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Calcium is a with the Ca and 20. As an , calcium is a reactive metal that forms a dark oxide-nitride layer when exposed to air. Its physical and chemical properties are most similar to its heavier homologues and . It is the fifth most abundant element in Earth's crust, and the third most abundant metal, after and . The most common calcium compound on Earth is , found in and the fossilised remnants of early sea life; , , , and are also sources of calcium. The name derives from ''calx'' "", which was obtained from heating limestone. Some calcium compounds were known to the ancients, though their chemistry was unknown until the seventeenth century. Pure calcium was isolated in 1808 via of its oxide by , who named the element. Calcium compounds are widely used in many industries: in foods and pharmaceuticals for , in the paper industry as bleaches, as components in cement and electrical insulators, and in the manufacture of soaps. On the other hand, the metal in pure form has few applications due to its high reactivity; still, in small quantities it is often used as an alloying component in steelmaking, and sometimes, as a calcium–lead alloy, in making automotive batteries. Calcium is the most abundant metal and the fifth-most abundant element in the . As s, calcium ions play a vital role in the and processes of organisms and s: in pathways where they act as a ; in release from ; in contraction of all cell types; as cofactors in many s; and in . Calcium ions outside cells are important for maintaining the across excitable s, synthesis, and bone formation.


Characteristics


Classification

Calcium is a very ductile silvery metal (sometimes described as pale yellow) whose properties are very similar to the heavier elements in its group, , , and . A calcium atom has twenty electrons, arranged in the s2. Like the other elements placed in group 2 of the periodic table, calcium has two s in the outermost s-orbital, which are very easily lost in chemical reactions to form a dipositive ion with the stable electron configuration of a , in this case . Hence, calcium is almost always in its compounds, which are usually . Hypothetical univalent salts of calcium would be stable with respect to their elements, but not to to the divalent salts and calcium metal, because the of MX2 is much higher than those of the hypothetical MX. This occurs because of the much greater afforded by the more highly charged Ca2+ cation compared to the hypothetical Ca+ cation.Greenwood and Earnshaw, pp. 112–3 Calcium, strontium, barium, and radium are always considered to be s; the lighter and , also in group 2 of the periodic table, are often included as well. Nevertheless, beryllium and magnesium differ significantly from the other members of the group in their physical and chemical behaviour: they behave more like and respectively and have some of the weaker metallic character of the s, which is why the traditional definition of the term "alkaline earth metal" excludes them. This classification is mostly obsolete in English-language sources, but is still used in other countries such as Japan. As a result, comparisons with strontium and barium are more germane to calcium chemistry than comparisons with magnesium.


Physical

Calcium metal melts at 842 °C and boils at 1494 °C; these values are higher than those for magnesium and strontium, the neighbouring group 2 metals. It crystallises in the arrangement like strontium; above 450 °C, it changes to an arrangement like magnesium. Its density of 1.55 g/cm3 is the lowest in its group. Calcium is harder than but can be cut with a knife with effort. While calcium is a poorer conductor of electricity than or by volume, it is a better conductor by mass than both due to its very low density. While calcium is infeasible as a conductor for most terrestrial applications as it reacts quickly with atmospheric oxygen, its use as such in space has been considered.Hluchan and Pomerantz, p. 484


Chemical

The chemistry of calcium is that of a typical heavy alkaline earth metal. For example, calcium spontaneously reacts with water more quickly than magnesium and less quickly than strontium to produce and hydrogen gas. It also reacts with the and in the air to form a mixture of and .C. R. Hammond ''The elements'' (p. 4–35) in When finely divided, it spontaneously burns in air to produce the nitride. In bulk, calcium is less reactive: it quickly forms a hydration coating in moist air, but below 30% it may be stored indefinitely at room temperature.Hluchan and Pomerantz, p. 483 Besides the simple oxide CaO, the CaO2 can be made by direct oxidation of calcium metal under a high pressure of oxygen, and there is some evidence for a yellow Ca(O2)2. Calcium hydroxide, Ca(OH)2, is a strong base, though it is not as strong as the hydroxides of strontium, barium or the alkali metals. All four dihalides of calcium are known. (CaCO3) and (CaSO4) are particularly abundant minerals.Greenwood and Earnshaw, pp. 122–5 Like strontium and barium, as well as the alkali metals and the divalent s and , calcium metal dissolves directly in liquid to give a dark blue solution. Due to the large size of the Ca2+ ion, high coordination numbers are common, up to 24 in some s such as CaZn13. Calcium is readily complexed by oxygen s such as and s, which are useful in and removing calcium ions from . In the absence of , smaller group 2 cations tend to form stronger complexes, but when large s are involved the trend is reversed. Although calcium is in the same group as magnesium and s are very commonly used throughout chemistry, organocalcium compounds are not similarly widespread because they are more difficult to make and more reactive, although they have recently been investigated as possible s. Organocalcium compounds tend to be more similar to organoytterbium compounds due to the similar of Yb2+ (102 pm) and Ca2+ (100 pm). Most of these compounds can only be prepared at low temperatures; bulky ligands tend to favor stability. For example, calcium di, Ca(C5H5)2, must be made by directly reacting calcium metal with or itself; replacing the C5H5 ligand with the bulkier C5(CH3)5 ligand on the other hand increases the compound's solubility, volatility, and kinetic stability.


Isotopes

Natural calcium is a mixture of five stable s (40Ca, 42Ca, 43Ca, 44Ca, and 46Ca) and one isotope with a half-life so long that it can be considered stable for all practical purposes (, with a half-life of about 4.3 × 1019 years). Calcium is the first (lightest) element to have six naturally occurring isotopes. By far the most common isotope of calcium in nature is 40Ca, which makes up 96.941% of all natural calcium. It is produced in the from fusion of s and is the heaviest stable nuclide with equal proton and neutron numbers; its occurrence is also supplemented slowly by the decay of . Adding another alpha particle leads to unstable 44Ti, which quickly decays via two successive s to stable 44Ca; this makes up 2.806% of all natural calcium and is the second-most common isotope. The other four natural isotopes, 42Ca, 43Ca, 46Ca, and 48Ca, are significantly rarer, each comprising less than 1% of all natural calcium. The four lighter isotopes are mainly products of the and silicon-burning processes, leaving the two heavier ones to be produced via processes. 46Ca is mostly produced in a "hot" , as its formation requires a rather high neutron flux to allow short-lived 45Ca to capture a neutron. 48Ca is produced by electron capture in the in e, where high neutron excess and low enough entropy ensures its survival. 46Ca and 48Ca are the first "classically stable" nuclides with a six-neutron or eight-neutron excess respectively. Although extremely neutron-rich for such a light element, 48Ca is very stable because it is a , having 20 protons and 28 neutrons arranged in closed shells. Its to 48 is very hindered because of the gross mismatch of : 48Ca has zero nuclear spin, being , while 48Sc has spin 6+, so the decay is by the conservation of . While two excited states of 48Sc are available for decay as well, they are also forbidden due to their high spins. As a result, when 48Ca does decay, it does so by to 48 instead, being the lightest nuclide known to undergo double beta decay. The heavy isotope 46Ca can also theoretically undergo double beta decay to 46Ti as well, but this has never been observed; the lightest and most common isotope 40Ca is also doubly magic and could undergo to 40, but this has likewise never been observed. Calcium is the only element to have two primordial doubly magic isotopes. The experimental lower limits for the half-lives of 40Ca and 46Ca are 5.9 × 1021 years and 2.8 × 1015 years respectively. Apart from the practically stable 48Ca, the longest lived of calcium is 41Ca. It decays by electron capture to stable 41 with a half-life of about a hundred thousand years. Its existence in the early Solar System as an has been inferred from excesses of 41K: traces of 41Ca also still exist today, as it is a , continuously reformed through of natural 40Ca. Many other calcium radioisotopes are known, ranging from 35Ca to 60Ca. They are all much shorter-lived than 41Ca, the most stable among them being 45Ca (half-life 163 days) and 47Ca (half-life 4.54 days). The isotopes lighter than 42Ca usually undergo to isotopes of potassium, and those heavier than 44Ca usually undergo to isotopes of , although near the s, and begin to be significant decay modes as well. Like other elements, a variety of processes alter the relative abundance of calcium isotopes. The best studied of these processes is the mass-dependent of calcium isotopes that accompanies the precipitation of calcium minerals such as , and from solution. Lighter isotopes are preferentially incorporated into these minerals, leaving the surrounding solution enriched in heavier isotopes at a magnitude of roughly 0.025% per atomic mass unit (amu) at room temperature. Mass-dependent differences in calcium isotope composition are conventionally expressed by the ratio of two isotopes (usually 44Ca/40Ca) in a sample compared to the same ratio in a standard reference material. 44Ca/40Ca varies by about 1% among common earth materials.


History

Calcium compounds were known for millennia, although their chemical makeup was not understood until the 17th century.Greenwood and Earnshaw, p. 108 Lime as a and as was used as far back as around 7000 BC. The first dated dates back to 2500 BC and was found in , . At about the same time, dehydrated (CaSO4·2H2O) was being used in the ; this material would later be used for the plaster in the tomb of . The s instead used lime mortars made by heating (CaCO3); the name "calcium" itself derives from the Latin word ''calx'' "lime". noted that the lime that resulted was lighter than the original limestone, attributing this to the boiling of the water; in 1755, proved that this was due to the loss of , which as a gas had not been recognised by the ancient Romans. In 1787, suspected that lime might be an oxide of a fundamental . In his table of the elements, Lavoisier listed five "salifiable earths" (i.e., ores that could be made to react with acids to produce salts (''salis'' = salt, in Latin): ''chaux'' (calcium oxide), ''magnésie'' (magnesia, magnesium oxide), ''baryte'' (barium sulfate), ''alumine'' (alumina, aluminium oxide), and ''silice'' (silica, silicon dioxide)). About these "elements", Lavoisier speculated: Calcium, along with its congeners magnesium, strontium, and barium, was first isolated by in 1808. Following the work of and on , Davy isolated calcium and magnesium by putting a mixture of the respective metal oxides with on a plate which was used as the anode, the cathode being a platinum wire partially submerged into mercury. Electrolysis then gave calcium–mercury and magnesium–mercury amalgams, and distilling off the mercury gave the metal. However, pure calcium cannot be prepared in bulk by this method and a workable commercial process for its production was not found until over a century later.


Occurrence and production

At 3%, calcium is the fifth , and the third most abundant metal behind and . It is also the fourth most abundant element in the . deposits pervade the Earth's surface as fossilized remains of past marine life; they occur in two forms, the (more common) and the (forming in more temperate seas). Minerals of the first type include , , , , and ; aragonite beds make up the , the , and the basins. s, s, and s are mostly made up of calcium carbonate. Among the other important minerals of calcium are (CaSO4·2H2O), (CaSO4), (CaF2), and ( a5(PO4)3F. The major producers of calcium are (about 10000 to 12000 s per year), (about 6000 to 8000 tonnes per year), and the (about 2000 to 4000 tonnes per year). and are also among the minor producers. In 2005, about 24000 tonnes of calcium were produced; about half of the world's extracted calcium is used by the United States, with about 80% of the output used each year. In Russia and China, Davy's method of electrolysis is still used, but is instead applied to molten . Since calcium is less reactive than strontium or barium, the oxide–nitride coating that results in air is stable and machining and other standard metallurgical techniques are suitable for calcium.Greenwood and Earnshaw, p. 110 In the United States and Canada, calcium is instead produced by reducing lime with aluminium at high temperatures.


Geochemical cycling

provides a link between , , and the . In the simplest terms, uplift of mountains exposes calcium-bearing rocks to chemical weathering and releases Ca2+ into surface water. These ions are transported to the ocean where they react with dissolved CO2 to form (), which in turn settles to the sea floor where it is incorporated into new rocks. Dissolved CO2, along with and ions, are termed "" (DIC). The actual reaction is more complicated and involves the bicarbonate ion (HCO) that forms when CO2 reacts with water at seawater : : + 2 → () + + At seawater pH, most of the CO2 is immediately converted back into . The reaction results in a net transport of one molecule of CO2 from the ocean/atmosphere into the . The result is that each Ca2+ ion released by chemical weathering ultimately removes one CO2 molecule from the surficial system (atmosphere, ocean, soils and living organisms), storing it in carbonate rocks where it is likely to stay for hundreds of millions of years. The weathering of calcium from rocks thus scrubs CO2 from the ocean and atmosphere, exerting a strong long-term effect on climate.


Uses

The largest use of metallic calcium is in , due to its strong for oxygen and . Its oxides and sulfides, once formed, give liquid lime and sulfide inclusions in steel which float out; on treatment, these inclusions disperse throughout the steel and became small and spherical, improving castability, cleanliness and general mechanical properties. Calcium is also used in maintenance-free , in which the use of 0.1% calcium– alloys instead of the usual –lead alloys leads to lower water loss and lower self-discharging. Due to the risk of expansion and cracking, is sometimes also incorporated into these alloys. These lead–calcium alloys are also used in casting, replacing lead–antimony alloys.Hluchan and Pomerantz, pp. 485–7 Calcium is also used to strengthen aluminium alloys used for bearings, for the control of graphitic in , and to remove impurities from lead. Calcium metal is found in some drain cleaners, where it functions to generate heat and that the fats and liquefies the proteins (for example, those in hair) that block drains. Besides metallurgy, the reactivity of calcium is exploited to remove from high-purity gas and as a for oxygen and nitrogen. It is also used as a reducing agent in the production of , , , and . It can also be used to store hydrogen gas, as it reacts with hydrogen to form solid , from which the hydrogen can easily be re-extracted. Calcium isotope fractionation during mineral formation has led to several applications of calcium isotopes. In particular, the 1997 observation by Skulan and DePaolo that calcium minerals are isotopically lighter than the solutions from which the minerals precipitate is the basis of analogous applications in medicine and in paleooceanography. In animals with skeletons mineralized with calcium, the calcium isotopic composition of soft tissues reflects the relative rate of formation and dissolution of skeletal mineral. In humans, changes in the calcium isotopic composition of urine have been shown to be related to changes in bone mineral balance. When the rate of bone formation exceeds the rate of bone resorption, the 44Ca/40Ca ratio in soft tissue rises and vice versa. Because of this relationship, calcium isotopic measurements of urine or blood may be useful in the early detection of metabolic bone diseases like . A similar system exists in seawater, where 44Ca/40Ca tends to rise when the rate of removal of Ca2+ by mineral precipitation exceeds the input of new calcium into the ocean. In 1997, Skulan and DePaolo presented the first evidence of change in seawater 44Ca/40Ca over geologic time, along with a theoretical explanation of these changes. More recent papers have confirmed this observation, demonstrating that seawater Ca2+ concentration is not constant, and that the ocean is never in a "steady state" with respect to calcium input and output. This has important climatological implications, as the marine calcium cycle is closely tied to the . Many calcium compounds are used in food, as pharmaceuticals, and in medicine, among others. For example, calcium and phosphorus are supplemented in foods through the addition of , , and . The last is also used as a polishing agent in and in s. is a white powder that is used as a suspending agent for pharmaceuticals. In baking, is used as a . is used as a bleach in papermaking and as a disinfectant, is used as a reinforcing agent in rubber, and is a component of and is used to make metallic soaps and synthetic resins. Calcium is on the .


Food sources

Foods rich in calcium include s, such as and , s, , products, , and s. Because of concerns for long-term adverse side effects, including calcification of arteries and s, both the U.S. (IOM) and the (EFSA) set (ULs) for combined dietary and supplemental calcium. From the IOM, people of ages 9–18 years are not to exceed 3 g/day combined intake; for ages 19–50, not to exceed 2.5 g/day; for ages 51 and older, not to exceed 2 g/day. EFSA set the UL for all adults at 2.5 g/day, but decided the information for children and adolescents was not sufficient to determine ULs.


Biological and pathological role


Function

Calcium is an needed in large quantities. The Ca2+ ion acts as an and is vital to the health of the muscular, circulatory, and digestive systems; is indispensable to the building of bone; and supports synthesis and function of blood cells. For example, it regulates the contraction of muscles, nerve conduction, and the clotting of blood. As a result, intra- and extracellular calcium levels are tightly regulated by the body. Calcium can play this role because the Ca2+ ion forms stable es with many organic compounds, especially s; it also forms compounds with a wide range of solubilities, enabling the formation of the . Sosa Torres, Martha; Kroneck, Peter M.H; "Introduction: From Rocks to Living Cells" pp 1-32 in "Metals, Microbes and Minerals: The Biogeochemical Side of Life" (2021) pp xiv + 341. Walter de Gruyter, Berlin. Editors Kroneck, Peter M.H. and Sosa Torres, Martha
DOI 10.1515/9783110589771-001
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Binding

Calcium ions may be complexed by proteins through binding the s of or residues; through interacting with , , or residues; or by being by γ-carboxylated amino acid residues. , a digestive enzyme, uses the first method; , a bone matrix protein, uses the third. Some other bone matrix proteins such as and use both the first and the second. Direct activation of enzymes by binding calcium is common; some other enzymes are activated by noncovalent association with direct calcium-binding enzymes. Calcium also binds to the layer of the , anchoring proteins associated with the cell surface.Hluchan and Pomerantz, pp. 489–94


Solubility

As an example of the wide range of solubility of calcium compounds, is very soluble in water, 85% of extracellular calcium is as with a solubility of 2.0  and the of bones in an organic matrix is at 100 μM.


Nutrition

Calcium is a common constituent of s, but the composition of calcium complexes in supplements may affect its bioavailability which varies by solubility of the salt involved: , , and are highly bioavailable, while the is less. Other calcium preparations include , , and . The intestine absorbs about one-third of calcium eaten as the , and plasma calcium level is then regulated by the s.


Hormonal regulation of bone formation and serum levels

and promote the formation of bone by allowing and enhancing the deposition of calcium ions there, allowing rapid bone turnover without affecting bone mass or mineral content. When plasma calcium levels fall, cell surface receptors are activated and the secretion of parathyroid hormone occurs; it then proceeds to stimulate the entry of calcium into the plasma pool by taking it from targeted kidney, gut, and bone cells, with the bone-forming action of parathyroid hormone being antagonised by , whose secretion increases with increasing plasma calcium levels.


Abnormal serum levels

Excess intake of calcium may cause . However, because calcium is absorbed rather inefficiently by the intestines, high serum calcium is more likely caused by excessive secretion of parathyroid hormone (PTH) or possibly by excessive intake of vitamin D, both of which facilitate calcium absorption. All these conditions result in excess calcium salts being deposited in the heart, blood vessels, or kidneys. Symptoms include anorexia, nausea, vomiting, memory loss, confusion, muscle weakness, increased urination, dehydration, and metabolic bone disease. Chronic hypercalcaemia typically leads to of soft tissue and its serious consequences: for example, calcification can cause loss of elasticity of s and disruption of laminar blood flow—and thence to and . Conversely, inadequate calcium or vitamin D intakes may result in , often caused also by inadequate secretion of parathyroid hormone or defective PTH receptors in cells. Symptoms include neuromuscular excitability, which potentially causes and disruption of conductivity in cardiac tissue.


Kidney stones


Bone disease

As calcium is required for bone development, many bone diseases can be traced to the organic matrix or the in molecular structure or organization of bone. is a reduction in mineral content of bone per unit volume, and can be treated by supplementation of calcium, vitamin D, and s. Inadequate amounts of calcium, vitamin D, or phosphates can lead to softening of bones, called .


Safety


Metallic calcium

Because calcium reacts exothermically with water and acids, calcium metal coming into contact with bodily moisture results in severe corrosive irritation. When swallowed, calcium metal has the same effect on the mouth, oesophagus, and stomach, and can be fatal.Rumack BH. POISINDEX. Information System Micromedex, Inc., Englewood, CO, 2010; CCIS Volume 143. Hall AH and Rumack BH (Eds) However, long-term exposure is not known to have distinct adverse effects.Hluchan and Pomerantz, pp. 487–9


See also


References


Bibliography

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