Calcium carbonate

Calcium carbonate is a chemical compound with the formula CaCO3. It is a common substance found in rocks as the minerals calcite and aragonite (most notably as limestone, which is a type of sedimentary rock consisting mainly of calcite) and is the main component of pearls and the shells of marine organisms, snails, and eggs. Calcium carbonate is the active ingredient in agricultural lime and is created when calcium ions in hard water react with carbonate ions to create limescale. It is medicinally used as a calcium supplement or as an antacid, but excessive consumption can be hazardous.

Calcium carbonate
Calcium carbonate
Calcium carbonate
IUPAC name
Calcium carbonate
Other names
3D model (JSmol)
ECHA InfoCard 100.006.765
EC Number 207-439-9
E number E170 (colours)
RTECS number FF9335000
Molar mass 100.0869 g/mol
Appearance Fine white powder; chalky taste
Odor odorless
Density 2.711 g/cm3 (calcite)
2.83 g/cm3 (aragonite)
Melting point 1,339 °C (2,442 °F; 1,612 K) (calcite)
825 °C (1,517 °F; 1,098 K) (aragonite)[4][5]
Boiling point decomposes
0.013 g/L (25 °C)[1][2]
Solubility in dilute acids soluble
Acidity (pKa) 9.0
−3.82×10−5 cm3/mol
93 J·mol−1·K−1[6]
−1207 kJ·mol−1[6]
A02AC01 (WHO) A12AA04 (WHO)
Safety data sheet ICSC 1193
NFPA 704
Flammability code 0: Will not burn. E.g., waterHealth code 0: Exposure under fire conditions would offer no hazard beyond that of ordinary combustible material. E.g., sodium chlorideReactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g., liquid nitrogenSpecial hazards (white): no codeNFPA 704 four-colored diamond
Lethal dose or concentration (LD, LC):
6450 mg/kg (oral, rat)
US health exposure limits (NIOSH):
PEL (Permissible)
TWA 15 mg/m3 (total) TWA 5 mg/m3 (resp)[7]
Related compounds
Other anions
Calcium bicarbonate
Other cations
Magnesium carbonate
Strontium carbonate
Barium carbonate
Related compounds
Calcium sulfate
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
Crystal structure of calcite


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

CaCO3(s) + 2 H+(aq) → Ca2+(aq) + CO2(g) + H2O(l)
CaCO3(s) → CaO(s) + CO2(g)

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

CaCO3(s) + CO2(g) + H2O(l) → Ca(HCO3)2(aq)

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

An unusual form of calcium carbonate is the hexahydrate, ikaite, CaCO3·6H2O. Ikaite is stable only below 8 °C.


The vast majority of calcium carbonate used in industry is extracted by mining or quarrying. Pure calcium carbonate (such as for food or pharmaceutical use), can be produced from a pure quarried source (usually marble).

Alternatively, calcium carbonate is prepared from calcium oxide. Water is added to give calcium hydroxide then carbon dioxide is passed through this solution to precipitate the desired calcium carbonate, referred to in the industry as precipitated calcium carbonate (PCC):[8]

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


The thermodynamically stable form of CaCO3 under normal conditions is hexagonal β-CaCO3 (the mineral calcite).[9] Other forms can be prepared, the denser (2.83 g/cm3) orthorhombic λ-CaCO3 (the mineral aragonite) and μ-CaCO3, occurring as the mineral vaterite.[9] The aragonite form can be prepared by precipitation at temperatures above 85 °C, the vaterite form can be prepared by precipitation at 60 °C.[9] Calcite contains calcium atoms coordinated by six oxygen atoms, in aragonite they are coordinated by nine oxygen atoms.[9] The vaterite structure is not fully understood.[10] Magnesium carbonate (MgCO3) has the calcite structure, whereas strontium carbonate and barium carbonate (SrCO3 and BaCO3) adopt the aragonite structure, reflecting their larger ionic radii.[9]


Calcite is the most stable polymorph of calcium carbonate. It is transparent to opaque. A transparent variety called Iceland spar (shown here) is used for optical purposes.

Geological sources

Calcite, aragonite and vaterite are pure calcium carbonate minerals. Industrially important source rocks which are predominantly calcium carbonate include limestone, chalk, marble and travertine.

Biological sources

Calcium carbonate chunks
Calcium carbonate chunks from clamshell

Eggshells, snail shells and most seashells are predominantly calcium carbonate and can be used as industrial sources of that chemical.[11] Oyster shells have enjoyed recent recognition as a source of dietary calcium, but are also a practical industrial source.[12][13] Dark green vegetables such as broccoli and kale contain dietarily significant amounts of calcium carbonate, however, they are not practical as an industrial source.[14]


Beyond Earth, strong evidence suggests the presence of calcium carbonate on Mars. Signs of calcium carbonate have been detected at more than one location (notably at Gusev and Huygens craters). This provides some evidence for the past presence of liquid water.[15][16]


Carbonate, is found frequently in geologic settings and constitutes an enormous carbon reservoir. Calcium carbonate occurs as aragonite, calcite and dolomite as significant constituents of the calcium cycle. The carbonate minerals form the rock types: limestone, chalk, marble, travertine, tufa, and others.

In warm, clear tropical waters corals are more abundant 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. Cold-water carbonates do exist at higher latitudes but have a very slow growth rate. The calcification processes are changed by ocean acidification.

Where the oceanic crust is subducted under a continental plate sediments will be carried down to warmer zones in the asthenosphere and lithosphere. Under these conditions calcium carbonate decomposes to produce carbon dioxide which, along with other gases, give rise to explosive volcanic 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 carbonate compensation depth can range from 4,000 to 6,000 meters below sea level.

Role in taphonomy

Calcium carbonate can preserve fossils through permineralization. Most of the vertebrate fossils of the Two Medicine Formation—a geologic formation known for its duck-billed dinosaur eggs—are preserved by CaCO3 permineralization.[17] This type of preservation conserves high levels of detail, even down to the microscopic level. However, it also leaves specimens vulnerable to weathering when exposed to the surface.[17]

Trilobite populations were once thought to have composed the majority of aquatic life during the Cambrian, due to the fact that their calcium carbonate-rich shells were more easily preserved than those of other species,[18] which had purely chitinous shells.


Industrial applications

The main use of calcium carbonate is in the construction industry, either as a building material, or limestone aggregate for road building, as an ingredient of cement, or as the starting material for the preparation of builders' lime by burning in a kiln. However, because of weathering mainly caused by acid rain,[19] calcium carbonate (in limestone form) is no longer used for building purposes on its own, but 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.[20]

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.[21]

It is also used as a raw material in the refining of sugar from sugar beet; it is calcined in a kiln with anthracite to produce calcium oxide and carbon dioxide. This burnt lime is then slaked in fresh water to produce a calcium hydroxide suspension for the precipitation of impurities in raw juice during carbonatation.[22]

Calcium carbonate in the form of chalk 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. 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.[23]

Fine ground calcium carbonate (GCC) is an essential ingredient in the microporous film used in 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 and PCC are used as a filler in paper because they are cheaper than wood fiber. In terms of market volume, GCC are the most important types of fillers currently used.[24] 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 used for paper filling and paper coatings is precipitated and prepared in a variety of shapes and sizes having characteristic narrow particle size distributions and equivalent spherical diameters of 0.4 to 3 micrometers.

Calcium carbonate is widely used as an extender in paints,[25] 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.[25] Some typical examples include around 15 to 20% loading of chalk in unplasticized polyvinyl chloride (uPVC) drainpipes, 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 usage temperatures.[26] Here the percentage is often 20–40%. It also routinely used as a filler in thermosetting resins (sheet and bulk molding compounds)[26] and has also been mixed with ABS, and other ingredients, to form some types of compression molded "clay" poker chips.[27] Precipitated calcium carbonate, made by dropping calcium oxide into water, is used by itself or with additives as a white paint, known as whitewashing.[28][29]

Calcium carbonate is added to a wide range of trade and do it yourself adhesives, sealants, and decorating fillers.[25] 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 ceramic glaze applications, calcium carbonate is known as whiting,[25] 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, 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.[30]

Health and dietary applications

500 mg calcium supplements with vitamin D
500-milligram calcium supplements made from calcium carbonate

Calcium carbonate is widely used medicinally as an inexpensive dietary calcium supplement for gastric antacid[31] (such as Tums). 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.[32]

Calcium carbonate is used in the production of calcium oxide as well as toothpaste and has seen a resurgence as a food preservative and color retainer, when used in or with products such as organic apples.[33]

Excess calcium from supplements, fortified food and high-calcium diets, can cause 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 hypercalcaemia, 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. Since the 1990s it has been most frequently reported in women taking calcium supplements above the recommended range of 1.2 to 1.5 grams daily, for prevention and treatment of osteoporosis,[34][35] 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.[36]

As a food additive it is designated E170,[37] and it has an INS number of 170. Used as an acidity regulator, anticaking agent, stabilizer or color it is approved for usage in the EU,[38] USA[39] and Australia and New Zealand.[40] It is used in some soy milk and almond 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.[41] Calcium carbonate is also used as a firming agent in many canned and bottled vegetable products.

Agricultural use

Agricultural lime, powdered chalk or limestone, is used as a cheap method for neutralising acidic soil, making it suitable for planting.[42]

Household use

Calcium carbonate is a key ingredient in many household cleaning powders like Comet and is used as a scrubbing agent.

Environmental applications

In 1989, a researcher, Ken Simmons, introduced CaCO3 into the Whetstone Brook in Massachusetts.[43] 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 amount 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.[44][45][46] 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.[47]

Calcium carbonate is also used in flue gas desulfurisation applications eliminating harmful SO2 and NO2 emissions from coal and other fossil fuels burnt in large fossil fuel power stations.[44]

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.[48]

The table shows that this partial 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) versus temperature (T).[49]
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


With varying CO2 pressure

Travertine calcium carbonate deposits from a hot spring

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+ + CO2−
Ksp = 3.7×10−9 to 8.7×10−9 at 25 °C

where the solubility product for [Ca2+][CO2−
is given as anywhere from Ksp = 3.7×10−9 to Ksp = 8.7×10−9 at 25 °C, depending upon the data source.[49][50] 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 CO2−
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 CO2−
combines with H+ in the solution according to:

⇌ H+ + CO2−
Ka2 = 5.61×10−11 at 25 °C

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 HCO
combines with H+ in solution according to:

H2CO3 ⇌ H+ + HCO
Ka1 = 2.5×10−4 at 25 °C

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

H2O + CO2(aq) ⇌ 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), PCO2 being the CO2 partial pressure.

For ambient air, PCO2 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 PCO2, 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 concentration. 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

Calcium ion solubility as a function of CO2 partial pressure at 25 °C (Ksp = 4.47×10−9)
PCO2 (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
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+] = [HCO
] + 2 [CO2−
] + [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 initial water solvent pH is not neutral, the equation is modified).

The adjacent table 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 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 PCO2, 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 PCO2 = 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, temperature and salinity: CaCO3 scaling in swimming pools


In contrast to the open equilibrium scenario above, many swimming pools are managed by addition of sodium bicarbonate (NaHCO3) to about 2 mM as a buffer, then control of pH through use of HCl, NaHSO4, Na2CO3, NaOH or chlorine formulations that are acidic or basic. In this situation, dissolved inorganic carbon (total inorganic carbon) is far from equilibrium with atmospheric CO2. Progress towards equilibrium through outgassing of CO2 is slowed by

  1. the slow reaction
    H2CO3 ⇌ CO2(aq) + H2O;[51]
  2. limited aeration in a deep water column; and
  3. periodic replenishment of bicarbonate to maintain buffer capacity (often estimated through measurement of ‘total alkalinity’).

In this situation, the dissociation constants for the much faster reactions

H2CO3 ⇌ H+ + HCO
⇌ 2 H+ + CO2−

allow the prediction of concentrations of each dissolved inorganic carbon species in solution, from the added concentration of HCO
(which constitutes more than 90% of Bjerrum plot species from pH 7 to pH 8 at 25 °C in fresh water).[52] Addition of HCO
will increase CO2−
concentration at any pH. Rearranging the equations given above, we can see that [Ca2+] = Ksp/[CO2−
, and [CO2−
] = Ka2 [HCO
. Therefore, when HCO
concentration is known, the maximum concentration of Ca2+ ions before scaling through CaCO3 precipitation can be predicted from the formula:

The solubility product for CaCO3 (Ksp) and the dissociation constants for the dissolved inorganic carbon species (including Ka2) are all substantially affected by temperature and salinity,[52] with the overall effect that [Ca2+]max increases from freshwater to saltwater, and decreases with rising temperature, pH, or added bicarbonate level, as illustrated in the accompanying graphs.

The trends are illustrative for pool management, but whether scaling occurs also depends on other factors including interactions with Mg2+, B(OH)
and other ions in the pool, as well as supersaturation effects.[53][54] Scaling is commonly observed in electrolytic chlorine generators, where there is a high pH near the cathode surface and scale deposition further increases temperature. This is one reason that some pool operators prefer borate over bicarbonate as the primary pH buffer, and avoid the use of pool chemicals containing calcium.[55]

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/mol):
[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/L 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 [Ca2+] ≈ 1/2 [A]. When the concentration decreases, [HCO
] 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 pKa = 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/L 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 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 [HCO
    ], [CO2−
    ], [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/L of acid)
62.0 7.39 0.874 0.123 0.0536 0.0477 0.0471 0.0471 0.0470
where [A] = [H3PO4] + [H
] + [HPO2−
] + [PO3−
] 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 [HPO2−
] is not negligible (see phosphoric acid).

See also


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External links


Aragonite is a carbonate mineral, one of the three most common naturally occurring crystal forms of calcium carbonate, CaCO3 (the other forms being the minerals calcite and vaterite). It is formed by biological and physical processes, including precipitation from marine and freshwater environments.

The crystal lattice of aragonite differs from that of calcite, resulting in a different crystal shape, an orthorhombic crystal system with acicular crystal. Repeated twinning results in pseudo-hexagonal forms. Aragonite may be columnar or fibrous, occasionally in branching stalactitic forms called flos-ferri ("flowers of iron") from their association with the ores at the Carinthian iron mines.

Biological pump

The biological pump, in its simplest form, is the ocean's biologically driven sequestration of carbon from the atmosphere to the ocean interior and seafloor sediments. It is the part of the oceanic carbon cycle responsible for the cycling of organic matter formed mainly by phytoplankton during photosynthesis (soft-tissue pump), as well as the cycling of calcium carbonate (CaCO3) formed into shells by certain organisms such as plankton and mollusks (carbonate pump).


Calcareous is an adjective meaning "mostly or partly composed of calcium carbonate", in other words, containing lime or being chalky. The term is used in a wide variety of scientific disciplines.

Calcareous sponge

The calcareous sponges of class Calcarea are members of the animal phylum Porifera, the cellular sponges. They are characterized by spicules made out of calcium carbonate in the form of calcite or aragonite. While the spicules in most species have three points, in some species they have either two or four points.


Calcite is a carbonate mineral and the most stable polymorph of calcium carbonate (CaCO3). The Mohs scale of mineral hardness, based on scratch hardness comparison, defines value 3 as "calcite".

Other polymorphs of calcium carbonate are the minerals aragonite and vaterite. Aragonite will change to calcite over timescales of days or less at temperatures exceeding 300 °C, and vaterite is even less stable.

Calcium bicarbonate

Calcium bicarbonate, also called calcium hydrogen carbonate, has a chemical formula Ca(HCO3)2. The term does not refer to a known solid compound; it exists only in aqueous solution containing the calcium (Ca2+), bicarbonate (HCO−3), and carbonate (CO2−3) ions, together with dissolved carbon dioxide (CO2). The relative concentrations of these carbon-containing species depend on the pH; bicarbonate predominates within the range 6.36–10.25 in fresh water.

All waters in contact with the atmosphere absorb carbon dioxide, and as these waters come into contact with rocks and sediments they acquire metal ions, most commonly calcium and magnesium, so most natural waters that come from streams, lakes, and especially wells, can be regarded as dilute solutions of these bicarbonates. These hard waters tend to form carbonate scale in pipes and boilers and they react with soaps to form an undesirable scum.

Attempts to prepare compounds such as solid calcium bicarbonate by evaporating its solution to dryness invariably yield instead the solid calcium carbonate:

Ca(HCO3)2(aq) → CO2(g) + H2O(l) + CaCO3(s).Very few solid bicarbonates other than those of the alkali metals and ammonium ion are known to exist.

The above reaction is very important to the formation of stalactites, stalagmites, columns, and other speleothems within caves, and for that matter, in the formation of the caves themselves. As water containing carbon dioxide (including extra CO2 acquired from soil organisms) passes through limestone or other calcium carbonate-containing minerals, it dissolves part of the calcium carbonate, hence becomes richer in bicarbonate. As the groundwater enters the cave, the excess carbon dioxide is released from the solution of the bicarbonate, causing the much less soluble calcium carbonate to be deposited.

In the reverse process, dissolved carbon dioxide (CO2) in rainwater (H2O) reacts with limestone calcium carbonate (CaCO3) to form soluble calcium bicarbonate (Ca(HCO3)2). This soluble compound is then washed away with the rainwater. This form of weathering is called carbonation.

In medicine, calcium bicarbonate is sometimes administered intravenously to immediately correct the cardiac depressor effects of hypokalemia by increasing calcium concentration in serum, and at the same time, correcting the acid usually present.


Caliche () is a sedimentary rock, a hardened natural cement of calcium carbonate that binds other materials—such as gravel, sand, clay, and silt. It occurs worldwide, in aridisol and mollisol soil orders—generally in arid or semiarid regions, including in central and western Australia, in the Kalahari Desert, in the High Plains of the western USA, in the Sonoran Desert and Mojave Desert, and in Eastern Saudi Arabia Al-Hasa. Caliche is also known as calcrete or kankar (in India). It belongs to the duricrusts. The term caliche is Spanish and is originally from the Latin calx, meaning lime.

Caliche is generally light-colored, but can range from white to light pink to reddish-brown, depending on the impurities present. It generally occurs on or near the surface, but can be found in deeper subsoil deposits, as well. Layers vary from a few inches to feet thick, and multiple layers can exist in a single location.

In northern Chile and Peru, caliche also refers to mineral deposits that include nitrate salts. Caliche can also refer to various claylike deposits in Mexico and Colombia. In addition, it has been used to describe some forms of quartzite, bauxite, kaolinite, laterite, chalcedony, opal, and soda niter.

A similar material, composed of calcium sulfate rather than calcium carbonate, is called gypcrust.


Chalk is a soft, white, porous, sedimentary carbonate rock, a form of limestone composed of the mineral calcite. Calcite is an ionic salt called calcium carbonate or CaCO3. It forms under reasonably deep marine conditions from the gradual accumulation of minute calcite shells (coccoliths) shed from micro-organisms called coccolithophores. Flint (a type of chert) is very common as bands parallel to the bedding or as nodules embedded in chalk. It is probably derived from sponge spicules or other siliceous organisms as water is expelled upwards during compaction. Flint is often deposited around larger fossils such as Echinoidea which may be silicified (i.e. replaced molecule by molecule by flint).

Chalk as seen in Cretaceous deposits of Western Europe is unusual among sedimentary limestones in the thickness of the beds. Most cliffs of chalk have very few obvious bedding planes unlike most thick sequences of limestone such as the Carboniferous Limestone or the Jurassic oolitic limestones. This presumably indicates very stable conditions over tens of millions of years.

Chalk has greater resistance to weathering and slumping than the clays with which it is usually associated, thus forming tall, steep cliffs where chalk ridges meet the sea. Chalk hills, known as chalk downland, usually form where bands of chalk reach the surface at an angle, so forming a scarp slope. Because chalk is well jointed it can hold a large volume of ground water, providing a natural reservoir that releases water slowly through dry seasons.


Coccoliths are individual plates of calcium carbonate formed by coccolithophores (single-celled algae such as Emiliania huxleyi) which are arranged around them in a coccosphere.

Elemental calcium

Elemental calcium is a term used on dietary supplement labels to refer to the amount of calcium in a product. Calcium pills contain calcium in a variety of molecules, such as calcium carbonate, calcium citrate, calcium citrate-maleate, etc. Each pill supplies a different amount of elemental calcium. For example, calcium carbonate is 40% elemental calcium by weight and calcium citrate is about 20%. Thus a 500 mg pill of calcium carbonate contains 200 mg of calcium and the container will indicate each pill has 200 mg of elemental calcium. This is the actual calcium content.

Hard water

Hard water is water that has high mineral content (in contrast with "soft water"). Hard water is formed when water percolates through deposits of limestone and chalk which are largely made up of calcium and magnesium carbonates.

Hard drinking water may have moderate health benefits, but can pose critical problems in industrial settings, where water hardness is monitored to avoid costly breakdowns in boilers, cooling towers, and other equipment that handles water. In domestic settings, hard water is often indicated by a lack of foam formation when soap is agitated in water, and by the formation of limescale in kettles and water heaters. Wherever water hardness is a concern, water softening is commonly used to reduce hard water's adverse effects.


Kankar or kunkur is a sedimentological term derived from Hindi, occasionally applied in India and the United States to detrital or residual rolled, often nodular calcium carbonate formed in soils of semi-arid regions.

It forms sheets across alluvial plains and can occur as discontinuous lines of nodular kankar or as indurated layers in stratigraphic profiles more commonly referred to as calcrete, hardpan or duricrust.

Lime (material)

Lime is a calcium-containing inorganic mineral composed primarily of oxides, and hydroxide, usually calcium oxide and/ or calcium hydroxide. It is also the name for calcium oxide which occurs as a product of coal seam fires and in altered limestone xenoliths in volcanic ejecta. The word lime originates with its earliest use as building mortar and has the sense of sticking or adhering.These materials are still used in large quantities as building and engineering materials (including limestone products, cement, concrete, and mortar), as chemical feedstocks, and for sugar refining, among other uses. Lime industries and the use of many of the resulting products date from prehistoric times in both the Old World and the New World. Lime is used extensively for wastewater treatment with ferrous sulfate.

The rocks and minerals from which these materials are derived, typically limestone or chalk, are composed primarily of calcium carbonate. They may be cut, crushed, or pulverized and chemically altered. Burning (calcination) converts them into the highly caustic material quicklime (calcium oxide) and, through subsequent addition of water, into the less caustic (but still strongly alkaline) slaked lime or hydrated lime (calcium hydroxide, Ca(OH)2), the process of which is called slaking of lime. Lime kilns are the kilns used for lime burning and slaking.

When the term is encountered in an agricultural context, it usually refers to agricultural lime, which is crushed limestone, not a product of a lime kiln. Otherwise it most commonly means slaked lime, as the more dangerous form is usually described more specifically as quicklime or burnt lime.


Limescale is the hard, off-white, chalky deposit found in kettles, hot-water boilers and the inside of hot water pipework.

It is also often found as a similar deposit on the inner surface of old pipe and other surfaces where "hard water" has evaporated. In addition to being unsightly and hard to clean, limescale seriously impairs the operation or damages various components.The type found deposited on the heating elements of water heaters has a main component of calcium carbonate. Hard water contains calcium (and often magnesium) bicarbonate or similar ions. Calcium salts, such as calcium bicarbonate and calcium carbonate are both more soluble in hot water than cold water. Thus, heating water does not cause calcium carbonate to precipitate per se. However, there is an equilibrium between dissolved calcium bicarbonate and dissolved calcium carbonate:

Ca2+ + 2HCO3− ⇋ Ca2+ + CO32− + CO2 + H2O

where the equilibrium is driven by the carbonate/bicarbonate, not the calcium. Note that the CO2 is dissolved in the water.

There is also an equilibrium of carbon dioxide between dissolved in water (dis) and the gaseous state (g):

CO2(dis) ⇋ CO2(g)The equilibrium of CO2 also moves to the right towards gaseous CO2 when the water temperature rises. When water that contains dissolved calcium carbonate is warmed, CO2 is removed from the water as gas causing the equilibrium of bicarbonate and carbonate to shift to the right, increasing the concentration of dissolved carbonate. As the concentration of carbonate increases, calcium carbonate precipitates as the salt: Ca2+ + CO32− ⇋ CaCO3.

As new cold water with dissolved calcium carbonate/bicarbonate is added and heated, CO2 gas is removed, carbonate concentration increases, and more calcium carbonate precipitates.

Descaling agents are used to remove scale. Prevention of scale build-up relies on the technologies of water softening.

Lindlar catalyst

A Lindlar catalyst is a heterogeneous catalyst that consists of palladium deposited on calcium carbonate which is then poisoned with various forms of lead or sulphur. It is used for the hydrogenation of alkynes to alkenes (i.e. without further reduction into alkanes) and is named after its inventor Herbert Lindlar.


Marl or marlstone is a calcium carbonate or lime-rich mud or mudstone which contains variable amounts of clays and silt. The dominant carbonate mineral in most marls is calcite, but other carbonate minerals such as aragonite, dolomite, and siderite may be present. Marl was originally an old term loosely applied to a variety of materials, most of which occur as loose, earthy deposits consisting chiefly of an intimate mixture of clay and calcium carbonate, formed under freshwater conditions; specifically an earthy substance containing 35–65% clay and 65–35% carbonate. It also describes a habit of coralline red alga. The term is today often used to describe indurated marine deposits and lacustrine (lake) sediments which more accurately should be named 'marlstone'. Marlstone is an indurated (resists crumbling or powdering) rock of about the same composition as marl, more correctly called an earthy or impure argillaceous limestone. It has a blocky subconchoidal fracture, and is less fissile than shale. The term 'marl' is widely used in English-language geology, while the terms Mergel and Seekreide (German for "lake chalk") are used in European references.

The lower stratigraphic units of the chalk cliffs of Dover consist of a sequence of glauconitic marls followed by rhythmically banded limestone and marl layers. Upper Cretaceous cyclic sequences in Germany and marl–opal-rich Tortonian-Messinian strata in the Sorbas basin related to multiple sea drawdown have been correlated with Milankovitch orbital forcing.Marl as lacustrine sediment is common in post-glacial lake-bed sediments, often found underlying peat bogs. It has been used as a soil conditioner and acid soil neutralizing agent.


A seashell or sea shell, also known simply as a shell, is a hard, protective outer layer created by an animal that lives in the sea. The shell is part of the body of the animal. Empty seashells are often found washed up on beaches by beachcombers. The shells are empty because the animal has died and the soft parts have been eaten by another animal or have decomposed.

A seashell is usually the exoskeleton of an invertebrate (an animal without a backbone), and is typically composed of calcium carbonate or chitin. Most shells that are found on beaches are the shells of marine mollusks, partly because these shells are usually made of calcium carbonate, and endure better than shells made of chitin.

Apart from mollusk shells, other shells that can be found on beaches are those of barnacles, horseshoe crabs and brachiopods. Marine annelid worms in the family Serpulidae create shells which are tubes made of calcium carbonate cemented onto other surfaces. The shells of sea urchins are called "tests", and the moulted shells of crabs and lobsters are exuviae. While most seashells are external, some cephalopods have internal shells.

Seashells have been used by humans for many different purposes throughout history and pre-history. However, seashells are not the only kind of shells; in various habitats, there are shells from freshwater animals such as freshwater mussels and freshwater snails, and shells of land snails.

Soft water

Soft water is surface water that contains low concentrations of ions and in particular is low in ions of calcium and magnesium. Soft water naturally occurs where rainfall and the drainage basin of rivers are formed of hard, impervious and calcium-poor rocks. Examples in the UK (United Kingdom) include Snowdonia in Wales and the Western Highlands in Scotland.

The term may also be used to describe water that has been produced by a water softening process although such water is more correctly termed softened water. In these cases the water may also contain elevated levels of sodium and bicarbonate ions.

Because soft water has few calcium ions, there is no inhibition of the lathering action of soaps and no soap scum is formed in normal washing. Similarly, soft water produces no calcium deposits in water heating systems. Water that is not soft is referred to as hard water.

In the UK, water is regarded as soft if the hardness is less than 50 mg/l of calcium carbonate. Water containing more than 50 mg/l of calcium carbonate is termed hard water. In the United States soft water is classified as having less than 60 mg/l of calcium carbonate.In USA, due to ancient sea beds with high limestone (calcium carbonate) concentrations, as much as 85% of water is hard, and many needing water softening treatment. Waters in Eastern and Southern Florida, North-western Texas the entire South Dakota are considered extremely hard, with the metropolitan cities hardest waters include Indianapolis, Las Vegas, Minneapolis, Phoenix, San Antonio and Tampa. Only the entire Mississippi and Maine have soft water.

Stone paper

Stone paper, also known as limestone paper, rock paper, bio-plastic paper, mineral paper and rich mineral paper, is a type of strong and durable paper-like material manufactured from calcium carbonate bonded with small amount of resin high-density polyethylene (HDPE). It is used for stationery, leaflets, posters, books, magazines, bags, packaging, wallpaper, adhesives, tags, in-mould labels, plates, trays, containers, and maps among other uses.

Calcium compounds
Magnesium (increases motility)
Aluminium (decreases motility)
Combinations and complexes
of aluminium, calcium and magnesium
Drugs for treatment of hyperkalemia and hyperphosphatemia (V03AE)
Potassium binders
Phosphate binders

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