| IUPAC name |
3D model (JSmol)
|E number||E170 (colours)|
CompTox Dashboard (EPA)
|Molar mass||100.0869 g/mol|
|Appearance||Fine white powder; chalky taste|
|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)
|0.013 g/L (25 °C)|
Solubility product (Ksp)
|Solubility in dilute acids||soluble|
Refractive index (nD)
Std enthalpy of
|A02AC01 ( WHO ) A12AA04 ( WHO )|
|Safety data sheet||ICSC 1193|
|NFPA 704 (fire diamond)|
|Lethal dose or concentration (LD, LC):|
LD50 (median dose)
|6450 mg/kg (oral, rat)|
|NIOSH (US health exposure limits):|
|TWA 15 mg/m3 (total) TWA 5 mg/m3 (resp)|
| Magnesium carbonate |
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
Calcium carbonate is a chemical compound with the formula Ca CO3. 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 and cause poor digestion.
Calcium carbonate shares the typical properties of other carbonates. Notably it
Calcium carbonate will react with water that is saturated with carbon dioxide to form the soluble calcium bicarbonate.
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):
The thermodynamically stable form of CaCO3 under normal conditions is hexagonal β-CaCO3 (the mineral calcite). 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. The vaterite structure is not fully understood. 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.Other forms can be prepared, the denser (2.83
Calcite, aragonite and vaterite are pure calcium carbonate minerals. Industrially important source rocks which are predominantly calcium carbonate include limestone, chalk, marble and travertine.
Eggshells, snail shells and most seashells are predominantly calcium carbonate and can be used as industrial sources of that chemical.Oyster shells have enjoyed recent recognition as a source of dietary calcium, but are also a practical industrial source. Dark green vegetables such as broccoli and kale contain dietarily significant amounts of calcium carbonate, however, they are not practical as an industrial source.
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.
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.
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.
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.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.
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,which had purely chitinous shells.
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,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.
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.
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.
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.
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. [ citation needed ]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, 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. Here the percentage is often 20–40%. It also routinely used as a filler in thermosetting resins (sheet and bulk molding compounds) and has also been mixed with ABS, and other ingredients, to form some types of compression molded "clay" poker chips. Precipitated calcium carbonate, made by dropping calcium oxide into water, is used by itself or with additives as a white paint, known as whitewashing.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. 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
Calcium carbonate is added to a wide range of trade and do it yourself adhesives, sealants, and decorating fillers. [ citation needed ]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,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.
Calcium carbonate is widely used medicinally as an inexpensive dietary calcium supplement for gastric antacid(such as Tums). 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.
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.
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, 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.
As a food additive it is designated E170,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, USA and Australia and New Zealand. 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. Calcium carbonate is also used as a firming agent in many canned and bottled vegetable products.
Agricultural lime, powdered chalk or limestone, is used as a cheap method for neutralising acidic soil, making it suitable for planting.
Calcium carbonate is a key ingredient in many household cleaning powders like Comet and is used as a scrubbing agent.
In 1989, a researcher, Ken Simmons, introduced CaCO3 into the Whetstone Brook in Massachusetts.His hope was that the calcium carbonate would counter the acid in the stream from acid rain and save the trout that had ceased to spawn. Although his experiment was a success, it did increase the 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. 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.
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.
Calcination of limestone using charcoal fires to produce quicklime has been practiced since antiquity by cultures all over the world. The temperature at which limestone yields calcium oxide is usually given as 825 °C, but stating an absolute threshold is misleading. Calcium carbonate exists in equilibrium with calcium oxide and carbon dioxide at any temperature. At each temperature there is a partial pressure of carbon dioxide that is in equilibrium with calcium carbonate. At room temperature the equilibrium overwhelmingly favors calcium carbonate, because the equilibrium CO2 pressure is only a tiny fraction of the partial CO2 pressure in air, which is about 0.035 kPa.
At temperatures above 550 °C the equilibrium CO2 pressure begins to exceed the CO2 pressure in air. So above 550 °C, calcium carbonate begins to outgas CO2 into air. However, in a charcoal fired kiln, the concentration of CO2 will be much higher than it is in air. Indeed, if all the oxygen in the kiln is consumed in the fire, then the partial pressure of CO2 in the kiln can be as high as 20 kPa.
The table shows that this 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.
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−
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. 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
3⇌ H+ + CO2−
|Ka2 = 5.61×10−11 at 25 °C|
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−|
|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 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
|PCO2 (atm)||pH||[Ca2+] (mol/L)|
|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,
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).
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.
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
In this situation, the dissociation constants for the much faster reactions
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). Addition of HCO−
3 will increase CO2−
3 concentration at any pH. Rearranging the equations given above, we can see that [Ca2+] = Ksp/, and [CO2−
3] = Ka2 [HCO−
3]/. Therefore, when HCO−
3 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,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)−
4 and other ions in the pool, as well as supersaturation effects. 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.
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.
(g/L of acid)
(g/L of acid)
(g/L of acid)
In inorganic chemistry, bicarbonate is an intermediate form in the deprotonation of carbonic acid. It is a polyatomic anion with the chemical formula HCO−
In chemistry, a carbonate is a salt of carbonic acid (H2CO3), characterized by the presence of the carbonate ion, a polyatomic ion with the formula of CO2−
3. The name may also refer to a carbonate ester, an organic compound containing the carbonate group C(=O)(O–)2.
Carbonic acid is a chemical compound with the chemical formula H2CO3 (equivalently: OC(OH)2). It is also a name sometimes given to solutions of carbon dioxide in water (carbonated water), because such solutions contain small amounts of H2CO3. In physiology, carbonic acid is described as volatile acid or respiratory acid because it is the only acid excreted as a gas by the lungs. It plays an important role in the bicarbonate buffer system to maintain acid–base homeostasis.
Sodium carbonate, Na2CO3, (also known as washing soda, soda ash and soda crystals) is the inorganic compound with the formula Na2CO3 and its various hydrates. All forms are white, water-soluble salts. All forms have a strongly alkaline taste and give moderately alkaline solutions in water. Historically it was extracted from the ashes of plants growing in sodium-rich soils. Because the ashes of these sodium-rich plants were noticeably different from ashes of wood (once used to produce potash), sodium carbonate became known as "soda ash." It is produced in large quantities from sodium chloride and limestone by the Solvay process.
Limewater is the common name for a dilute aqueous solution of calcium hydroxide. Calcium hydroxide, Ca(OH)2, is sparsely soluble at room temperature in water (1.5 g/L at 25 °C). "Pure" (i.e. less than or fully saturated) limewater is clear and colorless, with a slight earthy smell and an astringent/bitter taste. It is basic in nature with a pH of 12.4.
Hard water is water that has 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.
In chemistry, neutralization or neutralisation is a chemical reaction in which acid and a base react quantitatively with each other. In a reaction in water, neutralization results in there being no excess of hydrogen or hydroxide ions present in the solution. The pH of the neutralized solution depends on the acid strength of the reactants.
The Solvay process or ammonia-soda process is the major industrial process for the production of sodium carbonate (soda ash, Na2CO3). The ammonia-soda process was developed into its modern form by Ernest Solvay during the 1860s. The ingredients for this are readily available and inexpensive: salt brine (from inland sources or from the sea) and limestone (from quarries). The worldwide production of soda ash in 2005 has been estimated at 42 million metric tons, which is more than six kilograms (13 lb) per year for each person on Earth. Solvay-based chemical plants now produce roughly three-quarters of this supply, with the remaining being mined from natural deposits. This method superseded the Leblanc process.
Gravimetric analysis describes a set of methods used in analytical chemistry for the quantitative determination of an analyte based on its mass. The principle of this type of analysis is that once an ion's mass has been determined as a unique compound, that known measurement can then be used to determine the same analyte's mass in a mixture, as long as the relative quantities of the other constituents are known.
Limescale is a hard chalky deposit, consisting mainly of calcium carbonate (CaCO3), that often builds up inside kettles, hot water 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 evaporated.
Alkalinity is the capacity of water to resist changes in pH that would make the water more acidic. Alkalinity is the strength of a buffer solution composed of weak acids and their conjugate bases. It is measured by titrating the solution with a monoprotic acid such as HCl until its pH changes abruptly, or it reaches a known endpoint where that happens. Alkalinity is expressed in units of meq/L, which corresponds to the amount of monoprotic acid added as a titrant in millimoles per liter.
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.
Carbonate compensation depth (CCD) is the depth in the oceans below which the rate of supply of calcite lags behind the rate of solvation, such that no calcite is preserved. Aragonite compensation depth describes the same behaviour in reference to aragonitic carbonates. Aragonite is more soluble than calcite, so the aragonite compensation depth is generally shallower than the calcite compensation depth.
In marine and reef aquariums, a calcium reactor creates a balance of alkalinity. An acidic solution is produced by injecting carbon dioxide into a chamber with salt water and calcium rich media. The carbon dioxide lowers the pH by producing a solution high in carbonic acid, and dissolves calcium. The effluent is returned to the reef aquarium where the calcium is consumed by organisms, primarily corals when building skeletons. A calcium reactor is an efficient method to supply calcium 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.
Speleogenesis is the origin and development of caves, the primary process that determines essential features of the hydrogeology of karst and guides its evolution. It often deals with the development of caves through limestone, caused by the presence of water with carbon dioxide dissolved within it, producing carbonic acid which permits the dissociation of the calcium carbonate in the limestone.
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.
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.
The residual sodium carbonate (RSC) index of irrigation water or soil water is used to indicate the alkalinity hazard for soil. The RSC index is used to find the suitability of the water for irrigation in clay soils which have a high cation exchange capacity. When dissolved sodium in comparison with dissolved calcium and magnesium is high in water, clay soil swells or undergoes dispersion which drastically reduces its infiltration capacity.
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.
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.