — you would need to electrolyze molten NaCl (the Downs process) to plate sodium metal. In water, hydrogen always wins at the cathode. So the sodium ions never discharge; they just migrate toward the cathode and stay in solution, pairing with the freshly made OH⁻ to give NaOH. That is where the caustic soda comes from — not from any electrode reaction on sodium, but from spectator Na⁺ meeting electrolytically generated hydroxide.

Why chlorine wins: overpotential, not thermodynamics

This is the single most important idea in the whole process. The standard potentials say oxygen should evolve first. In practice, chlorine dominates because oxygen evolution is kinetically terrible and chlorine evolution is kinetically easy.

The actual voltage you must apply to drive a gas-evolving reaction is its equilibrium potential plus an overpotential (η) — extra voltage the electrode demands to overcome the sluggish multi-step mechanism. For oxygen evolution, four electrons and four proton transfers have to be marshalled through several bound intermediates (M-OH, M-O, M-OOH), and on most electrodes η(O₂) is a punishing 0.4–0.8 V. Chlorine evolution is a fast two-electron process with almost no overpotential on the right catalyst.

    Effective anode potential = E° + η

    Chlorine:   1.36 V + ~0.05 V  ≈  1.41 V   ← wins in practice
    Oxygen:     1.23 V + ~0.60 V  ≈  1.83 V   ← "cheaper" on paper, expensive in reality

Industry stacks the deck further. The workhorse anode is a Dimensionally Stable Anode (DSA) — a titanium plate coated with ruthenium and iridium oxides (RuO₂/IrO₂/TiO₂), invented by Henri Beer and commercialized by De Nora in the late 1960s. Its coating is chosen precisely because it is a great catalyst for Cl₂ evolution and a poor one for O₂ evolution, magnifying chlorine's kinetic edge. On top of the catalyst, three operating levers push selectivity toward chlorine:

• High chloride concentration. Saturated brine (~5 mol/L Cl⁻) means far more chloride than water is available at the electrode surface, and by the Nernst equation a high [Cl⁻] lowers the effective Cl₂ potential.
• Low anolyte pH. Keeping the anode compartment acidic (pH ~2–4) makes water oxidation harder (it needs to release H⁺ into an already-acidic solution) and suppresses the chlorate side reaction.
• Moderate temperature. Around 80–90 °C balances lower cell resistance against the fact that oxygen's kinetics improve faster with temperature than chlorine's.

The net result: modern cells run at 96–98% current efficiency for chlorine — only a few percent of the current leaks into oxygen.

The membrane cell: keeping three products apart

If chlorine gas, hydrogen gas, and hot caustic soda ever mixed inside the cell, they would immediately react — chlorine and hydroxide disproportionate; chlorine and hydrogen form explosive mixtures. So the entire cell is designed around separation. The dominant modern technology is the ion-exchange membrane cell, and its heart is a thin perfluorinated polymer sheet.

The membrane is a cation-exchange membrane (Nafion-type), a Teflon-like backbone hung with fixed negative charges — sulfonate (-SO₃⁻) on the anode-facing layer and carboxylate (-COO⁻) on the cathode-facing layer. Those fixed anionic sites let positive ions hop through but electrostatically repel negative ions. In this cell:

• Na⁺ passes through from anolyte to catholyte, carrying the ionic current across the gap.
• Cl⁻ is blocked — so it can't cross into the caustic and contaminate it with salt.
• OH⁻ is blocked — so it can't back-migrate into the anolyte, where it would react with chlorine.

Trace the full circuit. At the anode, Cl⁻ is oxidized to Cl₂ (which bubbles up and out). Losing chloride leaves the anolyte with a surplus of positive charge, so Na⁺ is pushed across the membrane to restore balance. It arrives in the catholyte just as water reduction there is generating H₂ gas and dumping OH⁻ into solution. Na⁺ + OH⁻ = NaOH, which accumulates as ~32–35% caustic soda. Every electron that flows moves one Cl⁻ toward oxidation, one Na⁺ across the membrane, and helps make one OH⁻ — the three products are stitched together by a single conserved current.

    ANODE side (anolyte)         MEMBRANE          CATHODE side (catholyte)
    ─────────────────────      ═══════════       ──────────────────────────
    2 Cl⁻ → Cl₂↑ + 2 e⁻     Na⁺ ──────→ Na⁺       2 H₂O + 2 e⁻ → H₂↑ + 2 OH⁻
    (depleted brine out)     (Cl⁻, OH⁻            Na⁺ + OH⁻ → NaOH (32–35%)
                              blocked)

Reagents, materials, and conditions

The chloralkali cell is a study in choosing materials that survive one of the most corrosive environments in industry — wet chlorine on one side, hot concentrated caustic on the other.

• Feed. Saturated brine, ~300 g/L NaCl (~25 wt%). It must be ruthlessly purified first: Ca²⁺, Mg²⁺, and especially trace Ba²⁺/Sr²⁺ are precipitated out (soda ash + caustic), then a chelating ion-exchange resin polishes hardness to below ~20 ppb — because those cations would foul the membrane and destroy it.
• Anode. Titanium coated with RuO₂/IrO₂/TiO₂ (a DSA). Titanium is used because it forms a protective oxide in the chlorine environment; the noble-metal-oxide coating catalyzes Cl₂ evolution.
• Cathode. Nickel or nickel-coated steel, often with a Ni-Ru-O or Raney-nickel coating to lower the hydrogen overpotential. Nickel resists hot caustic.
• Membrane. A bilayer perfluorosulfonic/perfluorocarboxylic acid film (e.g. DuPont Nafion, AGC Flemion, Asahi Kasei Aciplex), ~0.1–0.2 mm thick.
• Conditions. ~80–90 °C, cell voltage ~3.0–3.2 V, current density 5–6 kA/m². Anolyte held acidic (HCl dosed to pH ~2–4); catholyte leaves at 32–35% NaOH and is then evaporated to the commercial 50% grade.
• Energy. ~2,100–2,500 kWh per tonne of chlorine. The thermodynamic minimum cell voltage is 2.19 V, so every volt above that (overpotentials + ohmic IR drop) is pure energy overhead.

Three cell technologies compared

Three generations of cell have made chlorine and caustic. They differ in how they keep the anode and cathode products apart — and that single design choice cascades into energy use, product purity, and environmental footprint.

	Mercury cell	Diaphragm cell	Membrane cell
Separator	None — sodium amalgam is the moving barrier	Asbestos / polymer diaphragm (porous)	Ion-exchange membrane (Na⁺-selective)
Cathode	Flowing mercury (makes Na/Hg amalgam)	Steel (direct H₂ evolution)	Nickel (direct H₂ evolution)
NaOH concentration	~50% (very pure)	~11% (salt-contaminated, needs evaporation)	~32–35% (pure, low salt)
Energy (kWh/t Cl₂)	~3,400 (highest)	~2,600	~2,100–2,500 (lowest)
Product purity	Excellent, but Hg traces	Poor (salt in caustic)	Excellent
Environmental issue	Mercury emissions	Asbestos exposure	Cleanest of the three
Status	Banned / phased out (Minamata, EU 2017)	Declining, being retrofitted	Industry standard (>90% of new capacity)

Worked example: how much salt makes a tonne of chlorine?

The three products are locked together by stoichiometry — you cannot dial one up without the others. Start from the overall equation and work through Faraday's law.

    2 NaCl + 2 H₂O → Cl₂ + H₂ + 2 NaOH

    Per mole of Cl₂:  2 mol NaCl consumed, 2 mol NaOH made, 1 mol H₂ made,
                      and 2 mol of electrons must pass (n = 2).

Molar masses: Cl₂ = 70.9 g/mol, NaCl = 58.4 g/mol, NaOH = 40.0 g/mol, H₂ = 2.02 g/mol. So per tonne (10⁶ g) of chlorine:

• Moles of Cl₂: 10⁶ / 70.9 = 1.41 × 10⁴ mol.
• NaCl needed: 2 × 1.41 × 10⁴ × 58.4 = 1.65 tonnes of salt (before any recycle of unconverted brine).
• NaOH produced: 2 × 1.41 × 10⁴ × 40.0 = 1.13 tonnes of caustic soda.
• H₂ produced: 1.41 × 10⁴ × 2.02 = 0.028 tonnes ≈ 28 kg of hydrogen.
• Charge required (Faraday's law): Q = n·F·mol = 2 × 96 485 × 1.41 × 10⁴ ≈ 2.72 × 10⁹ C, i.e. ~756,000 ampere-hours per tonne of chlorine at 100% efficiency (a few percent more in reality).

This is why chlorine and caustic are sold as a coupled ECU (electrochemical unit): roughly one tonne of Cl₂ per 1.1 tonnes of NaOH, always. If demand for PVC (a chlorine sink) outruns demand for caustic, plants can't just make more chlorine — the caustic piles up, its price collapses, and the whole ECU economics reprice. The stoichiometry above is the reason.

Side reactions and limitations

• Oxygen co-evolution. The 2–4% of current that goes to O₂ instead of Cl₂ contaminates the chlorine gas and slowly corrodes the DSA coating. It is minimized by high [Cl⁻], low pH, and a good anode coating.
• The chlorate/hypochlorite reaction. If dissolved Cl₂ meets OH⁻, it disproportionates: Cl₂ + 2 OH⁻ → Cl⁻ + OCl⁻ + H₂O, and hypochlorite can be further oxidized to chlorate, ClO₃⁻. This wastes chlorine and poisons the product. The membrane keeps OH⁻ out of the anolyte; acidifying the anolyte keeps dissolved chlorine as molecular Cl₂ rather than letting it hydrolyze. It's the exact same chemistry that makes household bleach — useful there, unwanted here.
• Membrane fouling. Trace Ca²⁺/Mg²⁺/Ba²⁺ hydroxides precipitate inside the membrane, raise its resistance, and destroy selectivity. This is why brine purity to the ppb level is non-negotiable.
• The hydrogen–chlorine hazard. H₂ and Cl₂ form an explosive mixture that detonates on light or spark. The cell design must guarantee the two gas streams never meet; a membrane pinhole is a genuine safety event.
• Energy floor. Even a perfect cell needs 2.19 V; real cells run ~3 V, so ~25% of the electrical energy is unavoidably lost to overpotentials and IR drop. Oxygen-depolarized cathodes (which reduce O₂ to OH⁻ instead of evolving H₂) can cut the voltage by ~0.8–1.0 V, but then you sacrifice the hydrogen product.

History: from Castner-Kellner to the membrane

Electrolytic chlorine dates to the 1890s. In 1892, Hamilton Castner (American) and Karl Kellner (Austrian) independently patented the mercury-cathode cell; a patent dispute merged their work into the Castner-Kellner process, which used a rocking mercury cathode to amalgamate sodium and then decomposed the amalgam with water to yield exceptionally pure caustic. Diaphragm cells (the Griesheim and later Hooker designs) grew up alongside, trading purity for freedom from mercury.

The transformative innovation came in two steps. First, in the late 1960s, Henri Beer and the Oronzio De Nora company introduced the ruthenium-oxide-coated titanium Dimensionally Stable Anode, replacing the graphite anodes that had eroded and contaminated the chlorine for seventy years. Second, in the 1970s, DuPont's Nafion perfluorinated ion-exchange membrane (originally developed for fuel cells) made the fully membrane-separated cell practical. Asahi Chemical brought the first commercial membrane chloralkali plant online in Japan in 1975. Then the Minamata Convention on Mercury (2013) made the transition mandatory: the EU set a hard deadline of December 2017 to close mercury cells, and membrane technology — cleaner and less energy-hungry — became the global default.

Where the three products go

• Chlorine → PVC and beyond. The largest chlorine sink is vinyl chloride monomer (for PVC pipe, cable, and cladding). Chlorine also disinfects drinking water and pools, bleaches paper, and is the reagent behind thousands of chlorinated organics and pharmaceuticals. Roughly 85 million tonnes of chlorine are made worldwide each year.
• Sodium hydroxide (caustic soda) → everywhere. Pulp and paper (Kraft process), soap and detergent manufacture, alumina refining (the Bayer process), textiles, drain cleaner, CO₂ scrubbing, and pH control across the entire chemical industry.
• Hydrogen → fuel and reagent. Often burned on-site for process heat, reacted with the co-product chlorine to make ultra-pure HCl, hydrogenation of fats and oils, or increasingly fed to fuel cells as a low-carbon "byproduct hydrogen."

Because the plant's economics track electricity price so tightly and the products feed so many downstream industries, chlorine output is used as a leading economic indicator for the health of the chemical sector as a whole.