Calcium Carbonate Chemistry

The chemistry of calcium carbonate precipitation and dissolution governs where and how limestones form. Chemical weathering releases ions — including Ca²⁺, HCO₃^-, and CO₃^2- — from source rocks, and these ultimately accumulate in the ocean. Calcium ions have short residence times (~1,000,000 years) and carbonate ions have even shorter residence times (~110,000 years), meaning they are rapidly removed from solution to form carbonate rocks. ^[1] Bicarbonate ions (HCO₃^-) and carbonate ions (CO₃^2-) together make up almost 49 percent of total dissolved solids in average river water but less than 1 percent in ocean water, while Ca²⁺ makes up about 12 percent of dissolved solids in river water but only about 1 percent in ocean water — the discrepancy reflects how efficiently these ions are extracted from seawater to build carbonate rocks. ^[1] Chemically and biochemically deposited sedimentary rocks make up nearly one-fourth of all sedimentary rocks, and the majority of these are carbonate rocks. ^[1]

The CO₂ Equilibrium System

The dissolution and precipitation of calcite and aragonite are controlled chiefly by pH, which is in turn controlled by the partial pressure of dissolved CO₂. ^[1] When CO₂ dissolves in water the following chain of reactions occurs: CO₂ + H₂O → H₂CO₃ (carbonic acid); H₂CO₃ → H⁺ + HCO₃^- (bicarbonate ion); HCO₃^- → H⁺ + CO₃^2- (carbonate ion). ^[1] The second step generates H⁺ ions far faster than the third step generates carbonate ions, so adding CO₂ to water actually increases acidity and promotes dissolution of carbonate minerals — not precipitation. ^[1]

The key reversible reaction for carbonate rocks is: H₂O + CO₂ + CaCO₃ ⇌ Ca²⁺ + 2HCO₃^-. ^[1] When equilibrium is disturbed by loss of CO₂, H⁺ concentration decreases, pH rises, and the reaction shifts toward precipitation of solid CaCO₃. ^[1] This is the fundamental logic behind all inorganic carbonate precipitation: anything that removes CO₂ from the water — warming, pressure decrease, or photosynthesis — creates the conditions for precipitation.

Three Inorganic Triggers for Precipitation

Three water conditions promote inorganic CaCO₃ precipitation by decreasing carbonate solubility.

Temperature increase reduces the solubility of CO₂ (and other gases), causing CO₂ to escape, raising pH, and simultaneously reducing the solubility of calcite and aragonite directly — meaning warmer water is more likely to precipitate carbonate. ^[1] Calcium carbonate deposition is therefore favored in tropical areas where surface water temperatures may reach almost 30°C, compared to about 0°C in polar regions. ^[1]

Pressure decrease also allows CO₂ to escape: wave agitation by storms, wave breaking in the surf zone over shallow banks, or upwelling of deep pressurized waters can all lower pressure and release CO₂. ^[1]

Salinity decrease reduces ionic strength, lowering the concentration of foreign ions (including Mg²⁺) that interfere with carbonate crystal growth — so calcium carbonate solubility is several orders of magnitude lower in freshwater than in seawater and precipitates more readily in freshwater. ^[1] In the open ocean, however, salinity variation is modest (about 32–36 parts per thousand) and so has relatively little practical effect on precipitation. ^[1]

Why the Ocean Is Oversaturated but Does Not Precipitate Freely

Near-surface water in the modern ocean is oversaturated by more than six times with respect to calcite and more than four times with respect to aragonite — yet significant inorganic precipitation is inhibited by two factors. ^[1]

First, seawater is a well-buffered solution. The high alkalinity of ocean water — the high concentrations of bicarbonate and carbonate ions already present — inhibits breakdown of H₂CO₃, so the actual pH change in seawater from CO₂ gain or loss is comparatively small, and pH values in the open ocean rarely fall outside the range of 7.8 to 8.3. ^[1]

Second, Mg²⁺ ions at the concentration levels found in seawater strongly inhibit precipitation of calcite. Mg²⁺ is readily adsorbed onto calcite crystal surfaces and incorporated into the crystal structure; this nonequilibrium incorporation destabilizes the growing crystals and increases calcite solubility, preventing free nucleation and growth. ^[1] Aragonite has a different (orthorhombic) crystal structure in which Mg²⁺ ions are less prone to sorb to growing nuclei, so aragonite is less inhibited by Mg²⁺ and tends to precipitate preferentially in the presence of Mg²⁺. ^[1] Even aragonite, however, does not precipitate completely freely in ocean water, possibly because thin organophosphatic coatings form on aragonite seed nuclei and inhibit their growth. ^[1]

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