rising temperature, pH, or added bicarbonate level, as illustrated in the accompanying graphs. The trends are illustrative for pool management, but whether scaling occurs also depends on other factors including interactions with , and other ions in the pool, as well as supersaturation effects. Scaling is commonly observed in electrolytic chlorine generators, where there is a high pH near the cathode surface and scale deposition further increases temperature. This is one reason that some pool operators prefer borate over bicarbonate as the primary pH buffer, and avoid the use of pool chemicals containing calcium. Solubility in a strong or weak acid solution Solutions of strong (HCl), moderately strong (sulfamic) or weak (acetic, citric, sorbic, lactic, phosphoric) acids are commercially available. They are commonly used as descaling agents to remove limescale deposits. The maximum amount of that can be "dissolved" by one liter of an acid solution can be calculated using the above equilibrium equations. In the case of a strong monoacid with decreasing acid concentration [A] = [], we obtain (with molar mass = 100 g/mol): {| class="wikitable" |- ! width="160" |[A] (mol/L) | 1 | 10−1 | 10−2 | 10−3 | 10−4 | 10−5 | 10−6 | 10−7 | 10−10 |- ! width="160" |Initial pH | 0.00||1.00||2.00||3.00||4.00||5.00||6.00||6.79||7.00 |- ! width="160" |Final pH | 6.75||7.25||7.75||8.14||8.25||8.26||8.26||8.26||8.27 |- ! width="160" |Dissolved (g/L of acid) | 50.0||5.00||0.514||0.0849||0.0504||0.0474||0.0471||0.0470||0.0470 |} where the initial state is the acid solution with no (not taking into account possible dissolution) and the final state is the solution with saturated . For strong acid concentrations, all species have a negligible concentration in the final state with respect to and so that the neutrality equation reduces approximately to 2[] = [] yielding [] ≈ 0.5 []. When the concentration decreases, [] becomes non-negligible so that the preceding expression is no longer valid. For vanishing acid concentrations, one can recover the final pH and the solubility of in pure water. In the case of a weak monoacid (here we take acetic acid with pKa = 4.76) with decreasing total acid concentration [A] = [] + [AH], we obtain: {| class="wikitable" |- ! width="160" |[A] (mol/L) | [] ≈ 0.5 [] | 10−1 | 10−2 | 10−3 | 10−4 | 10−5 | 10−6 | 10−7 | 10−10 |- ! width="160" |Initial pH | 2.38||2.88||3.39||3.91||4.47||5.15||6.02||6.79||7.00 |- ! width="160" |Final pH | 6.75||7.25||7.75||8.14||8.25||8.26||8.26||8.26||8.27 |- ! width="160" |Dissolved (g/L of acid) | 49.5||4.99||0.513||0.0848||0.0504||0.0474||0.0471||0.0470||0.0470 |} For the