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Calcium carbonate

# Drugs & Medication

## Calcium carbonate

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Calcium carbonate
General
Systematic name Calcium carbonate
Other names Limestone,
calcite,
aragonite,
chalk,
marble
Molecular formula CaCO3
Molar mass 100.087 g/mol
Appearance White powder.
CAS number [471-34-1]
Properties
Density and phase 2.83 g/cm3, solid.
Solubility in water Insoluble
Melting point 825°C (1098 K)
Boiling point Decomposes
Thermochemistry
ΔfH0liquid

-1154 kJ/mol

ΔfH0solid

-1207 kJ/mol

S0solid

93 J/mol·K

Structure
Molecular shape Linear
Coordination
geometry
Tetrahedral
Dipole moment  ? D
Hazards
MSDS External MSDS
Main hazards Not hazardous.
NFPA 704
Flash point Non-flammable.
R/S statement R: R36, R37, R38
S: S26, S36
Supplementary data page
Structure and
properties
n, εr, etc.
Thermodynamic
data
Phase behaviour
Solid, liquid, gas
Related compounds
Other anions Calcium bicarbonate
Calcium sulfate
Other cations Magnesium carbonate (dolomite)
Strontium carbonate
Related compounds Calcium oxide
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 chemical formula CaCO3. It is commonly used medicinally as a calcium supplement or as an antacid. Calcium carbonate is the active ingredient in agricultural lime. It is a common substance found as rock in all parts of the world and is the main component of seashells and the shell of snails. It is usually the principal cause of hard water.

## Occurrence

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

• Aragonite
Calcite
Chalk
Limestone
Marble
Travertine

Eggshells are composed of approximately 95% calcium carbonate.

To test whether a mineral or rock contains calcium carbonate, strong acids, like hydrochloric acid, can be dropped with a dropper onto it. If it does contain the chemical, it will fizz and produce carbon dioxide; otherwise, it probably wouldn't react vigorously. For example, all of the rocks/mineral 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

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 825 °C in the case of CaCO3), to form calcium oxide, commonly called burnt 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 contaminate 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 nappies 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,[1] 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.

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, Dr. Simmons 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
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

Calcium carbonate is poorly soluble in water.

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 12.0 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

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. 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 extremly 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 water 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.

## References

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

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