Calcium carbonate 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 CO much higher than atmospheric. As such water percolates through calcium carbonate rock, the CaCO dissolves according to the second trend. When that same water then emerges from the tap, in time it comes into equilibrium with CO levels in the air by outgassing its excess CO. The calcium carbonate becomes less soluble as a result and the excess precipitates as lime scale. This same process is responsible for the formation of stalactites and stalagmites in limestone
Calcium carbonate caves. Two hydrated phases of calcium carbonate, monohydrocalcite, CaCO·HO and ikaite, CaCO·6HO, may precipitate from water at ambient conditions and persist as metastable phases. In contrast to the open equilibrium scenario above, many swimming pools are managed by addition of sodium bicarbonate (NaHCO) to about 2 mM as a buffer, then control of pH through use of HCl, NaHSO, NaCO, NaOH or chlorine formulations that are acidic or basic. In this situation, dissolved inorganic carbon (total inorganic carbon) is far from equilibrium with atmospheric CO. Progress towards equilibrium through outgassing of CO is slowed by (i) the slow reaction HCO