Calcium carbonate

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Calcium carbonate
Calcium carbonate.png
Calcium-carbonate-xtal-3D-SF.png
   Calcium, Ca
   Carbon, C
   Oxygen, O
Calcium carbonate.jpg
Names
IUPAC name
Calcium carbonate
Other names
Identifiers
3D model (JSmol)
ChEBI
ChEMBL
ChemSpider
DrugBank
ECHA InfoCard 100.006.765 OOjs UI icon edit-ltr-progressive.svg
EC Number
  • 207-439-9
E number E170 (colours)
KEGG
PubChem CID
RTECS number
  • FF9335000
UNII
  • InChI=1S/CH2O3.Ca/c2-1(3)4;/h(H2,2,3,4);/q;+2/p-2 Yes check.svgY
    Key: VTYYLEPIZMXCLO-UHFFFAOYSA-L Yes check.svgY
  • InChI=1/CH2O3.Ca/c2-1(3)4;/h(H2,2,3,4);/q;+2/p-2
    Key: VTYYLEPIZMXCLO-NUQVWONBAS
  • [Ca+2].[O-]C([O-])=O
  • C(=O)([O-])[O-].[Ca+2]
Properties
CaCO3
Molar mass 100.0869 g/mol
AppearanceFine white powder or colorless crystals; 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) [1] [2]
Boiling point decomposes
0.013 g/L (25 °C) [3] [4]
3.3×10−9 [5]
Solubility in dilute acidssoluble
Acidity (pKa)9.0
−3.82×10−5 cm3/mol
1.59
Structure
Trigonal
32/m
Thermochemistry
Std molar
entropy (S⦵298)
93 J/(mol·K) [6]
−1207 kJ/mol [6]
Pharmacology
A02AC01 ( WHO ) A12AA04 ( WHO )
Hazards
NFPA 704 (fire diamond)
NFPA 704.svgHealth 0: Exposure under fire conditions would offer no hazard beyond that of ordinary combustible material. E.g. sodium chlorideFlammability 0: Will not burn. E.g. waterInstability 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no code
0
0
0
Lethal dose or concentration (LD, LC):
6450 mg/kg (oral, rat)
NIOSH (US health exposure limits):
PEL (Permissible)
TWA 15 mg/m3 (total) TWA 5 mg/m3 (resp) [7]
Safety data sheet (SDS) ICSC 1193
Related compounds
Other anions
Calcium bicarbonate
Other cations
Related compounds
 Calcium sulfate
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
X mark.svgN  verify  (what is  Yes check.svgYX mark.svgN ?)
Crystal structure of calcite Calcite.png
Crystal structure of calcite

Calcium carbonate is a chemical compound with the chemical formula Ca CO3 . It is a common substance found in rocks as the minerals calcite and aragonite, most notably in chalk and limestone, eggshells, gastropod shells, shellfish skeletons and pearls. Materials containing much calcium carbonate or resembling it are described as calcareous. Calcium carbonate is the active ingredient in agricultural lime and is produced when calcium ions in hard water react with carbonate ions to form limescale. It has medical use as a calcium supplement or as an antacid, but excessive consumption can be hazardous and cause hypercalcemia and digestive issues. [8]

Chemistry

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

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

Calcium carbonate reacts 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.

Preparation

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) This process is called carbonatation: [9]

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

In a laboratory, calcium carbonate can easily be crystallized from calcium chloride (CaCl2), by placing an aqueous solution of CaCl2 in a desiccator alongside ammonium carbonate [NH4]2CO3. [10] In the desiccator, ammonium carbonate is exposed to air and decomposes into ammonia, carbon dioxide, and water. The carbon dioxide then diffuses into the aqueous solution of calcium chloride, reacts with the calcium ions and the water, and forms calcium carbonate.

Structure

The thermodynamically stable form of CaCO3 under normal conditions is hexagonal β-CaCO3 (the mineral calcite). Other forms can be prepared, the denser (2.83 g/cm3) orthorhombic λ-CaCO3 (the mineral aragonite) and hexagonal μ-CaCO3, occurring as the mineral vaterite. The aragonite form can be prepared by precipitation at temperatures above 85 °C; the vaterite form can be prepared by precipitation at 60 °C. Calcite contains calcium atoms coordinated by six oxygen atoms; in aragonite they are coordinated by nine oxygen atoms.[ citation needed ] The vaterite structure is not fully understood. [11] Magnesium carbonate (MgCO3) has the calcite structure, whereas strontium carbonate (SrCO3) and barium carbonate (BaCO3) adopt the aragonite structure, reflecting their larger ionic radii.[ citation needed ]

Polymorphs

Calcium carbonate crystallizes in three anhydrous polymorphs, [12] [13] of which calcite is the thermodynamically most stable at room temperature, aragonite is only slightly less so, and vaterite is the least stable. [14]

Crystal structure

The calcite crystal structure is trigonal, with space group R3c (No. 167 in the International Tables for Crystallography [15] ), and Pearson symbol hR10. [16] Aragonite is orthorhombic, with space group Pmcn (No 62), and Pearson Symbol oP20. [17] Vaterite is composed of at least two different coexisting crystallographic structures. The major structure exhibits hexagonal symmetry in space group P63/mmc, the minor structure is still unknown. [18]

Crystallization

Crystal structure of calcite and aragonite Calcite+Aragonite.png
Crystal structure of calcite and aragonite

All three polymorphs crystallize simultaneously from aqueous solutions under ambient conditions. [14] In additive-free aqueous solutions, calcite forms easily as the major product, while aragonite appears only as a minor product.

At high saturation, vaterite is typically the first phase precipitated, which is followed by a transformation of the vaterite to calcite. [19] This behavior seems to follow Ostwald's rule, in which the least stable polymorph crystallizes first, followed by the crystallization of different polymorphs via a sequence of increasingly stable phases. [20] However, aragonite, whose stability lies between those of vaterite and calcite, seems to be the exception to this rule, as aragonite does not form as a precursor to calcite under ambient conditions. [14]

Microscopic calcite and vaterite Calcite+Vaterite.png
Microscopic calcite and vaterite

Aragonite occurs in majority when the reaction conditions inhibit the formation of calcite and/or promote the nucleation of aragonite. For example, the formation of aragonite is promoted by the presence of magnesium ions, [21] or by using proteins and peptides derived from biological calcium carbonate. [22] Some polyamines such as cadaverine and Poly(ethylene imine) have been shown to facilitate the formation of aragonite over calcite. [14]

Selection by organisms

Organisms, such as molluscs and arthropods, have shown the ability to grow all three crystal polymorphs of calcium carbonate, mainly as protection (shells) and muscle attachments. [23] Moreover, they exhibit a remarkable capability of phase selection over calcite and aragonite, and some organisms can switch between the two polymorphs. The ability of phase selection is usually attributed to the use of specific macromolecules or combinations of macromolecules by such organisms. [24] [25] [26]

Occurrence

Calcite is the most stable polymorph of calcium carbonate. It is transparent to opaque. A transparent variety called Iceland spar (shown here) was used to create polarized light in the 19th century. Silfurberg.jpg
Calcite is the most stable polymorph of calcium carbonate. It is transparent to opaque. A transparent variety called Iceland spar (shown here) was used to create polarized light in the 19th century.

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 from clamshell Calcium carbonate chunks.JPG
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. [28] Oyster shells have enjoyed recent recognition as a source of dietary calcium, but are also a practical industrial source. [29] [30] Dark green vegetables such as broccoli and kale contain dietarily significant amounts of calcium carbonate, but they are not practical as an industrial source. [31]

Annelids in the family Lumbricidae, earthworms, possess a regionalization of the digestive track called calciferous glands, Kalkdrüsen, or glandes de Morren, that processes calcium and CO2 into calcium carbonate, which is later excreted into the dirt. [32] The function of these glands is unknown but is believed to serve as a CO2 regulation mechanism within the animals' tissues. [33] This process is ecologically significant, stabilizing the pH of acid soils. [34]

Extraterrestrial

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. [35] [36]

Geology

Surface precipitation of CaCO3 as tufa in Rubaksa, Ethiopia Rubaksa tufa plug.jpg
Surface precipitation of CaCO3 as tufa in Rubaksa, Ethiopia

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.

Tufa at Huanglong, Sichuan Water-of-Five-colored-Pond Huanglong Sichuan China.jpg
Tufa at Huanglong, Sichuan

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. Increasing pressure also increases the solubility of calcium carbonate. Calcium carbonate is unusual in that its solubility increases with decreasing temperature. [37] The carbonate compensation depth ranges from 4,000 to 6,000 meters below sea level in modern oceans, and the various polymorphs (calcite, aragonite) have different compensation depths based on their stability. [38]

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. [39] 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. [39]

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, [40] which had purely chitinous shells.

Uses

Construction

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, [41] 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. [42]

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. [43]

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. [44]

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 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. [45]

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. 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.[ citation needed ]

Calcium carbonate is widely used as an extender in paints, [46] 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. [46] Some typical examples include around 15–20% loading of chalk in unplasticized polyvinyl chloride (uPVC) drainpipes, 5–15% 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).[ citation needed ] Polypropylene compounds are often filled with calcium carbonate to increase rigidity, a requirement that becomes important at high usage temperatures. [47] Here the percentage is often 20–40%. It also routinely used as a filler in thermosetting resins (sheet and bulk molding compounds) [47] and has also been mixed with ABS, and other ingredients, to form some types of compression molded "clay" poker chips. [48] Precipitated calcium carbonate, made by dropping calcium oxide into water, is used by itself or with additives as a white paint, known as whitewashing. [49] [50]

Calcium carbonate is added to a wide range of trade and do it yourself adhesives, sealants, and decorating fillers. [46] 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. [51] [52] [53] [54]

In ceramic glaze applications, calcium carbonate is known as whiting, [46] 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. [55]

Health and diet

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

Calcium carbonate is widely used medicinally as an inexpensive dietary calcium supplement for gastric antacid [56] (such as Tums and Eno). It may be used as a phosphate binder for the treatment of hyperphosphatemia (primarily in patients with chronic kidney failure). It is used in the pharmaceutical industry as an inert filler for tablets and other pharmaceuticals. [57]

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. [58]

Calcium carbonate is used therapeutically as phosphate binder in patients on maintenance haemodialysis. It is the most common form of phosphate binder prescribed, particularly in non-dialysis chronic kidney disease. Calcium carbonate is the most commonly used phosphate binder, but clinicians are increasingly prescribing the more expensive, non-calcium-based phosphate binders, particularly sevelamer.

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 kidney 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, [59] [60] 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. [61]

As a food additive it is designated E170, [62] 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, [63] US [64] and Australia and New Zealand. [65] It is "added by law to all UK milled bread flour except wholemeal". [66] [67] It is used in some soy milk and almond milk products as a source of dietary calcium; at least one study suggests that calcium carbonate might be as bioavailable as the calcium in cow's milk. [68] Calcium carbonate is also used as a firming agent in many canned and bottled vegetable products.

Several calcium supplement formulations have been documented to contain the chemical element lead, [69] posing a public health concern. [70] Lead is commonly found in natural sources of calcium. [69]

Agriculture and aquaculture

Agricultural lime, powdered chalk or limestone, is used as a cheap method of neutralising acidic soil, making it suitable for planting, also used in aquaculture industry for pH regulation of pond soil before initiating culture. [71] There is interest in understanding whether or not it can affect pesticide adsorption and desorption in calcareous soil. [72]

Household cleaning

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

Pollution mitigation

In 1989, a researcher, Ken Simmons, introduced CaCO3 into the Whetstone Brook in Massachusetts. [73] 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. [74] [75] [76] 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. [77]

Calcium carbonate is also used in flue-gas desulfurization applications eliminating harmful SO2 and NO2 emissions from coal and other fossil fuels burnt in large fossil fuel power stations. [74]

Plastics

Calcium carbonate is commonly used in the plastic industry as a filler. When it is incorporated in a plastic material, it can improve the hardness, stiffness, dimensional stability and processability of the material. [78]

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. [79]

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). [80]
P (kPa)0.0550.130.311.805.99.3142434517280911011799013961
T (°C)55058760568072774877780083085287188189189893710821241

Solubility

With varying CO2 pressure

Travertine calcium carbonate deposits from a hot spring CanarySpring.jpg
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+ + CO2−3Ksp = 3.7×10−9 to 8.7×10−9 at 25 °C

where the solubility product for [Ca2+][CO2−3] is given as anywhere from Ksp = 3.7×10−9 to Ksp = 8.7×10−9 at 25 °C, depending upon the data source. [80] [81] 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 CO2−3 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 CO2−3 combines with H+ in the solution according to

HCO−3 ⇌ H+ + CO2−3  Ka2 = 5.61×10−11 at 25 °C

HCO−3 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−3 combines with H+ in solution according to

H2CO3 ⇌ H+ + HCO−3Ka1 = 2.5×10−4 at 25 °C

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

H2O + CO2(aq) ⇌ H2CO3Kh = 1.70×10−3 at 25 °C

And dissolved carbon dioxide is in equilibrium with atmospheric carbon dioxide according to

⁠PCO2/[CO2]⁠ = where = 29.76 atm/(mol/L) at 25 °C (Henry volatility), and PCO2 is the CO2 partial pressure.

For ambient air, PCO2 is around 3.5×10−4 atm (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 per 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)
10−1212.05.19×10−3
10−1011.31.12×10−3
10−810.72.55×10−4
10−69.831.20×10−4
10−48.623.16×10−4
3.5×10−48.274.70×10−4
10−37.966.62×10−4
10−27.301.42×10−3
10−16.633.05×10−3
15.966.58×10−3
105.301.42×10−2
H2O ⇌ H+ + OH−K = 10−14 at 25 °C

(which is true for all aqueous solutions), and the constraint that the solution must be electrically neutral, i.e., the overall charge of dissolved positive ions [Ca2+] + 2 [H+] must be cancelled out by the overall charge of dissolved negative ions [HCO−3] + [CO2−3] + [OH−], make it possible to solve simultaneously for the remaining five unknown concentrations (the previously mentioned form of the neutrality 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 balance is not neutral).

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).

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 one of the trends above. 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

CaCO3-pH.gif
CaCO3-Temp.gif

In contrast to the open equilibrium scenario above, many swimming pools are managed by addition of sodium bicarbonate (NaHCO3) to the concentration of about 2 mmol/L 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; [82]
  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−3 ⇌ 2 H+ + CO2−3

allow the prediction of concentrations of each dissolved inorganic carbon species in solution, from the added concentration of HCO−3 (which constitutes more than 90% of Bjerrum plot species from pH 7 to pH 8 at 25 °C in fresh water). [83] Addition of HCO−3 will increase CO2−3 concentration at any pH. Rearranging the equations given above, we can see that [Ca2+] = ⁠Ksp/[CO2−3]⁠, and [CO2−3] = ⁠Ka2 [HCO−3]/[H+]⁠. Therefore, when HCO−3 concentration is known, the maximum concentration of Ca2+ ions before scaling through CaCO3 precipitation can be predicted from the formula:

[Ca2+]max = ⁠Ksp/Ka2⁠ × ⁠[H+]/[HCO−3]⁠

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, [83] 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 Mg 2+, [B(OH)4]− and other ions in the pool, as well as supersaturation effects. [84] [85] 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. [86]

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.

[A] (mol/L)110−110−210−310−410−510−610−710−10
Initial pH0.001.002.003.004.005.006.006.797.00
Final pH6.757.257.758.148.258.268.268.268.27
Dissolved CaCO3 (g/L of acid)50.05.000.5140.08490.05040.04740.04710.04700.0470
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+] ≈ 0.5 [A−]. When the concentration decreases, [HCO−3] 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.
[A] (mol/L)[Ca2+] ≈ 0.5 [A−]
[ clarification needed ]		10−110−210−310−410−510−610−710−10
Initial pH2.382.883.393.914.475.156.026.797.00
Final pH6.757.257.758.148.258.268.268.268.27
Dissolved CaCO3 (g/L of acid)49.54.990.5130.08480.05040.04740.04710.04700.0470
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.
[A] (mol/L)110−110−210−310−410−510−610−710−10
Initial pH1.081.622.253.054.015.005.976.747.00
Final pH6.717.177.638.068.248.268.268.268.27
Dissolved CaCO3 (g/L of acid)62.07.390.8740.1230.05360.04770.04710.04710.0470
where [A] = [H3PO4] + [H2PO−4] + [HPO2−4] + [PO3−4] 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 [HPO2−4] is not negligible (see phosphoric acid).

See also

Electron micrograph of needle-like calcium carbonate crystals formed as limescale in a kettle Vodni kamen.jpg
Electron micrograph of needle-like calcium carbonate crystals formed as limescale in a kettle
Around 2 g of calcium-48 carbonate Calcium-48 carbonate.png
Around 2 g of calcium-48 carbonate

Related Research Articles

<span class="mw-page-title-main">Bicarbonate</span> Polyatomic anion

In inorganic chemistry, bicarbonate is an intermediate form in the deprotonation of carbonic acid. It is a polyatomic anion with the chemical formula HCO−
3.

<span class="mw-page-title-main">Carbonate</span> Salt or ester of carbonic acid

A carbonate is a salt of carbonic acid, H2CO3, characterized by the presence of the carbonate ion, a polyatomic ion with the formula CO2−3. The word "carbonate" may also refer to a carbonate ester, an organic compound containing the carbonate groupO=C(−O−)2.

<span class="mw-page-title-main">Calcite</span> Calcium carbonate mineral

Calcite is a carbonate mineral and the most stable polymorph of calcium carbonate (CaCO3). It is a very common mineral, particularly as a component of limestone. Calcite defines hardness 3 on the Mohs scale of mineral hardness, based on scratch hardness comparison. Large calcite crystals are used in optical equipment, and limestone composed mostly of calcite has numerous uses.

<span class="mw-page-title-main">Speleothem</span> Structure formed in a cave by the deposition of minerals from water

A speleothem is a geological formation by mineral deposits that accumulate over time in natural caves. Speleothems most commonly form in calcareous caves due to carbonate dissolution reactions. They can take a variety of forms, depending on their depositional history and environment. Their chemical composition, gradual growth, and preservation in caves make them useful paleoclimatic proxies.

<span class="mw-page-title-main">Hard water</span> Water that has a high mineral content

Hard water is water that has a high mineral content. Hard water is formed when water percolates through deposits of limestone, chalk or gypsum, which are largely made up of calcium and magnesium carbonates, bicarbonates and sulfates.

<span class="mw-page-title-main">Vaterite</span> Calcium carbonate mineral

Vaterite is a mineral, a polymorph of calcium carbonate (CaCO3). It was named after the German mineralogist Heinrich Vater. It is also known as mu-calcium carbonate (μ-CaCO3). Vaterite belongs to the hexagonal crystal system, whereas calcite is trigonal and aragonite is orthorhombic.

<span class="mw-page-title-main">Limescale</span> Hard, chalky deposit of calcium carbonate

Limescale is a hard, chalky deposit, consisting mainly of calcium carbonate (CaCO3). It often builds up inside kettles, boilers, and pipework, especially that for hot water. It is also often found as a similar deposit on the inner surfaces of old pipes and other surfaces where hard water has flowed. Limescale also forms as travertine or tufa in hard water springs.

<span class="mw-page-title-main">Alkalinity</span> Capacity of water to resist changes in pH that would make the water more acidic

Alkalinity (from Arabic: القلوية, romanized: al-qaly, lit. 'ashes of the saltwort') is the capacity of water to resist acidification. It should not be confused with basicity, which is an absolute measurement on the pH scale. Alkalinity is the strength of a buffer solution composed of weak acids and their conjugate bases. It is measured by titrating the solution with an acid such as HCl until its pH changes abruptly, or it reaches a known endpoint where that happens. Alkalinity is expressed in units of concentration, such as meq/L (milliequivalents per liter), μeq/kg (microequivalents per kilogram), or mg/L CaCO3 (milligrams per liter of calcium carbonate). Each of these measurements corresponds to an amount of acid added as a titrant.

Calcium bicarbonate, also called calcium hydrogencarbonate, has the chemical formula Ca(HCO3)2. The term does not refer to a known solid compound; it exists only in aqueous solution containing 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.

The carbonate compensation depth (CCD) is the depth, in the oceans, at which the rate of supply of calcium carbonates matches the rate of solvation. That is, solvation 'compensates' supply. Below the CCD solvation is faster, so that carbonate particles dissolve and the carbonate shells (tests) of animals are not preserved. Carbonate particles cannot accumulate in the sediments where the sea floor is below this depth.

<span class="mw-page-title-main">Calcium reactor</span>

A calcium reactor is an efficient method to supply calcium and trace elements to a reef aquarium. Reactors may be used in elaborate freshwater and brackish aquariums where freshwater clams and other invertebrates need a constant supply of calcium.

<span class="mw-page-title-main">Alkali soil</span> Soil type with pH > 8.5

Alkali, or Alkaline, soils are clay soils with high pH, a poor soil structure and a low infiltration capacity. Often they have a hard calcareous layer at 0.5 to 1 metre depth. Alkali soils owe their unfavorable physico-chemical properties mainly to the dominating presence of sodium carbonate, which causes the soil to swell and difficult to clarify/settle. They derive their name from the alkali metal group of elements, to which sodium belongs, and which can induce basicity. Sometimes these soils are also referred to as alkaline sodic soils. Alkaline soils are basic, but not all basic soils are alkaline.

<span class="mw-page-title-main">Concrete degradation</span> Damage to concrete affecting its mechanical strength and its durability

Concrete degradation may have many different causes. Concrete is mostly damaged by the corrosion of reinforcement bars due to the carbonatation of hardened cement paste or chloride attack under wet conditions. Chemical damage is caused by the formation of expansive products produced by chemical reactions, by aggressive chemical species present in groundwater and seawater, or by microorganisms Other damaging processes can also involve calcium leaching by water infiltration, physical phenomena initiating cracks formation and propagation, fire or radiant heat, aggregate expansion, sea water effects, leaching, and erosion by fast-flowing water.

<span class="mw-page-title-main">Shell growth in estuaries</span>

Shell growth in estuaries is an aspect of marine biology that has attracted a number of scientific research studies. Many groups of marine organisms produce calcified exoskeletons, commonly known as shells, hard calcium carbonate structures which the organisms rely on for various specialized structural and defensive purposes. The rate at which these shells form is greatly influenced by physical and chemical characteristics of the water in which these organisms live. Estuaries are dynamic habitats which expose their inhabitants to a wide array of rapidly changing physical conditions, exaggerating the differences in physical and chemical properties of the water.

Oilfield scale inhibition is the process of preventing the formation of scale from blocking or hindering fluid flow through pipelines, valves, and pumps used in oil production and processing. Scale inhibitors (SIs) are a class of specialty chemicals that are used to slow or prevent scaling in water systems. Oilfield scaling is the precipitation and accumulation of insoluble crystals (salts) from a mixture of incompatible aqueous phases in oil processing systems. Scale is a common term in the oil industry used to describe solid deposits that grow over time, blocking and hindering fluid flow through pipelines, valves, pumps etc. with significant reduction in production rates and equipment damages. Scaling represents a major challenge for flow assurance in the oil and gas industry. Examples of oilfield scales are calcium carbonate (limescale), iron sulfides, barium sulfate and strontium sulfate. Scale inhibition encompasses the processes or techniques employed to treat scaling problems.

<span class="mw-page-title-main">Calthemite</span> Secondary calcium carbonate deposit growing under man-made structures

Calthemite is a secondary deposit, derived from concrete, lime, mortar or other calcareous material outside the cave environment. Calthemites grow on or under man-made structures and mimic the shapes and forms of cave speleothems, such as stalactites, stalagmites, flowstone etc. Calthemite is derived from the Latin calx "lime" + Latin < Greek théma, "deposit" meaning ‘something laid down’, and the Latin –ita < Greek -itēs – used as a suffix indicating a mineral or rock. The term "speleothem", due to its definition can only be used to describe secondary deposits in caves and does not include secondary deposits outside the cave environment.

<span class="mw-page-title-main">Marine biogenic calcification</span> Shell formation mechanism

Marine biogenic calcification is the production of calcium carbonate by organisms in the global ocean.

<span class="mw-page-title-main">Calcium cycle</span>

The calcium cycle is a transfer of calcium between dissolved and solid phases. There is a continuous supply of calcium ions into waterways from rocks, organisms, and soils. Calcium ions are consumed and removed from aqueous environments as they react to form insoluble structures such as calcium carbonate and calcium silicate, which can deposit to form sediments or the exoskeletons of organisms. Calcium ions can also be utilized biologically, as calcium is essential to biological functions such as the production of bones and teeth or cellular function. The calcium cycle is a common thread between terrestrial, marine, geological, and biological processes. Calcium moves through these different media as it cycles throughout the Earth. The marine calcium cycle is affected by changing atmospheric carbon dioxide due to ocean acidification.

<span class="mw-page-title-main">Total inorganic carbon</span> Sum of the inorganic carbon species

Total inorganic carbon is the sum of the inorganic carbon species.

<span class="mw-page-title-main">Particulate inorganic carbon</span>

Particulate inorganic carbon (PIC) can be contrasted with dissolved inorganic carbon (DIC), the other form of inorganic carbon found in the ocean. These distinctions are important in chemical oceanography. Particulate inorganic carbon is sometimes called suspended inorganic carbon. In operational terms, it is defined as the inorganic carbon in particulate form that is too large to pass through the filter used to separate dissolved inorganic carbon.

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