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HOCl/OCl- Chemistry


(Note/discalimer: The chemical formulas shown in these pages are a "chemical shorthand" developed by chemists to represent our understanding of what is actually going on in the water. The actual processes are more complicated, involving other compounds than distilled water, and pure elements and compounds. Also, subscripting, superscripting, and some symbols are difficult to render in HTML and are therefore not entirely accurate on this page...)


When pure chlorine is added to water, it forms hypochlorous acid and hydrochloric acid:

Cl2 + H2O --> HOCl + HCl
chlorine plus water forms hypochlorous acid plus hydrochloric acid

Hypochlorous acid (HOCl) is the stronger form of free chlorine, and hydrochloric acid (HCl) lowers pH and alkalinity.

Hypochlorous acid further dissociates to hypochlorite ion (OCl–, the weaker form of free chlorine) and free hydrogen (H+):

HOCl --> OCl– + H+

This dissociation is reversable, and pH driven. As HOCl is used to kill algae, or as it evaporates, OCl– shifts back to HOCl to maintain the pH–mandated equilibrium. Representative pH levels and their corresponding percentages of HOCl and OCl– are:



 % as HOCl

% as OCl-























The full equation may be represented like this:

Cl2 + H2O --> HOCl + HCl
HOCl --> OCl– + H+

HOCl is, of course, the “active ingredient”. The OCl– is a bank, or reservoir of less active chlorine.

A common pool industry myth is that when HOCl is used in a pool only OCl– remains. In reality, when HOCl is used, OCl– immediately converts back to HOCl to maintain the percentage division mandated by the pH.
Although the actual reactions in water may be complex, a few illustrations in simplified language may help to describe what happens.
For example, if a pool contained 3 ppm total chlorine at a pH of 7.5, there would be about 1.5 ppm HOCl and 1.5 OCl–. If 1 ppm chlorine demand is introduced into the water, the myth would have us believe that as the demand is met, the HOCl is lowered to 0.5 ppm with the OCl– remaining at 1.5 ppm. Assuming the pH to be unchanged, what actually happens is that the total chlorine is lowered to 2 ppm, the HOCl is lowered to 1 ppm, and the OCl– is lowered to 1 ppm. This happens even though it may have actually been only the faster and more potent HOCl that satisfied the chlorine demand. The subsequent shift of 0.5 OCl– to HOCl is virtually instantaneous.

If, under the same circumstances, 2 ppm chlorine demand were introduced to the pool, the 3 ppm total chlorine is still sufficient to satisfy the demand. Since part of the OCl– may be used in the process, the kill rate may be slightly slower, but the demand is met. The remaining 1 ppm of chlorine in the pool shifts almost immediately to 0.5 HOCl and 0.5 OCl–.

Another myth or misunderstanding is that at higher pH levels chlorine is less effective. Actually, pH does not so much control chlorine’s effectiveness, but rather the percentage of chlorine in its most effective form (HOCl). Thus, if the target HOCl level in a pool is 2 ppm, examples of how that target level can be met include maintaining 3 ppm total chlorine (pH 7.2, HOCl 2.0, OCl– 1.0), 4 ppm total chlorine (pH 7.5, HOCl 2.0, OCl– 2.0), or 6 ppm total chlorine (pH 7.8, HOCl 2.0, OCl– 4.0).

In reference to the speed of the conversion of HOCl to OCl-, the following quaotes are from The Handbook of Chlorination by G. Clifford White (1972), considered the authority on water chlorination:

"The rate of dissociation of HOCl is so rapid that equilibrium between HOCl and the OCl- ion is maintained, even though the HOCl is being continuously used. For example, if water containing one ppm of titratable free available chlorine residual has introduced into it a reducing agent that consumes, say, 50 percent of the hypochlorous acid, the remaining residual will redistribute itself between HOCl and OCl- ion according to the values shown in Table 4-5 and Fig. 4-1." (page 217)

"In accordance with LaChatelier's principle, as soon as the HOCl in Eq. 8-6 is used up, more HOCl is immediately formed from the OCl- ion and the hydrogen ion to maintain the chemical equilibrium in Eq. 8-6." (page 475)

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