For reference, "electrochlorination" is just the electrolysis of a simple solution of water and table salt (NaCl) followed by reintroducing the chlorine gas produced.

But, the Wikipedia article on "electrochlorination" claims

The product of this process, sodium hypochlorite, contains 0.7% to 1% chlorine. Anything below the concentration of 1% chlorine is considered a non-hazardous chemical although still a very effective disinfectant. In addition, the sodium hypochlorite produced is in the pH range of 6-7.5. This means that the chemical is relatively neutral in regards to acidity or baseness. Also, at that pH range, the sodium hypochlorite is extremely stable and the electrochlorination extremely effective.

Is the yield really that small? It seems like we should be able to get much more! The electrolysis of NaCl in water $$\ce{2H+ (aq) + 2Cl- (aq) -> H2(g) + Cl2(g)}$$ upon recombining the chlorine gas $$\ce{2OH- (aq) + Cl2(g) -> OCl- (aq) + Cl- (aq) + H2O(\ell)}$$ should lead to a net increase in pH $$\ce{Cl- (aq) + H2O(\ell) -> OCl- (aq) + H2(g)}$$ as$\ce{\; OCl- (aq)\;}$is the conjugate base of a weak acid.

Consequently, it seems the resulting solution should have a much greater percent active chlorine, due to the hypochlorous stabilization effect of higher pH -- which would only get even higher as the electrolysis (given sufficient external voltage) continues!

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    $\begingroup$ About 20-21% is maximum possible and usually, its done to about 10-12%. Then some vacuum distillation if to be sold as 15% or more. Please don't ask for source, I work for an Alkali company and this is what is written on the Production records $\endgroup$ Aug 28, 2022 at 4:06

2 Answers 2


A little bit of Wikipedia research reveals that

Sodium hypochlorite can also be obtained as a crystalline pentahydrate $\ce{NaOCl·5 H2O}$, which is not explosive and is much more stable than the anhydrous compound.5 The formula is sometimes given as $\ce{2NaOCl·10H2O}$. The transparent light greenish yellow orthorombic[10][11] crystals contain 44% NaOCl by weight and melt at 25−27°C. The compound decomposes rapidly at room temperature, so it must be kept under refrigeration. At lower temperatures, however, it is quite stable: reportedly only 1% decomposition after 360 days at 7°C.6[12]

So, it looks like 44% w/v would be an upper bound to any solution of NaOCl in water, regardless of whether it was obtained via electrochlorination or not. This corresponds to approximately 43.1% active chlorine.

Also, note that that the term "active chlorine" measures overall oxidative capability, not necessarily hypochlorite concentration — hypochlorites can easily decompose into chlorates without altering the greater solution's overall oxidative capabilities $$\ce{3OCl- (aq) -> ClO3- (aq) + 2Cl- (aq)}$$ as demonstrated via their molar-equivalent iodometries $$\ce{3OCl- (aq) + 6I- (aq) + 6H+ (aq) -> 3I2(s) + 3Cl- (aq) + 3H2O (\ell) \\ ClO3- (aq) + 6I- (aq) + 6H+ (aq) -> 3I2(s) + Cl- (aq) + 3H2O (\ell)}$$ so more accurate measurements of hypochlorite concentration may likely require additional techniques (such as spectroscopy, for example) as well.

Finally, due to the unstable nature of solutions with high [NaOCl(aq)] and their natural tendency to $$\ce{2OCl- (aq) -> 2Cl- (aq) + O2 (g)}$$ maximizing [NaOCl(aq)] solely via electrolysis/electrochlorination will likely require a dedicated setup that simultaneously:

  • Limits atmospheric exposure around the anode to prevent loss of chlorine gas while allowing hydrogen gas to escape by the cathode (or, alternatively, reduce atmospheric oxygen to water)
  • Keeps the solution at low temperatures (to limit various other decomposition reactions)
  • Maintains some external "tending" voltage (to preserve the Nernst equilibrium)

For practical applications, this could be a very unwieldy form of storage!


I interpret the question "Maximum possible hypochlorite concentration from electrolysis?", to imply a safe (that is, likely non-explosive) manufacture thereof. Also, as cited in this source, increased related hypochlorous acid presence (from dissolved chlorine reacting with water) promulgates accelerated chloride and chlorate formation, and can actually consume all of the hypochlorous present, thereby limiting total sodium hypochlorite creation (depending on the electrolysis cell design and operating conditions).

In respect to safety, I would comment, for the record, that the actual weak link in a high concentration chlorine electrolysis cell, with respect to safe operations, actually relates not so much to the formed aqueous sodium hypochlorite, but from dichlorine oxide, as emanating from hypochlorous acid as created from the action of electrically released chlorine interacting with water (that is, as well, increasingly being removed due to any precipitating hydrates of sodium hypochlorite).

Why my safety concerns? Some quotes from Bretherick’s Handbook of Reactive Chemical Hazards, Seventh Edition, Volume 1, available free online:

the gas [dichlorine oxide] readily explodes on rapid the end of slow thermal decomposition. Kinetic data are summarised [3]. The spontaneously explosive decomposition of the gas was studied at 42—86 C, and induction periods up to several hours were noted [4].

And further:

Several instances of containers or drums igniting or erupting on reopening after previous use have been reported, and may well have involved formation by slow hydrolysis of a relatively high concentration of dichlorine monoxide in the containers. This may have been caused to decompose from sudden exposure to light, by friction on opening the drum, or by a static spark.

I suspect the presence of sodium chloride reduces the solubility of dichlorine monoxide, further contributing to potential issues.

As such, I recommend venting, so as per the incident noted above, to lessen dangerous explosive fumes, which is likely advisable, as the electrolysis solution, itself, is caustic so that even a mild explosive event may result in injuries.

Also, based on an old text, I surmised that one should not consider working, in general, with concentrations of hypochlorous acid exceeding 20%, being described as dangerous, which I believe, would also be advisable for a safe cell operation and performance. Support is provided by US Patent 5,108,560 citing working examples of hypochlorous solutions of 15% and 20%, and further per this already alluded to work where to quote:

The rate of decomposition of hypochlorous acid has been measured in aqueous solution in the presence of much sodium hypochlorite. The rate is nearly independent of the hypochlorite concentration, and proportional to the square of the hypochlorous acid concentration.

where the disproportionation results in an even more caustic chloric acid, and therefrom sodium chlorate.


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