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# Calcium carbonate

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

825 °C

Boiling point

Decomposes

Solubility in water Insoluble
Structure
Molecular shape Linear
Hazards
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.

## Occurrence

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

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

Calcium carbonate will react with water that is saturated with carbon dioxide to form the soluble calcium bicarbonate.

CaCO3 + CO2 + H2O → Ca(HCO3)2

This reaction is important in the erosion of carbonate rocks, forming caverns, and leads to hard water in many regions.

## 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):[1]

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

## Uses

### 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

Calcination of limestone using charcoal fires to produce quicklime has been practiced since antiquity by cultures all over the world. The temperature at which limestone yields calcium oxide is usually given 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 CO2 pressure is only a tiny fraction of the partial CO2 pressure in air, which is about 0.035 kPa.

At temperatures above 550 °C the equilibrium CO2 pressure begins to exceed the CO2 pressure in air. So above 550 °C, calcium carbonate begins to outgas CO2 into air. But in a charcoal fired kiln, the concentration of CO2 will be much higher than it is in air. Indeed if all the oxygen in the kiln is consumed in the fire, then the partial pressure of CO2 in the kiln can be as high as 20 kPa.

The table shows that this equilibrium pressure is not achieved until the temperature is nearly 800 °C. For the outgassing of CO2 from calcium carbonate to happen at an economically useful rate, the equilibrium pressure must significantly exceed the ambient pressure of CO2. And for it to happen rapidly, the equilibrium pressure must exceed total atmospheric pressure of 101 kPa, which happens at 898 °C.

## Solubility

### 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.

Some of the HCO3 combines with H+ in solution according to:

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

Some of the H2CO3 breaks up into water and dissolved carbon dioxide according to:

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

And dissolved carbon dioxide is in equilibrium with atmospheric carbon dioxide according to:

 $\frac{P_{\mathrm{CO}_2}}{[\mathrm{CO}_2]}\ =\ k_\mathrm{H}$ 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

(which is true for all aqueous solutions), and the fact that the solution must be electrically neutral,

2[Ca2+] + [H+] = [HCO3] + 2[CO32–] + [OH]

make it possible to solve simultaneously for the remaining five unknown concentrations (note that the above form of the neutrality equation is valid only if calcium carbonate has been put in contact with pure water or with a neutral pH solution; in the case where the origin water solvent pH is not neutral, the equation is modified).

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.
2) As ambient CO2 partial pressure increases to levels above atmospheric, pH drops, and much of the carbonate ion is converted to bicarbonate ion, which results in higher solubility of Ca2+.

The effect of the 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 CO2 much higher than atmospheric. As such water percolates through calcium carbonate rock, the CaCO3 dissolves according to the second trend. When that same water then emerges from the tap, in time it comes into equilibrium with CO2 levels in the air by outgassing its excess CO2. 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 caves.

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 7.2 7.4 7.6 7.8 8 8.2 8.27 8.4 [Ca2+]max (10-4mol/L or °F) 1590 635 253 101 40 15.9 6.35 4.7 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):
 [A] (mol/L) 1 10−1 10−2 10−3 10−4 10−5 10−6 10−7 10−10 Initial pH 0 1.00 2.00 3.00 4.00 5.00 6.00 6.79 7.00 Final pH 6.75 7.25 7.75 8.14 8.25 8.26 8.26 8.26 8.27 Dissolved CaCO3 (g per liter of acid) 50 5.00 0.514 0.0849 0.0504 0.0474 0.0471 0.0470 0.0470

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:
 [A] (mol/L) 1 10−1 10−2 10−3 10−4 10−5 10−6 10−7 10−10 Initial pH 2.38 2.88 3.39 3.91 4.47 5.15 6.02 6.79 7.00 Final pH 6.75 7.25 7.75 8.14 8.25 8.26 8.26 8.26 8.27 Dissolved CaCO3 (g per liter of acid) 49.5 4.99 0.513 0.0848 0.0504 0.0474 0.0471 0.0470 0.0470

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.

• 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], [CO32−], [Ca2+], [H+] and [OH]. The system may be reduced to a seventh degree equation for [H+] the numerical solution of which gives
 [A] (mol/L) 1 10−1 10−2 10−3 10−4 10−5 10−6 10−7 10−10 Initial pH 1.08 1.62 2.25 3.05 4.01 5.00 5.97 6.74 7.00 Final pH 6.71 7.17 7.63 8.06 8.24 8.26 8.26 8.26 8.27 Dissolved CaCO3 (g per liter of acid) 62 7.39 0.874 0.123 0.0536 0.0477 0.0471 0.0471 0.0470

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 ).

## References

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.
6. ^ Y. Zhao, B. R. Martin and C. M. Weaver (2005). "Calcium Bioavailability of Calcium Carbonate Fortified Soymilk Is Equivalent to Cow's Milk in Young Women". J. Nutr. 135 (10): 2379-2382.
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.