Redox Chemistry in Water-Treatment Disinfection
Chlorine, ozone and chlorine dioxide disinfect water for the same basic reason: they are all strong oxidising agents, and oxidising a microorganism's cell components is what disables it. Which one a treatment plant reaches for depends on more than tradition — it is grounded in the same standard reduction potentials you already use to predict whether one redox reaction will run ahead of another. This article uses real electrochemical data to rank these disinfectants by oxidising strength, works through the free-energy calculation behind that ranking, and is honest about the one thing E° values alone cannot tell you.
The formulas
Standard reduction potentials for common disinfectants
| Half-reaction | E° (V) |
|---|---|
| O₃ + 2H⁺ + 2e⁻ → O₂ + H₂O | +2.07 |
| HOCl + H⁺ + 2e⁻ → Cl⁻ + H₂O | +1.49 |
| Cl₂(g) + 2e⁻ → 2Cl⁻ | +1.36 |
| ClO₂ + e⁻ → ClO₂⁻ | +0.95 |
A higher E° means the species has a stronger thermodynamic pull to accept electrons — it is a stronger oxidising agent. Read straight off this table, ozone is the strongest of the four common disinfectants and chlorine dioxide the weakest, with chlorine (as HOCl, its active form in water) sitting in between.
Worked example 1 — comparing free energies
Using the standard hydrogen electrode as the implicit reference (E° = 0), find ΔG° for the reduction half-reactions of ozone and of chlorine gas, and compare them.
Ozone (n = 2): ΔG° = −nFE° = −(2)(96 485)(2.07)
= −399 527 J/mol ≈ −399.5 kJ/mol
Chlorine (n = 2): ΔG° = −(2)(96 485)(1.36)
= −262 439 J/mol ≈ −262.4 kJ/mol
Ozone's reduction has the more negative ΔG° — a larger thermodynamic driving force — confirming numerically what the E° comparison already suggested: ozone is the stronger oxidiser of the two.
Worked example 2 — can ozone oxidise chloride to chlorine?
Build a full cell from the O₃/O₂ and Cl₂/Cl⁻ half-reactions and check whether ozone can, in principle, oxidise chloride ion in water to chlorine gas.
Higher E° is reduced (cathode): O₃ + 2H⁺ + 2e⁻ → O₂ + H₂O, E° = +2.07 V
Lower E° is oxidised (anode, reverse of its reduction): 2Cl⁻ → Cl₂ + 2e⁻
E°cell = E°cathode − E°anode(reduction) = 2.07 − 1.36 = +0.71 V
Positive E°cell means the reaction is spontaneous under standard conditions. Confirm with ΔG°: ΔG° = −nFE° = −(2)(96 485)(0.71) = −137 009 J/mol ≈ −137.0 kJ/mol
Yes — thermodynamically, ozone is capable of oxidising chloride to chlorine, which is one reason ozone and chlorine-based disinfection are sometimes used together rather than treated as interchangeable alternatives.
Where this is actually used
Chlorination remains the most widely used disinfection method because it is inexpensive and, crucially, it leaves a measurable residual — some active chlorine remains dissolved in the water as it travels through the distribution network, continuing to suppress microbial regrowth all the way to the tap. Ozone is a stronger, faster-acting oxidiser and is effective against some pathogens (such as Cryptosporidium oocysts) that are comparatively resistant to chlorine at achievable doses, but ozone decomposes quickly in water and leaves essentially no residual protection, so plants using ozone for primary disinfection commonly follow it with a small chlorine dose specifically to maintain residual protection in distribution. Chlorine dioxide, despite its lower E° in the table above, is valued for a different practical reason: it reacts less readily with the dissolved organic matter naturally present in water, which reduces the formation of chlorinated disinfection by-products (trihalomethanes) that arise when chlorine reacts with that same organic matter.
Common mistakes that cost marks
- Assuming a higher E° always means "the better disinfectant in practice." E° measures thermodynamic oxidising strength, not real-world disinfection performance, which also depends on dose, contact time (often summarised as the "CT value" — concentration multiplied by contact time), pathogen resistance and whether a disinfectant leaves a useful residual. A stronger oxidiser is not automatically the right engineering choice.
- Forgetting to reverse the sign when a half-reaction is run as an oxidation. Tables always list reduction potentials; the species being oxidised at the anode uses its reduction E° with a sign flip when computing E°cell, not the table value directly.
- Getting the sign of ΔG° = −nFE° wrong. A positive E°cell must give a negative (spontaneous) ΔG°; a positive ΔG° from a positive E° signals a sign error, not a genuinely non-spontaneous reaction.
- Ignoring H⁺ and H₂O when balancing these half-reactions. Several of the disinfectant half-reactions above involve H⁺ explicitly, and their potentials are therefore pH-dependent in practice — a fact worth stating rather than silently ignoring when discussing real water chemistry.
Exam relevance
| Exam | Typical use |
|---|---|
| IIT-JAM / CUET-PG Physical Chemistry | Standard reduction potentials, ΔG°–E° conversions, cell-potential numericals |
| GATE Chemistry | Electrochemistry, predicting spontaneity of redox reactions |
| CSIR-NET Physical Chemistry | Electrochemical series applications, environmental redox chemistry |
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