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 (usuallymarble).
Alternatively, calcium carbonate is prepared fromcalcium oxide. Water is added to givecalcium hydroxide thencarbon dioxide is passed through this solution toprecipitate the desired calcium carbonate, referred to in the industry as precipitated calcium carbonate (PCC). This process is calledcarbonatation:[9]
CaO + H2O → Ca(OH)2
Ca(OH)2 + CO2 → CaCO3 + H2O
In a laboratory, calcium carbonate can easily be crystallized fromcalcium chloride (CaCl2), by placing anaqueous solution ofCaCl2 in adesiccator alongsideammonium carbonate[NH4]2CO3.[10] In the desiccator, ammonium carbonate is exposed to air and decomposes intoammonia, carbon dioxide, andwater. The carbon dioxide then diffuses into the aqueous solution of calcium chloride, reacts with the calcium ions and the water, and forms calcium carbonate.
The thermodynamically stable form ofCaCO3 under normal conditions ishexagonal β-CaCO3 (the mineralcalcite). Other forms can be prepared, the denser (2.83 g/cm3)orthorhombic λ-CaCO3 (the mineralaragonite) and hexagonal μ-CaCO3, occurring as the mineralvaterite. 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, whereasstrontium carbonate (SrCO3) andbarium carbonate (BaCO3) adopt the aragonite structure, reflecting their largerionic radii.[citation needed]
Thecalcite crystal structure istrigonal, withspace group R3c (No. 167 in the International Tables for Crystallography[15]), andPearson symbol hR10.[16]Aragonite isorthorhombic, 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 exhibitshexagonal symmetry in space group P63/mmc, the minor structure is still unknown.[18]
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 followOstwald'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
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 ascadaverine andPoly(ethylene imine) have been shown to facilitate the formation of aragonite over calcite.[14] Solid‑state NMR analysis has revealed that poly‑aspartate-stabilized ACC contains water molecules that undergo millisecond-timescale flips, illustrating dynamic hydration as a key factor in delaying crystallization.[23]
Organisms, such asmolluscs andarthropods, have shown the ability to grow all three crystal polymorphs of calcium carbonate, mainly as protection (shells) and muscle attachments.[24] 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.[25][26][27]
Calcite is the most stablepolymorph of calcium carbonate. It is transparent to opaque. A transparent variety calledIceland spar (shown here) was used to createpolarized light in the 19th century.[28]
Eggshells,snail shells and mostseashells are predominantly calcium carbonate and can be used as industrial sources of that chemical.[29]Oyster shells have enjoyed recent recognition as a source of dietary calcium, but are also a practical industrial source.[30][31] Darkgreen vegetables such asbroccoli andkale contain dietarily significant amounts of calcium carbonate, but they are not practical as an industrial source.[32]
Annelids in the familyLumbricidae, earthworms, possess a regionalization of the digestive track calledcalciferous glands, Kalkdrüsen, or glandes de Morren, that processes calcium andCO2 into calcium carbonate, which is later excreted into the dirt.[33] The function of these glands is unknown but is believed to serve as aCO2 regulation mechanism within the animals' tissues.[34] This process is ecologically significant, stabilizing thepH of acid soils.[35]
Beyond Earth, strong evidence suggests the presence of calcium carbonate onMars. Signs of calcium carbonate have been detected at more than one location (notably atGusev andHuygens craters). This provides some evidence for the past presence of liquid water.[36][37]
In warm, clear tropical waterscorals are more abundant than towards the poles where the waters are cold. Calcium carbonate contributors, includingplankton (such ascoccoliths and plankticforaminifera),coralline algae,sponges,brachiopods,echinoderms,bryozoa andmollusks, 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. Thecalcification processes are changed byocean acidification.
Thecarbonate 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.[38] 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.[39]
Trilobite populations were once thought to have composed the majority of aquatic life during theCambrian, due to the fact that their calcium carbonate-rich shells were more easily preserved than those of other species,[41] which had purelychitinous shells.
The main use of calcium carbonate is in the construction industry, either as a building material, or limestoneaggregate for road building, as an ingredient ofcement, or as the starting material for the preparation ofbuilders' lime by burning in akiln. However, because of weathering mainly caused byacid rain,[42] 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 ofiron fromiron ore in ablast furnace. The carbonate iscalcinedin situ to givecalcium oxide, which forms aslag with various impurities present, and separates from the purified iron.[43]
In theoil industry, calcium carbonate is added todrilling 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 toswimming pools, as apH corrector for maintainingalkalinity and offsetting the acidic properties of thedisinfectant agent.[44]
It is also used as a raw material in therefining of sugar fromsugar beet; it is calcined in a kiln withanthracite to produce calcium oxide and carbon dioxide. This burnt lime is then slaked in fresh water to produce a calcium hydroxidesuspension for the precipitation of impurities in raw juice duringcarbonatation.[45]
Calcium carbonate in the form ofchalk has traditionally been a major component ofblackboard chalk. However, modern manufactured chalk is mostlygypsum, hydratedcalcium sulfateCaSO4·2H2O. Calcium carbonate is a main source for growingbiorock. Precipitated calcium carbonate (PCC), pre-dispersed inslurry form, is a common filler material forlatex gloves with the aim of achieving maximum saving in material and production costs.[46]
Fine ground calcium carbonate (GCC) is an essential ingredient in the microporous film used indiapers 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 inpaper because they are cheaper thanwood fiber. Printing and writing paper can contain 10–20% calcium carbonate. In North America, calcium carbonate has begun to replacekaolin in the production ofglossy paper. Europe has been practicing this as alkalinepapermaking 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 anextender inpaints,[47] in particularmatteemulsion paint where typically 30% by weight of the paint is either chalk or marble. It is also a popular filler in plastics.[47] Some typical examples include around 15–20% loading of chalk inunplasticized polyvinyl chloride (uPVC)drainpipes, 5–15% loading ofstearate-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.[48] Here the percentage is often 20–40%. It also routinely used as a filler inthermosetting resins (sheet and bulk molding compounds)[48] and has also been mixed withABS, and other ingredients, to form some types of compression molded "clay"poker chips.[49] Precipitated calcium carbonate, made by droppingcalcium oxide into water, is used by itself or with additives as a white paint, known aswhitewashing.[50][51]
Calcium carbonate is added to a wide range of trade anddo it yourself adhesives, sealants, and decorating fillers.[47] 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 settingstained glass windows, and as a resist to prevent glass from sticking to kiln shelves when firing glazes and paints at high temperature.[52][53][54][55]
Inceramic glaze applications, calcium carbonate is known aswhiting,[47] 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 aflux material in the glaze. Ground calcium carbonate is anabrasive (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 theMohs scale, and will therefore not scratchglass and most otherceramics,enamel,bronze,iron, andsteel, and have a moderate effect on softer metals likealuminium andcopper. A paste made from calcium carbonate anddeionized water can be used to cleantarnish onsilver.[56]
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.[59]
Calcium carbonate is used therapeutically as phosphate binder in patients on maintenancehaemodialysis. 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, particularlysevelamer.
Excess calcium from supplements, fortified food, and high-calcium diets can causemilk-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 inkidney failure,alkalosis, andhypercalcaemia, 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 forpeptic 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,[60][61] and is exacerbated bydehydration. 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.[62]
Several calcium supplement formulations have been documented to contain the chemical elementlead,[70] posing apublic health concern.[71] Lead is commonly found in natural sources of calcium.[70]
Agricultural lime, powdered chalk or limestone, is used as a cheap method of neutralisingacidic soil, making it suitable for planting, also used in aquaculture industry for pH regulation of pond soil before initiating culture.[72] There is interest in understanding whether or not it can affect pesticide adsorption and desorption in calcareous soil.[73]
In 1989, a researcher, Ken Simmons, introducedCaCO3 into the Whetstone Brook inMassachusetts.[74] 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 thatCaCO3 can be added to neutralize the effects of acid rain inriver ecosystems. Currently calcium carbonate is used to neutralize acidic conditions in both soil and water.[75][76] Since the 1970s, suchliming 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 influe-gas desulfurization applications eliminating harmfulSO2 andNO2 emissions from coal and other fossil fuels burnt in large fossil fuel power stations.[75]
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 oflimestone usingcharcoal fires to producequicklime 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 andcarbon dioxide at any temperature. At each temperature there is apartial pressure of carbon dioxide that is in equilibrium with calcium carbonate. At room temperature the equilibrium overwhelmingly favors calcium carbonate, because the equilibriumCO2 pressure is only a tiny fraction of the partialCO2 pressure in air, which is about 0.035 kPa.
At temperatures above 550 °C the equilibriumCO2 pressure begins to exceed theCO2 pressure in air. So above 550 °C, calcium carbonate begins to outgasCO2 into air. However, in a charcoal fired kiln, the concentration ofCO2 will be much higher than it is in air. Indeed, if all theoxygen in the kiln is consumed in the fire, then the partial pressure ofCO2 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 ofCO2 from calcium carbonate to happen at an economically useful rate, the equilibrium pressure must significantly exceed the ambient pressure ofCO2. And for it to happen rapidly, the equilibrium pressure must exceed total atmospheric pressure of 101 kPa, which happens at 898 °C.
Equilibrium pressure ofCO2 overCaCO3 (P) versus temperature (T).[80]
Calcium carbonate is poorly soluble in pure water (47 mg/L at normal atmosphericCO2 partial pressure as shown below).
The equilibrium of its solution is given by the equation (with dissolved calcium carbonate on the right):
CaCO3 ⇌ Ca2+ + CO2−3
Ksp =3.7×10−9 to8.7×10−9 at 25 °C
where thesolubility product for[Ca2+][CO2−3] is given as anywhere fromKsp =3.7×10−9 toKsp =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 dissolvedCa2+ per liter of solution) with the molar concentration of dissolvedCO2−3 cannot exceed the value ofKsp. This seemingly simple solubility equation, however, must be taken along with the more complicated equilibrium ofcarbon dioxide withwater (seecarbonic acid). Some of theCO2−3 combines withH+ in the solution according to
HCO−3 ⇌ H+ + CO2−3
Ka2 =5.61×10−11 at 25 °C
HCO−3 is known as thebicarbonate ion.Calcium bicarbonate is many times more soluble in water than calcium carbonate—indeed it existsonly in solution.
Some of theHCO−3 combines withH+ in solution according to
H2CO3 ⇌ H+ + HCO−3
Ka1 =2.5×10−4 at 25 °C
Some of theH2CO3 breaks up into water and dissolved carbon dioxide according to
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), andPCO2 is theCO2 partial pressure.
For ambient air,PCO2 is around3.5×10−4atm (or equivalently 35 Pa). The last equation above fixes the concentration of dissolvedCO2 as a function ofPCO2, independent of the concentration of dissolvedCaCO3. At atmospheric partial pressure ofCO2, dissolvedCO2 concentration is1.2×10−5 moles per liter. The equation before that fixes the concentration ofH2CO3 as a function ofCO2 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 ofCO2partial pressure at 25 °C (Ksp = 4.47×10−9)
(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 ofCO2 (Ksp =4.47×10−9 has been taken for the calculation).
At atmospheric levels of ambientCO2 the table indicates that the solution will be slightly alkaline with a maximumCaCO3 solubility of 47 mg/L.
As ambientCO2 partial pressure is reduced below atmospheric levels, the solution becomes more and more alkaline. At extremely lowPCO2, dissolvedCO2, bicarbonate ion, and carbonate ion largely evaporate from the solution, leaving a highly alkaline solution ofcalcium hydroxide, which is more soluble thanCaCO3. ForPCO2 = 10−12 atm, the[Ca2+][OH−]2 product is still below the solubility product ofCa(OH)2 (8×10−6). For still lowerCO2 pressure,Ca(OH)2 precipitation will occur beforeCaCO3 precipitation.
As ambientCO2 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 ofCa2+.
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 ofCO2 much higher than atmospheric. As such, water percolates through calcium carbonate rock, theCaCO3 dissolves according to one of the trends above. When that same water then emerges from the tap, in time, it comes into equilibrium withCO2 levels in the air by outgassing its excessCO2. The calcium carbonate becomes less soluble as a result, and the excess precipitates as lime scale. This same process is responsible for the formation ofstalactites andstalagmites in limestone caves.
Two hydrated phases of calcium carbonate,monohydrocalciteCaCO3·H2O andikaiteCaCO3·6H2O, mayprecipitate 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 ofsodium 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 atmosphericCO2. Progress towards equilibrium through outgassing ofCO2 is slowed by
periodic replenishment of bicarbonate to maintain buffer capacity (often estimated through measurement oftotal 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 ofHCO−3 (which constitutes more than 90% ofBjerrum plot species from pH 7 to pH 8 at 25 °C in fresh water).[83] Addition ofHCO−3 will increaseCO2−3 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, whenHCO−3 concentration is known, the maximum concentration ofCa2+ ions before scaling throughCaCO3 precipitation can be predicted from the formula:
[Ca2+]max =Ksp/Ka2 ×[H+]/[HCO−3]
The solubility product forCaCO3 (Ksp) and the dissociation constants for the dissolved inorganic carbon species (includingKa2) are all substantially affected by temperature andsalinity,[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 withMg2+,[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]
Solutions ofstrong (HCl), moderately strong (sulfamic) orweak (acetic,citric,sorbic,lactic,phosphoric) acids are commercially available. They are commonly used asdescaling agents to removelimescale deposits. The maximum amount ofCaCO3 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 (withCaCO3 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
DissolvedCaCO3 (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 noCa2+ (not taking into account possibleCO2 dissolution) and the final state is the solution with saturatedCa2+. For strong acid concentrations, all species have a negligible concentration in the final state with respect toCa2+ andA− 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 ofCaCO3 in pure water.
In the case of a weak monoacid (here we take acetic acid withpKa = 4.76) with decreasing total acid concentration [A] = [A−] + [AH], we obtain:
For the same total acid concentration, the initial pH of the weak acid is less acid than that of the strong acid; however, the maximum amount ofCaCO3 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 manyH+ ions as the strong acid to "dissolve" the calcium carbonate.
The calculation in the case ofphosphoric 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−3], [CO2−3], [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
DissolvedCaCO3 (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] + [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 (seephosphoric acid).
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