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 as 825 °C, but stating an absolute threshold is misleading. Calcium carbonate exists in equilibrium with calcium oxide and carbon dioxide at any temperature. At each temperature there is a partial pressure of carbon dioxide that is in equilibrium with calcium carbonate. At room temperature the equilibrium overwhelmingly favors calcium carbonate, because the equilibrium CO pressure is only a tiny fraction of the partial CO pressure in air, which is about 0.035 kPa. At temperatures above 550 °C the equilibrium CO pressure begins to exceed the CO pressure in air. So above 550 °C, calcium carbonate begins to outgas CO