Information about 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)

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
• 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.

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) or it can be prepared by passing carbon dioxide into a solution of calcium hydroxide: the calcium carbonate precipitates out, and this grade of product is referred to as a precipitate (abbreviated to PCC).

Ca(OH)₂ + CO₂ → CaCO₃ + H₂O

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 quick lime:
   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.

Uses

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 . A common contaminant is magnesium carbonate.

Calcium carbonate is widely used as an extender in paints, 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. 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. 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 widely used medicinally as an inexpensive dietary calcium supplement, antacid, and/or phosphate binder. It is also used in the pharmaceutical industry as a base material for tablets of other pharmaceuticals.

Calcium carbonate is known as whiting in ceramics/glazing applications, 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.

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.

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 micron in diameter, useful in coatings for paper.

As a food additive, it is used in some soy milk products as a source of dietary calcium.

In 1989, a researcher introduced CaCO₃ into the Whetstone Brook in Massachusetts . 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 aluminum 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. Nowadays, calcium carbonate is used to neutralise acidic conditions in both soil and water.

Calcination Equilibrium

Equilibrium Pressure of CO2 over CaCO3[1]
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 answer to the question, "how hot does the fire have to be?" 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 of calcium carbonate in water

Solubility in pure water with varying CO₂ pressure

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⁺][CO₃^2–] is given as anywhere from K_sp = 3.7×10^–9 to K_sp = 8.7×10^–9 at 25 °C, depending upon the data source.^[2][3] 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 CO₃^2– cannot exceed the value of K_sp. 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 CO₃^2– 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:

:

where kH = 29.76 atm/(mol/L) at 25°C (Henry constant), being the CO2 partial pressure.

Calcium Ion Solubility as a function of CO2 partial pressure at 25 °C
(atm)	pH	[Ca2+] (mol/L)
10−12	12.0	5.19 × 10−3
10−10	11.3	1.12 × 10−3
10−8	10.7	2.55 × 10−4
10−6	9.83	1.20 × 10−4
10−4	8.62	3.16 × 10−4
3.5 × 10−4	8.27	4.70 × 10−4
10−3	7.96	6.62 × 10−4
10−2	7.30	1.42 × 10−3
10−1	6.63	3.05 × 10−3
1	5.96	6.58 × 10−3
10	5.30	1.42 × 10−2

For ambient air, 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 , 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[CO₃^2–] + [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 (K_sp = 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 , 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.

Solubility at atmospheric CO₂ pressure with varying pH

We now consider the problem of the maximum solubility of calcium carbonate in normal atmospheric conditions ( = 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:

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

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.

References

1. ^ CRC Handbook of Chemistry and Physics 44th ed., p2292
2. ^ CSUDH
3. ^ CRC Handbook of Chemistry and Physics, 44th ed.

See also

• Aragonite
• Calcite
• Limestone
• Ocean acidification
• Gesso
• Cuttlefish
• Cuttlebone

External links

• International Chemical Safety Card 1193
• CID 516889 from PubChem
• ATC codes: A02AC01 and A12AA04

Limestone is a sedimentary rock composed largely of the mineral calcite (calcium carbonate: CaCO₃). Limestone often contains variable amounts of silica in the form of chert or flint, as well as varying amounts of clay, silt and sand as disseminations, nodules, or layers
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calcite]] The carbonate mineral, calcite, is a chemical or biochemical calcium carbonate corresponding to the formula CaCO₃ and is one of the most widely distributed minerals on the Earth's surface.
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Aragonite is a carbonate mineral, one of the two common, naturally occurring polymorphs of calcium carbonate, CaCO₃. The other is the mineral calcite. Aragonite's crystal lattice differs from that of calcite, resulting in a different crystal shape, an orthorhombic
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Chalk (IPA: /ˈtʃɔːk/) is a soft, white, porous sedimentary rock, a form of limestone composed of the mineral calcite.
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Marble is a nonfoliated metamorphic rock resulting from the metamorphism of limestone, composed mostly of calcite (a crystalline form of calcium carbonate, CaCO₃). It is extensively used for sculpture, as a building material, and in many other applications.
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CAS registry numbers are unique numerical identifiers for chemical compounds, polymers, biological sequences, mixtures and alloys. They are also referred to as CAS numbers, CAS RNs or CAS #s.
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A chemical formula is a concise way of expressing information about the atoms that constitute a particular chemical compound. A chemical formula is also a short way of showing how a chemical reaction occurs.
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Molar mass, symbol M,^[1] is the mass of one mole of a substance (chemical element or chemical compound).^[2] It is a physical property which is characteristic of each pure substance.
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In physics, density is mass m per unit volume V—how heavy something is compared to its size. A small, heavy object, such as a rock or a lump of lead, is denser than a lighter object of the same size or a larger object of the same weight, such as pieces of
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The melting point of a crystalline solid is the temperature range at which it changes state from solid to liquid. Although the phrase would suggest a specific temperature and is commonly and incorrectly used as such in most textbooks and literature, most crystalline compounds
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boiling point of a liquid is the temperature at which the vapor pressure of the liquid equals the environmental pressure surrounding the liquid.^[1][2][3][4]
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Solubility is a physical property referring to the ability for a given substance, the solute, to dissolve in a solvent.^[1] It is measured in terms of the maximum amount of solute dissolved in a solvent at equilibrium. The resulting solution is called a saturated solution.
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standard state of a material is its state at 1 bar (100 kilopascals exactly). This pressure was changed from 1 atm (101.325 kilopascals) by IUPAC in 1990.^[1] The standard state of a material can be defined at any given temperature, most commonly 25 degrees Celsius,
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A chemical formula is a concise way of expressing information about the atoms that constitute a particular chemical compound. A chemical formula is also a short way of showing how a chemical reaction occurs.
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Calcium (IPA: /ˈkalsiəm/) is the chemical element in the periodic table that has the symbol Ca and atomic number 20. It has an atomic mass of 40.078.
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4, 2
(mildly acidic oxide)
Electronegativity 2.55 (Pauling scale)
Ionization energies
(more) 1st: 1086.5 kJmol⁻¹
2nd: 2352.6 kJmol⁻¹
3rd: 4620.5 kJmol⁻¹

Atomic radius 70 pm
Atomic radius (calc.
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2, −1
(neutral oxide)
Electronegativity 3.44 (Pauling scale)
Ionization energies
(more) 1st: 1313.9 kJmol⁻¹
2nd: 3388.3 kJmol⁻¹
3rd: 5300.5 kJmol⁻¹

Atomic radius 60 pm
Atomic radius (calc.
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Balanced Rock stands in Garden of the Gods park in Colorado Springs, CO]] A rock is a naturally occurring aggregate of minerals and/or mineraloids. The Earth's lithosphere is made of rock. In general rocks are of three types, namely, igneous, sedimentary, and metamorphic.
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shell is a hard, rigid outer layer, which has evolved in a very wide variety of different animals, including mollusks, sea urchins, crustaceans, turtles and tortoises, armadillos, etc.
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snail is loosely applied to almost all members of the molluscan class Gastropoda which have coiled shells in the adult stage.

The class Gastropoda is the second largest class of invertebrates, second only to the insects.
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eggshell is a term for the outer covering of a hard-shelled egg, and some forms of eggs with soft outer coats.

The generalized eggshell structure, which varies widely among species, is a protein matrix lined with mineral crystals, usually of a calcium compound such as
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Agricultural lime is a soil additive made from pulverized limestone or chalk. The primary active component is calcium carbonate. Additional chemicals vary depending on the mineral source and may include calcium oxide, magnesium oxide and magnesium carbonate.
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Calcium (IPA: /ˈkalsiəm/) is the chemical element in the periodic table that has the symbol Ca and atomic number 20. It has an atomic mass of 40.078.
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