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Calcium carbonate

Calcium carbonate
Other names Limestone; calcite; aragonite; chalk; marble
CAS number 471-34-1
Molecular formula CaCO3
Molar mass 100.087 g/mol
Appearance White powder.
Density 2.83 g/cm³, solid.
Melting point

825 °C

Boiling point


Solubility in water Insoluble
Molecular shape Linear
Main hazards Not hazardous.
R-phrases R36, R37, R38
S-phrases S26, S36
Flash point Non-flammable.
Except where noted otherwise, data are given for
materials in their standard state
(at 25 °C, 100 kPa)

Infobox disclaimer and references

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.

Additional recommended knowledge



Calcium carbonate is found naturally as the following minerals and rocks:

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
  2. it releases carbon dioxide on heating (to above 840 °C in the case of CaCO3), to form calcium oxide, commonly called quicklime:
    CaCO3 → CaO + CO2


CaCO3 + CO2 + H2O → Ca(HCO3)2




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):[1]

CaCO3 → CaO + CO2
CaO + H2O → Ca(OH)2
Ca(OH)2 + CO2 → CaCO3 + H2O


Industrial applications

The main use of calcium carbonate is in the construction industry, either as a building material in its own right (e.g. marble) or limestone aggregate for roadbuilding or as an ingredient of cement or as the starting material for the preparation of builder's lime by burning in a kiln.

Calcium carbonate is also used in the purification of iron from iron ore in a blast furnace. Calcium carbonate is calcined in situ to give calcium oxide, which forms a slag with various impurities present, and separates from the purified iron.[2]

Calcium carbonate is widely used as an extender in paints,[3] in particular matte emulsion paint where typically 30% by weight of the paint is either chalk or marble.

Calcium carbonate is also widely used as a filler in plastics.[3] Some typical examples include around 15 to 20% loading of chalk in uPVC drain pipe, 5 to 15% loading of stearate coated chalk or marble in uPVC window profile. Fine ground calcium carbonate is an essential ingredient in the microporous film used in babies' diapers and some building films as the pores are nucleated around the calcium carbonate particles during the manufacture of the film by biaxial stretching.

Calcium carbonate is also used in a wide range of trade and DIY adhesives, sealants, and decorating fillers.[3] Ceramic tile adhesives typically contain 70 to 80% limestone. Decorating crack fillers contain similar levels of marble or dolomite. It is also mixed with putty in setting stained glass windows, and as a resist to prevent glass from sticking to kiln shelves when firing glazes and paints at high temperature.

Calcium carbonate is known as whiting in ceramics/glazing applications,[3] where it is used as a common ingredient for many glazes in its white powdered form. When a glaze containing this material is fired in a kiln, the whiting acts as a flux material in the glaze.

In North America, calcium carbonate has begun to replace kaolin in the production of glossy paper. Europe has been practicing this as alkaline papermaking or acid-free papermaking for some decades. Carbonates are available in forms: ground calcium carbonate (GCC) or precipitated calcium carbonate (PCC). The latter has a very fine and controlled particle size, on the order of 2 micrometres in diameter, useful in coatings for paper.

Used in swimming pools as a pH corrector for maintaining alkalinity "buffer" to offset the acidic properties of the disinfectant agent.

It is commonly called chalk as it has been a major component of blackboard chalk. Chalk may consist of either calcium carbonate or gypsum, hydrated calcium sulfate CaSO4·2H2O.

Health and dietary applications

Calcium carbonate is widely used medicinally as an inexpensive dietary calcium supplement or antacid.[4] 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.[5]

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.[6]

Ecological applications

In 1989, a researcher, Ken Simmons, introduced CaCO3 into the Whetstone Brook in Massachusetts.[7] His hope was that the calcium carbonate would counter the acid in the stream from acid rain and save the trout that had ceased to spawn. Although his experiment was a success, it did increase the amounts of aluminium ions in the area of the brook that was not treated with the limestone. This shows that CaCO3 can be added to neutralize the effects of acid rain in river ecosystems. Currently calcium carbonate is used to neutralize acidic conditions in both soil and water.[8][9]

Calcination equilibrium

Equilibrium Pressure of CO2 over CaCO3[10]
550 °C 0.055 kPa
587 °C 0.13 kPa
605 °C 0.31 kPa
680 °C 1.80 kPa
727 °C 5.9 kPa
748 °C 9.3 kPa
777 °C 14 kPa
800 °C 24 kPa
830 °C 34 kPa
852 °C 51 kPa
871 °C 72 kPa
881 °C 80 kPa
891 °C 91 kPa
898 °C 101 kPa
937 °C 179 kPa
1082 °C 901 kPa
1241 °C 3961 kPa





With varying CO2 pressure

Calcium ion solubility
as a function of CO2 partial pressure at 25 °C
\scriptstyle P_{\mathrm{CO}_2} (atm) pH [Ca2+] (mol/L)
10−12 12.05.19 × 10−3
10−10 11.31.12 × 10−3
10−8 10.72.55 × 10−4
10−6 9.831.20 × 10−4
10−4 8.623.16 × 10−4
3.5 × 10−4 8.274.70 × 10−4
10−3 7.966.62 × 10−4
10−2 7.301.42 × 10−3
10−1 6.633.05 × 10−3
1 5.966.58 × 10−3
10 5.301.42 × 10−2

Calcium carbonate is poorly soluble in pure water. The equilibrium of its solution is given by the equation (with dissolved calcium carbonate on the right):

CaCO3 Ca2+ + CO32– Ksp = 3.7×10–9 to 8.7×10–9 at 25 °C

where the solubility product for [Ca2+][CO32–] is given as anywhere from Ksp = 3.7×10–9 to Ksp = 8.7×10–9 at 25 °C, depending upon the data source.[11][10] What the equation means is that the product of molar concentration of calcium ions (moles of dissolved Ca2+ per liter of solution) with the molar concentration of dissolved CO32– cannot exceed the value of Ksp. This seemingly simple solubility equation, however, must be taken along with the more complicated equilibrium of carbon dioxide with water (see carbonic acid). Some of the CO32– combines with H+ in the solution according to:

HCO3 H+ + CO32–    Ka2 = 5.61×10–11 at 25 °C

HCO3 is known as the bicarbonate ion. Calcium bicarbonate is many times more soluble in water than calcium carbonate -- indeed it exists only in solution.


H2CO3 H+ + HCO3    Ka1 = 2.5×10–4 at 25 °C


H2O + CO2(dissolved) H2CO3    Kh = 1.70×10–3 at 25 °C


\frac{P_{\mathrm{CO}_2}}{[\mathrm{CO}_2]}\ =\ k_\mathrm{H} where kH = 29.76 atm/(mol/L) at 25 °C (Henry constant), \scriptstyle P_{\mathrm{CO}_2} being the CO2 partial pressure.

For ambient air, \scriptstyle P_{\mathrm{CO}_2} is around 3.5×10–4 atmospheres (or equivalently 35 Pa). The last equation above fixes the concentration of dissolved CO2 as a function of \scriptstyle P_{\mathrm{CO}_2}, independent of the concentration of dissolved CaCO3. At atmospheric partial pressure of CO2, dissolved CO2 concentration is 1.2×10–5 moles/liter. The equation before that fixes the concentration of H2CO3 as a function of [CO2]. For [CO2]=1.2×10–5, it results in [H2CO3]=2.0×10–8 moles per liter. When [H2CO3] is known, the remaining three equations together with

H2O H+ + OH K = 10–14 at 25 °C




The table on the right shows the result for [Ca2+] and [H+] (in the form of pH) as a function of ambient partial pressure of CO2 (Ksp = 4.47×10−9 has been taken for the calculation). At atmospheric levels of ambient CO2 the table indicates the solution will be slightly alkaline. The trends the table shows are

1) As ambient CO2 partial pressure is reduced below atmospheric levels, the solution becomes more and more alkaline. At extremely low \scriptstyle P_{\mathrm{CO}_2}, dissolved CO2, bicarbonate ion, and carbonate ion largely evaporate from the solution, leaving a highly alkaline solution of calcium hydroxide, which is more soluble than CaCO3.


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

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 × 10−4 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

pH 7.0 7.2 7.4 7.6 7.8 8.0 8.2 8.27 8.4
[Ca2+]max (10-4mol/L or °F) 1590 635 253 101 40.0 15.9 6.35 4.70 2.53
[Ca2+]max (mg/L) 6390 2540 1010 403 160 63.9 25.4 18.9 10.1


  • 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], we obtain (with CaCO3 molar mass = 100 g):
Initial pH
Final pH 6.757.257.758.
Dissolved CaCO3 (g per liter of acid)

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 so that the neutrality equation reduces approximately to 2[Ca2+] = [A] yielding \scriptstyle[\mathrm{Ca}^{2+}] \simeq \frac{[\mathrm{A}^-]}{2}. When the concentration decreases, [HCO3] 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]+[AH], we obtain:
Initial pH 2.382.883.393.914.475.156.026.797.00
Final pH 6.757.257.758.
Dissolved CaCO3 (g per liter of acid) 49.54.990.5130.08480.05040.04740.04710.04700.0470

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.

Initial pH 1.081.622.
Final pH 6.717.177.638.
Dissolved CaCO3 (g per liter of acid) 62.07.390.8740.1230.05360.04770.04710.04710.0470

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

See also



  1. ^ Solvay Precipitated Calcium Carbonate: Production. Solvay S. A. (2007-03-09). Retrieved on 2007-12-30.
  2. ^ Blast Furnace. Science Aid. Retrieved on 2007-12-30.
  3. ^ a b c d Calcium Carbonate Powder. Reade Advanced Materials (2006-02-04). Retrieved on 2007-12-30.
  4. ^ Calcium Carbonate. Medline Plus. National Institutes of Health (2005-10-01). Retrieved on 2007-12-30.
  5. ^ Herbert A. Lieberman, Leon Lachman, Joseph B. Schwartz (1990). Pharmaceutical Dosage Forms: Tablets, 153. ISBN 0824780442. 
  7. ^ Associated Press. "Limestone Dispenser Fights Acid Rain in Stream", New York Times, 1989-06-13. 
  8. ^ R. K. Schreiber (1988). "Cooperative federal-state liming research on surface waters impacted by acidic deposition". Water, Air, & Soil Pollution 41 (1): 53-73. doi:10.1007/BF00160344.
  9. ^ Dan Kircheis; Richard Dill (2006). Effects of low pH and high aluminum on Atlantic salmon smolts in Eastern Maine and liming project feasibility analysis (reprinted at Downeast Salmon Federation). National Marine Fisheries Service and Maine Atlantic Salmon Commission.
  10. ^ a b CRC Handbook of Chemistry and Physics 44th ed.
  11. ^ Selected Solubility Products and Formation Constants at 25 °C. California State University, Dominguez Hills.
This article is licensed under the GNU Free Documentation License. It uses material from the Wikipedia article "Calcium_carbonate". A list of authors is available in Wikipedia.