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!


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!


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