Cyanuric Acid, Chlorine, pH

            Cyanuric Acid


            Cyanuric Acid is a very weak organic acid with the following chemical structure: [C(O)NH]3

            It is triazine, consisting of a ring of 3 Cyanic Acid [CONH] molecules. The hydrogen atoms shift positions, causing the molecule to exist in 2 different structures (tautomers) that readily interconvert:


            1200px-Cyanuric_acid.png

            Cyanuric acid is a component of many stabilized chlorine compounds used in pool sanitation.


            “Trichlor” (Trichlor-s-triazinetrione/trichloroisocyanuric acid)

            “Trichlor” is most commonly used as slow dissolving chlorine tablets. It has the most available chlorine by weight and has the following chemical structure and formula:


            Trichloorisocyanuurzuur.png


            [C(O)NCl]3


            “Dichlor” (Dichloro-s-triazinetrione/dichloroisocyanuric acid)

            “Dichlor” is a quick dissolving chlorine, typically used for “shock” treatments (where enough chlorine is added to raise levels to 10 PPM or more). It has the following chemical structure and formula:


            Troclosene.svg.png

            [C(O)NCl]2[C(O)NH]


            How Chlorinated Cyanurates Release Chlorine


            The release of chlorine by the addition of "Trichlor" can be looked at as a multistage process. Each reaction exists in an equilibrium, influenced primarily by the pH of the water. The overall reaction is as follows:


            Trichlorisocyanuric Acid ←→ Dichlorisocyanuric Acid + Chloride ←→ Monochlorisocyanuric Acid + Chloride ←→ Cyanuric Acid

            Addition of any of the listed compounds will cause a shift in equilibrium. For example, the addition of Cyanuric Acid will cause the reaction to shift more to the left. However, it takes more and more chemical energy to shift the reaction through subsequent equilibrium points.

            Influence of pH

            Looking at each reaction, we see that the formation of Hypochlorites is heavily influenced by pH; the higher the pH the more available hydroxides as a reactant, shifting the equilibrium to the right. Inversely, the lower the pH, the more the reaction shifts to the left forming more chlorinated cyanurates.


            1. [C(O)NCl]3 + OH- ←→ [C(O)NCl]2[C(O)NH] + OCl-

            2. [C(O)NCl]2[C(O)NH]  + OH- ←→ [C(O)NH]2[C(O)NCl]  + OCl-

            3. [C(O)NH]2[C(O)NCl]  + OH- ←→ [C(O)NH]3 + OCl-


            pK values of Chlorinated Cyanurates

            Various chlorinated cyanurates have different dissociation constants, influencing the concentration of reactants. Looking at pK values which are derived from dissociation constants make it relatively easy to compare how these compounds dissociate. Where pK = pH, we can assume that 50% of the reaction exists as the chlorinated cyanurate, while the remaining exists as hypochlorites.


            For example, Trichlor has a very low pK value, indicating that it is a strong acid. This means that it fully dissociates in water. Although it is technically an equilibrium reaction and in theory could exist in solution with a low enough pH, in practice Trichlor does not exist in solution.


            Below are reference pK values for each of the above equilibrium reactions:


            1. [C(O)NCl]3 + OH- ←→ [C(O)NCl]2[C(O)NH] + OCl- (pK = <0.3)

            2. [C(O)NCl]2[C(O)NH]  + OH- ←→ [C(O)NH]2[C(O)NCl]  + OCl-  (pK = 3.1)

            3. [C(O)NH]2[C(O)NCl]  + OH- ←→ [C(O)NH]3 + OCl-  (pK = 4.1)


            Dichlor and Monochlor isocyanurates are considered relatively weak acids, with some degree existing in solution. Based on pK values, we can determine that these weak acids exist in 50% concentration where pK = pH. This means Dichlor releases 50% hypochlorite ions at a pH of 3.1, while Monochlor releases 50% at a pH of 4.1. As pH rises, so does the amount of available hypochlorite.


            Cyanuric Acid Oxidation by Hypochlorites

            Hypochlorites are able to oxidize cyanuric acid, resulting in their decomposition and their reactants gassing out of solution.


            2[C(O)NH]3 + 9 OCl- ←→ 3N2 + 6CO2 + 9Cl- + 3H2O


            While also an equilibrium reaction, both carbon dioxide and nitrogen gas are released from the water resulting in a very linear reaction. In theory, the addition of the reactants could result in the formation of Cyanuric Acid, however this is not observed in pool environments due to the lack of the required atmospheric pressure to lead to the reaction in any meaningful amount.


            However, this does outline how cyanuric acid decomposes overtime. This is slow due to the required high ratio of hypochlorites to cyanuric acid. For example, 5 ppm of hypochlorites (shocking with cyanuric acid free chlorine to achieve 10 ppm total chlorine at around a pH of 7.5) will only reduce Cyanuric Acid 1 ppm.


            Implications of Balancing


            Chlorine “Lock”

            Based on the above chemistry, we can conclude that the phenomenon of Chlorine “lock” (whereby there is no available chlorine due to high Cyanuric Acid levels) is largely a myth. However, it is undeniable that cyanuric acid does have some degree of impact. Yet this impact is not only due to the amount of cyanuric acid present, but also the pH maintained (both of which affect equilibrium).


            pH Ranges

            Most pools are maintained in the pH range of 7.2 to 7.6 according to the Langelier Saturation Index. This range is favored due to the increased presence of hypochlorous acid which is well documented:


            HClO ⇌ ClO + H+  (pKa = 7.53)


            At a pH of 7.53, 50% exists at Hypochlorous Acid. Just as in any other equilibrium reaction, the reaction will shift as amounts of reactants change; consuming hypochlorous acid will result in hypochlorites reacting to form more hypochlorous acid to maintain equilibrium.


            However, the lower the pH the more chlorinated cyanurates exist in equilibrium as well. This means that while more hypochlorous acid will exist at a lower pH, more chlorine will be “unavailable” at any single time in the form of chlorinated cyanurates.


            Combining these two factors creates a case for adjusting pH values upwards to maintain higher total available chlorine; cyanuric acid’s impact is reduced as more hypochlorites are available to form hypochlorous acid as equilibrium dictates (i.e. through the consumption of hypochlorous acid). In addition, this can marginally accelerate the decomposition of chlorinated cyanurates through oxidation of hypochlorites.


            Conclusion


            While cyanuric acid levels do affect the availability of hypochlorous acid, so does pH. Maintaining a higher pH, as stipulated in the Hamilton IndexTM of 7.8 to 8.2, can help mitigate the reaction of cyanuric acid and available chlorine by shifting equilibrium points.

            Regardless of the pH, however, there is always a portion of chlorine available, and as that chlorine is consumed more chlorine will be made available by mechanisms of equilibrium.




            Updated: 07 May 2018 10:34 AM
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