Calcium carbonate

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Calcium carbonate is a chemical compound, with the chemical formula CaCO3. It is a common substance found as rock in all parts of the world, and is the main component of shells of marine organisms, snails, and eggshells. Calcium carbonate is the active ingredient in agricultural lime, and is usually the principal cause of hard water. It is commonly used medicinally as a calcium supplement or as an antacid.

Occurrence
Calcium carbonate is found naturally as the following minerals and rocks: Aragonite, Calcite, Vaterite or (μ-CaCO3), Chalk, Limestone, Marble and Travertine

To test whether a mineral or rock contains calcium carbonate, strong acids, such as hydrochloric acid, can be added to it. If the sample does contain calcium carbonate, it will fizz and produce carbon dioxide and water. Weak acids such as acetic acid will react, albeit less vigorously. All of the rocks/minerals mentioned above will react with acid.

Chemical properties

 * See also: Carbonate

Calcium carbonate shares the typical properties of other carbonates. Notably:
 * 1) it reacts with strong acids, releasing carbon dioxide: CaCO3 + 2HCl → CaCl2 + CO2 + H2O

Calcium carbonate will react with water that is saturated with carbon dioxide to form the soluble calcium bicarbonate.
 * CaCO3 + CO2 + H2O → Ca(HCO3)2

Preparation
The vast majority of calcium carbonate used in industry is extracted by mining or quarrying. Pure calcium carbonate (e.g. for food or pharmaceutical use), can be produced from a pure quarried source (usually marble).

Alternatively, calcium oxide is prepared by calcining crude calcium carbonate. Water is added to give calcium hydroxide, and carbon dioxide is passed through this solution to precipitate the desired calcium carbonate, referred to in the industry as precipitated calcium carbonate (PCC):


 * CaCO3 &rarr; CaO + CO2
 * CaO + H2O &rarr; Ca(OH)2
 * Ca(OH)2 + CO2 → CaCO3 + H2O

Health and dietary applications
Calcium carbonate is widely used medicinally as an inexpensive dietary calcium supplement or antacid. It may be used as a phosphate binder for the treatment of hyperphosphatemia (primarily in patients with chronic renal failure) when lanthanum carbonate is not prescribed. It is also used in the pharmaceutical industry as an inert filler for tablets and other pharmaceuticals.

As a food additive, it is used in some soy milk products as a source of dietary calcium; one study concludes that calcium carbonate is as bioavailable as ordinary cow's milk.

With varying pH
We now consider the problem of the maximum solubility of calcium carbonate in normal atmospheric conditions ($$\scriptstyle P_{\mathrm{CO}_2}$$ = 3.5 × 10&minus;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 NaHSO4 to decrease the pH or of NaHCO3 to increase it). From the above equations for the solubility product, the hydratation reaction and the two acid reactions, the following expression for the maximum [Ca2+] can be easily deduced:
 * $$[\mathrm{Ca}^{2+}]_\mathrm{max} = \frac{K_\mathrm{sp}k_\mathrm{H}} {K_\mathrm{h}K_\mathrm{a1}K_\mathrm{a2}} \frac{[\mathrm{H}^+]^2}{P_{\mathrm{CO}_2}}$$

showing a quadratic dependence in [H+]. The numerical application with the above values of the constants gives

Comments:
 * decreasing the pH from 8 to 7 increases the maximum Ca2+ concentration by a factor 100
 * note that the Ca2+ concentration of the previous table is recovered for pH = 8.27
 * keeping the pH to 7.4 in a swimming pool (which gives optimum HClO/OCl− ratio in the case of "chlorine" maintenance) results in a maximum Ca2+ concentration of 1010 mg/L. This means that successive cycles of water evaporation and partial renewing may result in a very hard water before CaCO3 precipitates. Addition of a calcium sequestrant or complete renewing of the water will solve the problem.

Solubility in a strong or weak acid solution
Solutions of strong (HCl) or weak (acetic, phosphoric) acids are commercially available. They are commonly used to remove limescale deposits. The maximum amount of CaCO3 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 concentration [A] = [A&minus;], we obtain (with CaCO3 molar mass = 100 g):

where the initial state is the acid solution with no Ca2+ (not taking into account possible CO2 dissolution) and the final state is the solution with saturated Ca2+. For strong acid concentrations, all species have a negligible concentration in the final state with respect to Ca2+ and A&minus; so that the neutrality equation reduces approximately to 2[Ca2+] = [A&minus;] yielding $$\scriptstyle[\mathrm{Ca}^{2+}] \simeq \frac{[\mathrm{A}^-]}{2}$$. When the concentration decreases, [HCO3&minus;] becomes non negligible so that the preceding expression is no longer valid. For vanishing acid concentrations, we recover the final pH and the solubility of CaCO3 in pure water.


 * In the case of a weak monoacid (here we take acetic acid with pKA = 4.76) with decreasing concentration [A] = [A&minus;]+[AH], we obtain:

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 that the pKA, so that the weak acid is almost completely dissociated, yielding in the end as many H+ ions as the strong acid to "dissolve" the calcium carbonate.


 * The calculation in the case of phosphoric acid (which is the most widely used for domestic applications) is more complicated since the concentrations of the four dissociation states corresponding to this acid must be calculated together with [HCO3&minus;], [CO32&minus;], [Ca2+], [H+] and [OH&minus;]. The system may be reduced to a seventh degree equation for [H+] the numerical solution of which gives

where [A] = [H3PO4] + [H2PO4&minus;] + [HPO42&minus;] + [PO43&minus;]. We see that phosphoric acid is more efficient than a monoacid since at the final almost neutral pH, the second dissociated state concentration [HPO42&minus;] is not negligible (see phosphoric acid ).