Calcium carbonate

Calcium carbonate
IUPAC name Calcium carbonate
Other names Limestone; calcite; aragonite; chalk; marble; pearl
Identifiers
CAS number	471-34-1 Y
ChemSpider	9708 Y
UNII	H0G9379FGK Y
KEGG	D00932 Y
ChEBI	CHEBI:3311 Y
RTECS number	FF9335000
Jmol-3D images	Image 1 Image 2
SMILES [Ca+2].[O-]C([O-])=O C(=O)([O-])[O-].[Ca+2]
InChI InChI=1S/CH2O3.Ca/c2-1(3)4;/h(H2,2,3,4);/q;+2/p-2 Y Key: VTYYLEPIZMXCLO-UHFFFAOYSA-L Y InChI=1/CH2O3.Ca/c2-1(3)4;/h(H2,2,3,4);/q;+2/p-2 Key: VTYYLEPIZMXCLO-NUQVWONBAS
Properties
Molecular formula	CaCO3
Exact mass	100.0869 g/mol
Appearance	Fine white powder
Density	2.71 g/cm3 (calcite) 2.83 g/cm3 (aragonite)
Melting point	825 °C (aragonite) 1339 °C (calcite)[2]
Boiling point	decomposes
Solubility in water	0.00015 mol/L (25°C)
Solubility product, Ksp	4.8×10−9[1]
Solubility in dilute acids	soluble
Acidity (pKa)	9.0
Refractive index (nD)	1.59
Structure
Crystal structure	Trigonal
Space group	32/m
Hazards
MSDS	ICSC 1193
EU Index	Not listed
NFPA 704	0 0 0
Flash point	Non-flammable
Related compounds
Other anions	Calcium bicarbonate
Other cations	Magnesium carbonate Strontium carbonate Barium carbonate
Related compounds	Calcium sulfate
Y carbonate (verify) (what is: Y/N?) Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Crystal structure of calcite

Calcium carbonate is a chemical compound with the formula CaCO₃. It is a common substance found in rocks in all parts of the world, and is the main component of shells of marine organisms, snails, coal balls, pearls, 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 excessive consumption can be hazardous.

Chemical properties

See also: Carbonate

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

• it reacts with strong acids, releasing carbon dioxide:

CaCO_3(s) + 2 HCl_(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 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 carbonate is prepared by calcining crude calcium oxide. 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):^[3]

CaCO₃ → CaO + CO₂

CaO + H₂O → Ca(OH)₂

Ca(OH)₂ + CO₂ → CaCO₃ + H₂O

Occurrence

Calcium carbonate is found naturally as the following minerals in the form of polymorphs:

• Aragonite
• Calcite
• Vaterite or (μ-CaCO₃)

The trigonal crystal structure of calcite is most common.

The calcium carbonate minerals occur in the following rocks:

• Chalk
• Limestone
• Marble
• Travertine

Calcite   Aragonite   Marble   Travertine

Geology

Carbonate is found frequently in geologic settings and constitute an enormous carbon reservoir. Calcium carbonate occurs as the polymorphs aragonite and calcite. A polymorph is a mineral with the same chemical formula but different chemical structure. The carbonate minerals form the rock types: limestone, chalk, marble, travertine, tufa, and others. Calcite commonly occurs as sediments in marine settings. Calcite is typically found around the warm tropic environments. Calcite precipitates in warmer shallow environments more than it does under colder environments because warmer environments do not favor the dissolution of CO₂. This is analogous to CO₂ being dissolved in soda. When you take the cap off of a soda bottle, the CO₂ rushes out. As the soda warms up, carbon dioxide is released. This same principle can be applied to calcite in the ocean. Cold-water carbonates do exist at higher latitudes but have a very slow growth rate.

In tropic settings, the waters are warm and clear. Corals are more abundant in this environment than towards the poles where the waters are cold. Calcium carbonate contributors, including plankton (such as coccoliths and planktic foraminifera), coralline algae, sponges, brachiopods, echinoderms, bryozoa and mollusks, are typically found in shallow water environments where sunlight and filterable food are more abundant. The calcification processes are changed by the ocean acidification.

Where the oceanic crust is subducted under a continental plate sediments will be carried down to warmer zones in the astenosphere and mesosphere where the calcium carbonate is decomposed to carbon dioxide which will give rise to explosive vulcanic eruptions.

Carbonate compensation depth

The carbonate compensation depth (CCD) is the point in the ocean where the rate of precipitation of calcium carbonate is balanced by the rate of dissolution due to the conditions present. Deep in the ocean, the temperature drops and pressure increases. Calcium carbonate is unusual in that its solubility increases with decreasing temperature. Increasing pressure also increases the solubility of calcium carbonate. The CCD can range from 4–6 km below sea level.

Taphonomy

Calcium carbonate can preserve fossils through permineralization. Most of the vertebrate fossils of the Two Medicine Formation, known for its duck-billed dinosaur eggs, are preserved by CaCO₃ permineralization.^[4] This type of preservation preserves high levels of detail, even down to the microscopic level.^[4] However, it also leaves specimens vulnerable to weathering when exposed to the surface.^[4]

Uses

Industrial applications

The main use of calcium carbonate is in the construction industry, either as a building material 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. However, due to weathering mainly caused by acid rain, calcium carbonate (in limestone form) is no longer used for building purposes on its own, and only as a raw/primary substance for building materials.

Calcium carbonate is also used in the purification of iron from iron ore in a blast furnace. The carbonate is calcined in situ to give calcium oxide, which forms a slag with various impurities present, and separates from the purified iron.^[5]

In the oil industry, calcium carbonate is added to drilling fluids as a formation-bridging and filtercake-sealing agent; it is also a weighting material which increases the density of drilling fluids to control the downhole pressure. Calcium carbonate is added to swimming pools, as a pH corrector for maintaining alkalinity and offsetting the acidic properties of the disinfectant agent.

Calcium carbonate has traditionally been a major component of blackboard chalk. However, modern manufactured chalk is mostly gypsum, hydrated calcium sulfate CaSO₄·2H₂O. Calcium carbonate is a main source for growing Seacrete, or Biorock. Precipitated calcium carbonate (PCC), pre-dispersed in slurry form, is a common filler material for latex gloves with the aim of achieving maximum saving in material and production costs.^[6]

Fine ground calcium carbonate (GCC) 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. GCC or PCC is used as a filler in paper because they are cheaper than wood fiber. Printing and writing paper can contain 10–20% calcium carbonate. 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. PCC has a very fine and controlled particle size, on the order of 2 micrometres in diameter, useful in coatings for paper.

Calcium carbonate is widely used as an extender in paints,^[7] in particular matte emulsion paint where typically 30% by weight of the paint is either chalk or marble. It is also a popular filler in plastics.^[7] Some typical examples include around 15 to 20% loading of chalk in unplasticized polyvinyl chloride (uPVC) drain pipe, 5 to 15% loading of stearate coated chalk or marble in uPVC window profile. PVC cables can use calcium carbonate at loadings of up to 70 phr (parts per hundred parts of resin) to improve mechanical properties (tensile strength and elongation) and electrical properties (volume resistivity). Polypropylene compounds are often filled with calcium carbonate to increase rigidity, a requirement that becomes important at high use temperatures.^[8] It also routinely used as a filler in thermosetting resins (sheet and bulk molding compounds)^[8] and has also been mixed with ABS, and other ingredients, to form some types of compression molded "clay" poker chips. Precipitated calcium carbonate, made by dropping calcium oxide into water, is used by itself or with additives as a white paint, known as whitewashing.

Calcium carbonate is added to a wide range of trade and do it yourself adhesives, sealants, and decorating fillers.^[7] 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.

In ceramics/glazing applications, calcium carbonate is known as whiting,^[7] and is 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. Ground calcium carbonate is an abrasive (both as scouring powder and as an ingredient of household scouring creams), in particular in its calcite form, which has the relatively low hardness level of 3 on the Mohs scale of mineral hardness, and will therefore not scratch glass and most other ceramics, enamel, bronze, iron, and steel, and have a moderate effect on softer metals like aluminium and copper. A paste made from calcium carbonate and deionized water can be used to clean tarnish on silver.^[9]

Health and dietary applications

500-milligram calcium supplements made from calcium carbonate

Calcium carbonate is widely used medicinally as an inexpensive dietary calcium supplement or gastric antacid.^[10] 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.^[11]

Calcium carbonate is known among IBS sufferers to help reduce diarrhea. Some individuals report being symptom-free since starting supplementation. The process in which calcium carbonate reduces diarrhea is by binding water in the bowel, which creates a stool that is firmer and better formed. Calcium carbonate supplements are often combined with magnesium in various proportions. This should be taken into account as magnesium is known to cause diarrhea.

Calcium carbonate is used in the production of toothpaste and has seen a resurgence as a food preservative and color retainer, when used in or with products such as organic apples or food.^[12]

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 hypercalcemia, 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 arose. During the past 15 years, it has been reported in women taking calcium supplements above the recommended range of 1.2 to 1.5 g 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.^[13]

As a food additive it is designated E170^[14]; INS number 170. Used as an acidity regulator, anticaking agent, stabiliser or colour it is approved for usage in the EU,^[15] USA^[16] and Australia and New Zealand.^[17] It is used in some soy milk products as a source of dietary calcium; one study suggests that calcium carbonate might be as bioavailable as the calcium in cow's milk.^[18] Calcium carbonate is also used as a firming agent in many canned or bottled vegetable products.

Environmental applications

In 1989, a researcher, Ken Simmons, introduced CaCO₃ into the Whetstone Brook in Massachusetts.^[19] 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.^[20][21] Since the 1970s, such liming has been practiced on a large scale in Sweden to mitigate acidification and several thousand lakes and streams are limed repeatedly.^[22]

Calcination equilibrium

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. However, 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.

Equilibrium pressure of CO2 over CaCO3 (P) vs. temperature (T).[23]
P (kPa)	0.055	0.13	0.31	1.80	5.9	9.3	14	24	34	51	72	80	91	101	179	901	3961
T (°C)	550	587	605	680	727	748	777	800	830	852	871	881	891	898	937	1082	1241

Solubility

With varying CO₂ pressure

Calcium ion solubility as a function of CO2 partial pressure at 25 °C (Ksp = 4.47×10−9)
$\scriptstyle P_{\text{CO}_2}$ (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

Calcium carbonate is poorly soluble in pure water (47 mg/L at normal atmospheric CO₂ partial pressure as shown below).

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.^[23][24] 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:

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

where kH = 29.76 atm/(mol/L) at 25 °C (Henry constant), $\scriptstyle P_{\text{CO}_2}$ being the CO2 partial pressure.

For ambient air, $\scriptstyle P_{\text{CO}_2}$ 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 $\scriptstyle P_{\text{CO}_2}$, 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[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).

Travertine calcium carbonate deposits from a hot spring

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 with a maximum CaCO₃ solubility of 47 mg/L.

• As ambient CO₂ partial pressure is reduced below atmospheric levels, the solution becomes more and more alkaline. At extremely low $\scriptstyle P_{\text{CO}_2}$, 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. Note that for $\scriptstyle P_{\text{CO}_2}$ = 10⁻¹² atm, the [Ca²⁺][OH⁻]² product is still below the solubility product of Ca(OH)₂ (8×10⁻⁶). For still lower CO₂ pressure, Ca(OH)₂ precipitation will occur before CaCO₃ precipitation.

• 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

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 sodium bisulfate NaHSO₄ to decrease the pH or of sodium bicarbonate NaHCO₃ to increase it). From the above equations for the solubility product, the hydration reaction and the two acid reactions, the following expression for the maximum [Ca²⁺] can be easily deduced:

$[\text{Ca}^{2+}]_\text{max} = \frac{K_\text{sp}} {K_\text{h}K_\text{a1}K_\text{a2}k_\text{H}} \frac{[\text{H}^+]^2}{P_{\text{CO}_2}}$

showing a quadratic dependence in [H⁺]. The numerical application with the above values of the constants gives^{[citation needed]}

pH		7.0	7.2	7.4	7.6	7.8	8.0	8.2	8.27	8.4
[Ca2+]max (10−6mol/L)		180	71.7	28.5	11.4	4.52	1.80	0.717	0.519	0.285
[Ca2+]max (mg/L)		7.21	2.87	1.14	0.455	0.181	0.0721	0.0287	0.0208	0.0114

Comments:

• decreasing the pH from 8 to 7 increases the maximum Ca²⁺ concentration by a factor 100. Water with a pH maintained to 7 can dissolve up to 15.9 mg/L of CaCO₃. This explains the high Ca²⁺ concentration in some mineral waters with pH close to 7.
• 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/ClO⁻ 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 (water with a Ca²⁺ concentration above 120 mg/L is considered very hard). Addition of a calcium sequestering agent or complete renewing of the water will solve the problem.

Solubility in a strong or weak acid solution

Solutions of strong (HCl), moderately strong (sulfamic) or weak (acetic, citric, sorbic, lactic, phosphoric) acids are commercially available. They are commonly used as descaling agents 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 acid 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		0.00	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)		50.0	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, one can 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 pK_A = 4.76) with decreasing total acid 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

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 than the pK_A, 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)		62.0	7.39	0.874	0.123	0.0536	0.0477	0.0471	0.0471	0.0470

where [A] = [H₃PO₄] + [H₂PO₄⁻] + [HPO₄²⁻] + [PO₄³⁻] is the total acid concentration. Thus 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

• Cuttlebone
• Cuttlefish
• Gesso
• Limescale
• Marble
• Ocean acidification

References

1. ^ Patnaik, Pradyot (2003). Handbook of Inorganic Chemical Compounds. McGraw-Hill. ISBN 0070494398. http://books.google.com/?id=Xqj-TTzkvTEC. Retrieved 2009-06-06. 
2. ^ "Occupational safety and health guidline for calcium carbonate". US Dept. of Health and Human Services. http://www.cdc.gov/niosh/docs/81-123/pdfs/0090.pdf. Retrieved 31 March 2011. 
3. ^ "Solvay Precipitated Calcium Carbonate: Production". Solvay S. A.. 2007-03-09. http://www.solvaypcc.com/safety_environment/0,0,1000044-_EN,00.html. Retrieved 2007-12-30. 
4. ^ ^{a b c} Trexler, D., 2001, Two Medicine Formation, Montana: geology and fauna: In: Mesozoic Vertebrate Life, edited by Tanke, D. H., and Carpenter, K., Indiana University Press, pp. 298–309. ISBN 0253339073
5. ^ "Blast Furnace". Science Aid. http://www.scienceaid.co.uk/chemistry/industrial/blastfurnace.html. Retrieved 2007-12-30. 
6. ^ "Precipitated Calcium Carbonate uses". http://www.aristocratholding.com/calris-5.html. 
7. ^ ^{a b c d} "Calcium Carbonate Powder". Reade Advanced Materials. 2006-02-04. http://www.reade.com/Products/Minerals_and_Ores/calcium_carbonate.html. Retrieved 2007-12-30. 
8. ^ ^{a b} "Calcium carbonate in plastic applications". Imerys Performance Minerals. http://www.imerys-perfmins.com/calcium-carbonate/eu/calcium-carbonate-plastic.htm. Retrieved 2008-08-01. 
9. ^ "Ohio Historical Society Blog: Make It Shine". Ohio Historical Society. http://ohiohistory.wordpress.com/2011/06/02/making-it-shine/. 
10. ^ "Calcium Carbonate". Medline Plus. National Institutes of Health. 2005-10-01. Archived from the original on 2007-10-17. http://web.archive.org/web/20071017031324/http://www.nlm.nih.gov/medlineplus/druginfo/medmaster/a601032.html. Retrieved 2007-12-30. 
11. ^ Herbert A. Lieberman, Leon Lachman, Joseph B. Schwartz (1990). Pharmaceutical Dosage Forms: Tablets. New York: Dekker. p. 153. ISBN 0824780442. 
12. ^ Food Additives – Names Starting with C
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External links

• International Chemical Safety Card 1193
• CID 516889 from PubChem
• ATC codes: A02AC01 and A12AA04
• The British Calcium Carbonate Association – What is calcium carbonate