"කැල්සියම් කාබනේට්" හි සංශෝධන අතර වෙනස්කම්

Calciumකැල්සියම් carbonateකාබනේට් is poorly soluble in pure water (47 mg/L at normal atmospheric CO₂ partial pressure as shown below).

The equilibrium of its solution is given by the equation (with dissolved calciumකැල්සියම් carbonateකාබනේට් on the right):

HCO₃^– is known as the [[bicarbonate]] ion. [[Calcium bicarbonate]] is many times more soluble in water than calciumකැල්සියම් carbonate—indeedකාබනේට්—indeed it exists ''only'' in solution.

make it possible to solve simultaneously for the remaining five unknown concentrations (note that the above form of the neutrality equation is valid only if calciumකැල්සියම් carbonateකාබනේට් has been put in contact with pure water or with a neutral pH solution; in the case where the origin water solvent pH is not neutral, the equation is modified).

[[Image:CanarySpring.jpg|thumb|right|Travertine calciumකැල්සියම් carbonateකාබනේට් deposits from a [[hot spring]]]]

The effect of the latter is especially evident in day to day life of people who have hard water. Water in aquifers underground can be exposed to levels of CO₂ much higher than atmospheric. As such water percolates through calciumකැල්සියම් carbonateකාබනේට් rock, the CaCO₃ dissolves according to the second trend. When that same water then emerges from the tap, in time it comes into equilibrium with CO₂ levels in the air by outgassing its excess CO₂. The calciumකැල්සියම් carbonateකාබනේට් becomes less soluble as a result and the excess precipitates as lime scale. This same process is responsible for the formation of [[stalactites]] and [[stalagmite]]s in limestone caves.

Two hydrated phases of calciumකැල්සියම් carbonateකාබනේට්, [[monohydrocalcite]], CaCO₃·H₂O and [[ikaite]], CaCO₃·6H₂O, may [[precipitate]] from water at ambient conditions and persist as metastable phases.

We now consider the problem of the maximum solubility of calciumකැල්සියම් carbonateකාබනේට් in normal atmospheric conditions (<math>\scriptstyle P_{\mathrm{CO}_2}</math> = 3.5 × 10⁻⁴ atm) when the pH of the solution is adjusted. This is for example the case in a swimming pool where the pH is maintained between 7 and 8 (by addition of [[sodium bisulfate]] NaHSO₄ to decrease the pH or of [[sodium bicarbonate]] NaHCO₃ to increase it). From the above equations for the solubility product, the hydration reaction and the two acid reactions, the following expression for the maximum [Ca²⁺] can be easily deduced:

We see that for the same total acid concentration, the initial pH of the weak acid is less acid than the one of the strong acid; however, the maximum amount of CaCO₃ which can be dissolved is approximately the same. This is because in the final state, the pH is larger than the p''K''_A, so that the weak acid is almost completely dissociated, yielding in the end as many H⁺ ions as the strong acid to "dissolve" the calciumකැල්සියම් carbonateකාබනේට්.