Calcium carbonate Totally Explained

Everything about Calcium Carbonate totally explained

| Section2 = | Section3 = | Section7 =,, | SPhrases =, }} }} Calcium carbonate is a chemical compound with the chemical formula Ca C O₃. 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, but high consumption can be hazardous.

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 or sulfuric acid, can be added to it. If the sample does contain calcium carbonate, it'll 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

ť See also: Carbonate

Calcium carbonate shares the typical properties of other carbonates. Notably:

it reacts with strong acids, releasing carbon dioxide:
CaCO_3(s) + 2HCl_(aq) → CaCl_2(aq) + CO_2(g) + H₂O_(l)

it releases carbon dioxide on heating (to above 840 °C in the case of CaCO₃), to form calcium oxide, commonly called quicklime, with reaction enthalpy 178 kJ / mole:
CaCO₃ → CaO + CO₂ Calcium carbonate will react with water that's saturated with carbon dioxide to form the soluble calcium bicarbonate. ť 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 (for example 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):
ť CaCO₃ → CaO + CO₂

CaO + H₂O → Ca(OH)₂ ť 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 (for example 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.
   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. It may be used as a phosphate binder for the treatment of hyperphosphatemia (primarily in patients with chronic renal failure). It is also used in the pharmaceutical industry as an inert filler for tablets and other pharmaceuticals. Calcium carbonate is also used in Homeopathy. It is one of the constitutional remedies. Excess calcium from supplements, fortified food and high-calcium diets, can cause the "milk alkali syndrome," which has serious toxicity and can be fatal. In 1915, Bertram Sippy introduced the "Sippy regimen" of hourly ingestion of milk and cream, and the gradual addition of eggs and cooked cereal, for 10 days, combined with alkaline powders, which provided symptomatic relief for peptic ulcer disease. Over the next several decades, the Sippy regimen resulted in renal failure, alkalosis, and hypercalemia, mostly in men with peptic ulcer disease. These adverse effects were reversed when the regimen stopped, but it was fatal in some patients with protracted vomiting. Milk alkali syndrome declined in men after effective treatments for peptic ulcer disease. But during the past 15 years, it has been reported in women taking calcium supplements above the recommended range of 1200 to 1500 mg daily, for prevention and treatment of osteoporosis, and is exacerbated by dehydration. Calcium has been added to over-the-counter products, which contributes to inadvertent excessive intake. Excessive calcium intake can lead to hypercalcemia, complications of which include vomiting, abdominal pain and altered mental status.
   A form of food additive is designated as E170. It is used in some soy milk products as a source of dietary calcium; one study suggests that calcium carbonate might be bioavailable as the calcium in cow's milk.

Ecological applications

In 1989, a researcher, Ken 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 aluminium ions in the area of the brook that wasn't 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.

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 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's a partial pressure of carbon dioxide that's 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's 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 isn't 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
$scriptstyle P_;">$ [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)	62.0	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⁻] isn't negligible (see phosphoric acid ).

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