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{{Infobox calcium|engvar=en-GB}}
'''Calcium''' is a [[Chemistry:Chemical element|chemical element]]; it has [[Chemistry:Symbol|symbol]] '''Ca''' and [[Physics:Atomic number|atomic number]] 20. As an [[Chemistry:Alkaline earth metal|alkaline earth metal]], 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 [[Chemistry:Strontium|strontium]] and [[Chemistry:Barium|barium]]. It is the fifth most abundant element in Earth's crust, and the third most abundant metal, after [[Chemistry:Iron|iron]] and [[Chemistry:Aluminium|aluminium]]. The most common calcium compound on Earth is [[Chemistry:Calcium carbonate|calcium carbonate]], found in [[Earth:Limestone|limestone]] and the fossilised remnants of early sea life; [[Chemistry:Gypsum|gypsum]], [[Chemistry:Anhydrite|anhydrite]], [[Chemistry:Fluorite|fluorite]], and [[Chemistry:Apatite|apatite]] are also sources of calcium. The name derives from [[Social:Latin language|Latin]] ''calx'' "[[Chemistry:Lime (material)|lime]]", 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 [[Chemistry:Electrolysis|electrolysis]] of its oxide by [[Biography:Humphry Davy|Humphry Davy]], who named the element. Calcium compounds are widely used in many industries: in foods and pharmaceuticals for calcium supplementation, 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 [[Medicine:Human body#Composition|human body]].<ref name="lpi">{{cite web | url=https://lpi.oregonstate.edu/mic/minerals/calcium | title=Calcium | publisher=Linus Pauling Institute, Oregon State University, Corvallis, Oregon | date=1 September 2017 | access-date=31 August 2019}}</ref> As [[Physics:Electrolyte|electrolyte]]s, [[Biology:Calcium in biology|calcium ions]] (Ca<sup>2+</sup>) play a vital role in the [[Biology:Physiology|physiological]] and [[Biology:Biochemistry|biochemical]] processes of organisms and [[Biology:Cell|cell]]s: in [[Biology:Signal transduction|signal transduction]] pathways where they act as a second messenger; in [[Biology:Neurotransmitter|neurotransmitter]] release from neurons; in contraction of all [[Biology:Muscle cell|muscle cell]] types; as [[Biology:Cofactor (biochemistry)|cofactors]] in many [[Biology:Enzyme|enzyme]]s; and in fertilization.<ref name="lpi" /> Calcium ions outside cells are important for maintaining the [[Physics:Potential difference|potential difference]] across excitable [[Biology:Cell membrane|cell membrane]]s, [[Biology:Protein|protein]] synthesis, and bone formation.<ref name="lpi" /><ref name="ods">{{cite web |title=Calcium: Fact Sheet for Health Professionals |url=https://ods.od.nih.gov/factsheets/Calcium-HealthProfessional/ |publisher=Office of Dietary Supplements, US National Institutes of Health |access-date=31 August 2019 |date=9 July 2019}}</ref>

==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, [[Chemistry:Strontium|strontium]], [[Chemistry:Barium|barium]], and [[Chemistry:Radium|radium]]. A calcium atom has twenty electrons, arranged in the [[Physics:Electron configuration|electron configuration]] [Ar]4s<sup>2</sup>. Like the other elements placed in group 2 of the periodic table, calcium has two [[Physics:Valence electron|valence electron]]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 [[Chemistry:Noble gas|noble gas]], in this case [[Chemistry:Argon|argon]].<ref name="Greenwood112"/>

Hence, calcium is almost always divalent in its compounds, which are usually [[Physics:Ionic compound|ionic]]. Hypothetical univalent salts of calcium would be stable with respect to their elements, but not to [[Biology:Disproportionation|disproportionation]] to the divalent salts and calcium metal, because the enthalpy of formation of MX<sub>2</sub> is much higher than those of the hypothetical MX. This occurs because of the much greater [[Physics:Lattice energy|lattice energy]] afforded by the more highly charged Ca<sup>2+</sup> cation compared to the hypothetical Ca<sup>+</sup> cation.<ref name="Greenwood112">Greenwood and Earnshaw, pp. 112–13</ref>

Calcium, strontium, barium, and radium are always considered to be [[Chemistry:Alkaline earth metal|alkaline earth metal]]s; the lighter [[Chemistry:Beryllium|beryllium]] and [[Chemistry:Magnesium|magnesium]], 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 [[Chemistry:Aluminium|aluminium]] and [[Chemistry:Zinc|zinc]] respectively and have some of the weaker metallic character of the [[Chemistry:Post-transition metal|post-transition metal]]s, which is why the traditional definition of the term "alkaline earth metal" excludes them.<ref>{{cite book |last=Parish |first=R. V. |date=1977 |title=The Metallic Elements |location=London |publisher=Longman |page=[https://archive.org/details/metallicelements0000pari/page/34 34] |isbn=978-0-582-44278-8 |url=https://archive.org/details/metallicelements0000pari/page/34 }}</ref>

===Physical properties===
Calcium metal melts at 842&nbsp;°C and boils at 1494&nbsp;°C; these values are higher than those for magnesium and strontium, the neighbouring group 2 metals. It crystallises in the face-centered cubic arrangement like strontium; above 450&nbsp;°C, it changes to an [[Anisotropy|anisotropic]] hexagonal close-packed arrangement like magnesium. Its density of 1.55&nbsp;g/cm<sup>3</sup> is the lowest in its group.<ref name="Greenwood112" />

Calcium is harder than [[Chemistry:Lead|lead]] but can be cut with a knife with effort. While calcium is a poorer conductor of electricity than [[Chemistry:Copper|copper]] or [[Chemistry:Aluminium|aluminium]] by volume, it is a better conductor by mass than both due to its very low density.<ref name=Ullmann484/> 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.<ref name="Ullmann484">Hluchan and Pomerantz, p. 484</ref>

===Chemical properties===
[[File:Ca(aq)6 improved image.tif|thumb|left|Structure of the polymeric [Ca(H<sub>2</sub>O)<sub>6</sub>]<sup>2+</sup> center in hydrated calcium chloride, illustrating the high coordination number typical for calcium complexes.]]

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 [[Chemistry:Calcium hydroxide|calcium hydroxide]] and hydrogen gas. It also reacts with the [[Chemistry:Oxygen|oxygen]] and [[Chemistry:Nitrogen|nitrogen]] in the air to form a mixture of [[Chemistry:Calcium oxide|calcium oxide]] and [[Chemistry:Calcium nitride|calcium nitride]].<ref name="CRC">C. R. Hammond ''The elements'' (pp. 4–35) in {{RubberBible86th}}</ref> 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% [[Physics:Relative humidity|relative humidity]] it may be stored indefinitely at room temperature.<ref name="Ullmann483">Hluchan and Pomerantz, p. 483</ref>

Besides the simple oxide CaO, the [[Chemistry:Peroxide|peroxide]] [[Chemistry:Calcium peroxide|CaO<sub>2</sub>]] can be made by direct oxidation of calcium metal under a high pressure of oxygen, and there is some evidence for a yellow [[Chemistry:Superoxide|superoxide]] Ca(O<sub>2</sub>)<sub>2</sub>.<ref>Greenwood and Earnshaw, p. 119</ref> Calcium hydroxide, Ca(OH)<sub>2</sub>, is a strong base, though it is not as strong as the hydroxides of strontium, barium or the alkali metals.<ref>Greenwood and Earnshaw, p. 121</ref> All four dihalides of calcium are known.<ref>Greenwood and Earnshaw, p. 117</ref> [[Chemistry:Calcium carbonate|Calcium carbonate]] (CaCO<sub>3</sub>) and [[Chemistry:Calcium sulfate|calcium sulfate]] (CaSO<sub>4</sub>) are particularly abundant minerals.<ref name="Greenwood122">Greenwood and Earnshaw, pp. 122–15</ref> Like strontium and barium, as well as the alkali metals and the divalent [[Chemistry:Lanthanide|lanthanide]]s [[Chemistry:Europium|europium]] and [[Chemistry:Ytterbium|ytterbium]], calcium metal dissolves directly in liquid [[Chemistry:Ammonia|ammonia]] to give a dark blue solution.<ref name="Greenwood112" />

Due to the large size of the calcium ion (Ca<sup>2+</sup>), high coordination numbers are common, up to 24 in some intermetallic compounds such as CaZn<sub>13</sub>.<ref>Greenwood and Earnshaw, p. 115</ref> Calcium is readily complexed by oxygen chelates such as [[Chemistry:Ethylenediaminetetraacetic acid|EDTA]] and [[Chemistry:Polyphosphate|polyphosphate]]s, which are useful in analytic chemistry and removing calcium ions from [[Physics:Hard water|hard water]]. In the absence of steric hindrance, smaller group 2 cations tend to form stronger complexes, but when large polydentate [[Chemistry:Macrocycle|macrocycle]]s are involved the trend is reversed.<ref name="Greenwood122" />

Although calcium is in the same group as magnesium and organomagnesium compounds 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 [[Chemistry:Catalyst|catalyst]]s.<ref>{{cite journal|last1=Harder|first1=S.|last2=Feil|first2=F.|last3=Knoll|first3=K.|year=2001|title=Novel Calcium Half-Sandwich Complexes for the Living and Stereoselective Polymerization of Styrene|journal=Angew. Chem. Int. Ed.|volume=40|issue=22|pages=4261–64|doi=10.1002/1521-3773(20011119)40:22<4261::AID-ANIE4261>3.0.CO;2-J|pmid=29712082}}</ref><ref>{{cite journal|last1=Crimmin|first1=Mark R.|last2=Casely|first2=Ian J.|last3=Hill|first3=Michael S.|title=Calcium-Mediated Intramolecular Hydroamination Catalysis|journal=[[Organization:Journal of the American Chemical Society|Journal of the American Chemical Society]]|year=2005|volume=127|issue=7|pages=2042–43|doi=10.1021/ja043576n|pmid=15713071}}</ref><ref>{{cite journal|last1=Jenter|first1=Jelena|last2=Köppe|first2=Ralf|last3=Roesky|first3=Peter W.|title=2,5-Bis{''N''-(2,6-diisopropylphenyl)iminomethyl}pyrrolyl Complexes of the Heavy Alkaline Earth Metals: Synthesis, Structures, and Hydroamination Catalysis|journal=Organometallics|year=2011|volume=30|issue=6|pages=1404–13|doi=10.1021/om100937c}}</ref><ref>{{cite journal|last1=Arrowsmith|first1=Merle|last2=Crimmin|first2=Mark R.|last3=Barrett|first3=Anthony G. M.|last4=Hill|first4=Michael S.|last5=Kociok-Köhn|first5=Gabriele|last6=Procopiou|first6=Panayiotis A.|title=Cation Charge Density and Precatalyst Selection in Group 2-Catalyzed Aminoalkene Hydroamination|journal=Organometallics|year=2011|volume=30|issue=6|pages=1493–1506|doi=10.1021/om101063m}}</ref><ref>{{cite journal|last1=Penafiel|first1=J.|last2=Maron|first2=L.|last3=Harder|first3=S.|year=2014|title=Early Main Group Metal Catalysis: How Important is the Metal?|journal=Angew. Chem. Int. Ed.|volume=54|issue=1|pages=201–06|doi=10.1002/anie.201408814|pmid=25376952|url=https://pure.rug.nl/ws/files/83571601/Early_Main_Group_Metal_Catalysis_How_Important_is_the_Metal.pdf }}</ref> Organocalcium compounds tend to be more similar to organoytterbium compounds due to the similar [[Chemistry:Ionic radius|ionic radii]] of Yb<sup>2+</sup> (102&nbsp;pm) and Ca<sup>2+</sup> (100&nbsp;pm).<ref name="Greenwood and Earnshaw, pp. 136–37">Greenwood and Earnshaw, pp. 136–37</ref>

Most of these compounds can only be prepared at low temperatures; bulky ligands tend to favor stability. For example, calcium dicyclopentadienyl, Ca(C<sub>5</sub>H<sub>5</sub>)<sub>2</sub>, must be made by directly reacting calcium metal with mercurocene or [[Chemistry:Cyclopentadiene|cyclopentadiene]] itself; replacing the C<sub>5</sub>H<sub>5</sub> ligand with the bulkier C<sub>5</sub>(CH<sub>3</sub>)<sub>5</sub> ligand on the other hand increases the compound's solubility, volatility, and kinetic stability.<ref name="Greenwood and Earnshaw, pp. 136–37"/>

===Isotopes===
{{main|Physics:Isotopes of calcium}}

Natural calcium is a mixture of five stable [[Physics:Isotope|isotope]]s (<sup>40</sup>Ca, <sup>42</sup>Ca, <sup>43</sup>Ca, <sup>44</sup>Ca, and <sup>46</sup>Ca) and one isotope with a half-life so long that it can be considered stable for all practical purposes ([[Physics:Calcium-48|<sup>48</sup>Ca]], with a half-life of about 4.3&nbsp;×&nbsp;10<sup>19</sup>&nbsp;years). Calcium is the first (lightest) element to have six naturally occurring isotopes.<ref name="CRC" />

By far the most common isotope of calcium in nature is <sup>40</sup>Ca, which makes up 96.941% of all natural calcium. It is produced in the [[Astronomy:Silicon-burning process|silicon-burning process]] from fusion of [[Physics:Alpha particle|alpha particle]]s and is the heaviest stable nuclide with equal proton and neutron numbers; its occurrence is also supplemented slowly by the decay of [[Chemistry:Primordial nuclide|primordial]] [[Physics:Potassium-40|<sup>40</sup>K]].  Adding another alpha particle leads to unstable <sup>44</sup>Ti, which quickly decays via two successive [[Physics:Electron capture|electron capture]]s to stable <sup>44</sup>Ca; this makes up 2.806% of all natural calcium and is the second-most common isotope.<ref name="Cameron"/><ref name="Clayton"/>

The other four natural isotopes, <sup>42</sup>Ca, <sup>43</sup>Ca, <sup>46</sup>Ca, and <sup>48</sup>Ca, are significantly rarer, each comprising less than 1% of all natural calcium.  The four lighter isotopes are mainly products of the [[Astronomy:Oxygen-burning process|oxygen-burning]] and silicon-burning processes, leaving the two heavier ones to be produced via [[Physics:Neutron capture|neutron capture]] processes.  <sup>46</sup>Ca is mostly produced in a "hot" [[Physics:S-process|s-process]], as its formation requires a rather high neutron flux to allow short-lived <sup>45</sup>Ca to capture a neutron. <sup>48</sup>Ca is produced by electron capture in the [[Physics:R-process|r-process]] in [[Astronomy:Type Ia supernova|type Ia supernova]]e, where high neutron excess and low enough entropy ensures its survival.<ref name="Cameron">{{cite journal | last1 = Cameron |first1 = A. G. W. | year = 1973 | title = Abundance of the Elements in the Solar System | url = https://pubs.giss.nasa.gov/docs/1973/1973_Cameron_ca06310p.pdf | journal = Space Science Reviews | volume = 15 |issue = 1 | pages = 121–46 | doi = 10.1007/BF00172440 | bibcode = 1973SSRv...15..121C |s2cid = 120201972 }}</ref><ref name="Clayton">{{cite book |last=Clayton |first=Donald |date=2003 |title=Handbook of Isotopes in the Cosmos: Hydrogen to Gallium |publisher=Cambridge University Press |pages=184–98 |isbn=9780521530835}}</ref>

<sup>46</sup>Ca and <sup>48</sup>Ca are the first "classically stable" nuclides with a six-neutron or eight-neutron excess respectively. Although extremely neutron-rich for such a light element, <sup>48</sup>Ca is very stable because it is a [[Physics:Magic number|doubly magic nucleus]], having 20 protons and 28 neutrons arranged in closed shells. Its [[Physics:Beta decay|beta decay]] to <sup>48</sup>[[Chemistry:Scandium|Sc]] is very hindered because of the gross mismatch of [[Physics:Nuclear spin|nuclear spin]]: <sup>48</sup>Ca has zero nuclear spin, being [[Physics:Even and odd atomic nuclei|even–even]], while <sup>48</sup>Sc has spin 6+, so the decay is [[Physics:Forbidden mechanism|forbidden]] by the conservation of [[Physics:Angular momentum|angular momentum]]. While two excited states of <sup>48</sup>Sc are available for decay as well, they are also forbidden due to their high spins. As a result, when <sup>48</sup>Ca does decay, it does so by [[Physics:Double beta decay|double beta decay]] to <sup>48</sup>[[Chemistry:Titanium|Ti]] instead, being the lightest nuclide known to undergo double beta decay.{{NUBASE2016|ref}}<ref>{{Cite journal
 |last1=Arnold |first1=R.
 |display-authors=0
 |year=2016
 |collaboration=NEMO-3 Collaboration
 |title= Measurement of the double-beta decay half-life and search for the neutrinoless double-beta decay of <sup>48</sup>Ca with the NEMO-3 detector
 |journal=[[Physics:Physical Review D|Physical Review D]]
 |volume=93 |issue=11
 |page=112008
 |doi= 10.1103/PhysRevD.93.112008
|arxiv=1604.01710|bibcode=2016PhRvD..93k2008A|s2cid=55485404
 }}</ref>

The heavy isotope <sup>46</sup>Ca can also theoretically undergo double beta decay to <sup>46</sup>Ti as well, but this has never been observed. The lightest and most common isotope <sup>40</sup>Ca is also doubly magic and could undergo [[Physics:Double electron capture|double electron capture]] to <sup>40</sup>[[Chemistry:Argon|Ar]], 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 <sup>40</sup>Ca and <sup>46</sup>Ca are 5.9&nbsp;×&nbsp;10<sup>21</sup>&nbsp;years and 2.8&nbsp;×&nbsp;10<sup>15</sup>&nbsp;years respectively.{{NUBASE2016|ref}}

Apart from the practically stable <sup>48</sup>Ca, the longest lived radioisotope of calcium is <sup>41</sup>Ca. It decays by electron capture to stable <sup>41</sup>[[Chemistry:Potassium|K]] with a half-life of about a hundred thousand years. Its existence in the early Solar System as an [[Chemistry:Extinct radionuclide|extinct radionuclide]] has been inferred from excesses of <sup>41</sup>K: traces of <sup>41</sup>Ca also still exist today, as it is a [[Physics:Cosmogenic nuclide|cosmogenic nuclide]], continuously reformed through [[Physics:Neutron activation|neutron activation]] of natural <sup>40</sup>Ca.<ref name="Clayton" />

Many other calcium radioisotopes are known, ranging from <sup>35</sup>Ca to <sup>60</sup>Ca. They are all much shorter-lived than <sup>41</sup>Ca, the most stable among them being <sup>45</sup>Ca (half-life 163&nbsp;days) and <sup>47</sup>Ca (half-life 4.54&nbsp;days). The isotopes lighter than <sup>42</sup>Ca usually undergo beta plus decay to isotopes of potassium, and those heavier than <sup>44</sup>Ca usually undergo beta minus decay to isotopes of [[Chemistry:Scandium|scandium]], although near the [[Physics:Nuclear drip line|nuclear drip line]]s, [[Physics:Proton emission|proton emission]] and [[Physics:Neutron emission|neutron emission]] begin to be significant decay modes as well.{{NUBASE2016|ref}}

Like other elements, a variety of processes alter the relative abundance of calcium isotopes.<ref>{{Cite journal|last1=Russell|first1=W. A.|last2=Papanastassiou|first2=D. A.|last3=Tombrello|first3=T. A.|title=Ca isotope fractionation on the earth and other solar system materials|journal=Geochim Cosmochim Acta|date=1978|volume=42|pages=1075–90|doi=10.1016/0016-7037(78)90105-9|issue=8|bibcode = 1978GeCoA..42.1075R }}</ref> The best studied of these processes is the mass-dependent [[Physics:Isotope fractionation|fractionation]] of calcium isotopes that accompanies the precipitation of calcium minerals such as [[Chemistry:Calcite|calcite]], [[Chemistry:Aragonite|aragonite]] and [[Chemistry:Apatite|apatite]] 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 <sup>44</sup>Ca/<sup>40</sup>Ca) in a sample compared to the same ratio in a standard reference material. <sup>44</sup>Ca/<sup>40</sup>Ca varies by about 1% among common earth materials.<ref>{{Cite journal|last1=Skulan|first1=J.|last2=Depaolo|first2=D. J.|title=Calcium isotope fractionation between soft and mineralized tissues as a monitor of calcium use in vertebrates|journal=Proc Natl Acad Sci USA|date=1999|volume=96|pages=13709–13|doi=10.1073/pnas.96.24.13709|pmid=10570137|issue=24|pmc=24129 |bibcode = 1999PNAS...9613709S |doi-access=free}}</ref>

==History==
[[File:Ein Ghazal Venus.jpg|thumb|upright=0.4|One of the 'Ain Ghazal Statues, made from lime plaster]]
Calcium compounds were known for millennia, although their chemical makeup was not understood until the 17th century.<ref name="Greenwood108">Greenwood and Earnshaw, p. 108</ref> Lime as a [[Chemistry:Lime (material)|building material]]<ref name="minerals.usgs">{{cite web |title = Commodity report:Lime |publisher = United States Geological Survey | first = Lori E. | last = Apodaca |date=2021| url = http://minerals.usgs.gov/minerals/pubs/commodity/lime/390498.pdf | access-date= 2012-03-06}}</ref> and as plaster for statues was used as far back as around 7000&nbsp;BC.<ref>{{Cite journal |title=Burnt Lime Products and Social Implications in the Pre-Pottery Neolithic B Villages of the Near East |journal=Paléorient |first=Yosef |last=Garfinkel |date=1987 | doi = 10.3406/paleo.1987.4417 |jstor=41492234 |volume=13 |issue=1 |pages=69–76}}</ref> The first dated [[Chemistry:Lime kiln|lime kiln]] dates back to 2500&nbsp;BC and was found in Khafajah, [[Place:Mesopotamia|Mesopotamia]].<ref>{{cite book | page =4 | title =Lime Kilns and Lime Burning | isbn =978-0-7478-0596-0 | last1 =Williams | first1 =Richard | date =2004 | publisher =Bloomsbury USA }}</ref><ref>{{cite book | url = https://books.google.com/books?id=vHQsGAKAdYoC | title = Lime and Limestone: Chemistry and Technology, Production and Uses | isbn = 978-3-527-61201-7 | last1 = Oates | first1 = J. A. H | year = 2008| publisher = Wiley }}</ref>

At about the same time, dehydrated [[Chemistry:Gypsum|gypsum]] (CaSO<sub>4</sub>·2H<sub>2</sub>O) was being used in the Great Pyramid of Giza. This material would later be used for the plaster in the tomb of Tutankhamun. The ancient Romans instead used lime mortars made by heating [[Earth:Limestone|limestone]] (CaCO<sub>3</sub>). The name "calcium" itself derives from the Latin word ''calx'' "lime".<ref name="Greenwood108" />

Vitruvius noted that the lime that resulted was lighter than the original limestone, attributing this to the boiling of the water. In 1755, [[Biography:Joseph Black|Joseph Black]] proved that this was due to the loss of [[Chemistry:Carbon dioxide|carbon dioxide]], which as a gas had not been recognised by the ancient Romans.<ref name="Weeks" />

In 1789, [[Biography:Antoine Lavoisier|Antoine Lavoisier]] suspected that lime might be an oxide of a fundamental [[Chemistry:Chemical element|chemical element]]. 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 reasoned: {{blockquote|We are probably only acquainted as yet with a part of the metallic substances existing in nature, as all those which have a stronger affinity to oxygen than carbon possesses, are incapable, hitherto, of being reduced to a metallic state, and consequently, being only presented to our observation under the form of oxyds, are confounded with earths. It is extremely probable that barytes, which we have just now arranged with earths, is in this situation; for in many experiments it exhibits properties nearly approaching to those of metallic bodies. It is even possible that all the substances we call earths may be only metallic oxyds, irreducible by any hitherto known process.<ref>Lavoisier, Antoine; Kerr, Robert (translator) (1799) ''Elements of Chemistry'', 4th ed. Edinburgh, Scotland: William Creech. [https://archive.org/details/elementschemist00kerrgoog/page/n217 p. 218]. The original passage appears in: Lavoisier, Antoine (1789) ''[[Chemistry:Traité Élémentaire de Chimie|Traité Élémentaire de Chimie]]''. Paris, France: Cuchet. Vol. 1. [https://books.google.com/books?id=hZch3yOrayUC&pg=PA174 p. 174].</ref>}}

Calcium, along with its congeners magnesium, strontium, and barium, was first isolated by [[Biography:Humphry Davy|Humphry Davy]] in 1808. Following the work of Jöns Jakob Berzelius and Magnus Martin af Pontin on [[Chemistry:Electrolysis|electrolysis]], Davy isolated calcium and magnesium by putting a mixture of the respective metal oxides with [[Chemistry:Mercury(II) oxide|mercury(II) oxide]] on a [[Chemistry:Platinum|platinum]] 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.<ref name="Greenwood108" /><ref>{{cite journal | last1 = Davy | first1 = H. | date = 1808 | title = Electro-chemical researches on the decomposition of the earths; with observations on the metals obtained from the alkaline earths, and on the amalgam procured from ammonia | url = https://books.google.com/books?id=gpwEAAAAYAAJ&pg=102 | journal = Philosophical Transactions of the Royal Society of London | volume = 98 | pages = 333–70 | doi=10.1098/rstl.1808.0023| doi-access =  | bibcode = 1808RSPT...98..333D | s2cid = 96364168 }}</ref> 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.<ref name="Weeks">{{Cite book
|last = Weeks
|first = Mary Elvira
|year = 1968
|title = Discovery of the Elements
|publisher = Journal of Chemical Education
|pages =505–10
|location = Easton, PA
|lccn = 68-15217
|ref = CITEREFWeeks1968
|isbn = 978-0-7661-3872-8
}}</ref>

==Occurrence and production==
[[File:Pamukkale Hierapolis Travertine pools.JPG|thumb|[[Earth:Travertine|Travertine]] terraces in Pamukkale, Turkey]]
At 3%, calcium is the fifth [[Earth:Abundance of elements in Earth's crust|most abundant element in the Earth's crust]], and the third most abundant metal behind [[Chemistry:Aluminium|aluminium]] and [[Chemistry:Iron|iron]].<ref name="Greenwood108" /> It is also the fourth most abundant element in the lunar highlands.<ref name="Ullmann483" /> Sedimentary [[Chemistry:Calcium carbonate|calcium carbonate]] deposits pervade the Earth's surface as fossilized remains of past marine life; they occur in two forms, the rhombohedral [[Chemistry:Calcite|calcite]] (more common) and the [[Chemistry:Orthorhombic|orthorhombic]] [[Chemistry:Aragonite|aragonite]] (forming in more temperate seas). Minerals of the first type include [[Earth:Limestone|limestone]], [[Chemistry:Dolomite (mineral)|dolomite]], [[Chemistry:Marble|marble]], [[Earth:Chalk|chalk]], and [[Physics:Iceland spar|iceland spar]]; aragonite beds make up the Bahamas, the Florida Keys, and the Red Sea basins. [[Biology:Coral|Coral]]s, sea shells, and [[Chemistry:Pearl|pearl]]s are mostly made up of calcium carbonate. Among the other important minerals of calcium are [[Chemistry:Gypsum|gypsum]] (CaSO<sub>4</sub>·2H<sub>2</sub>O), [[Chemistry:Anhydrite|anhydrite]] (CaSO<sub>4</sub>), [[Chemistry:Fluorite|fluorite]] (CaF<sub>2</sub>), and [[Chemistry:Apatite|apatite]] ([Ca<sub>5</sub>(PO<sub>4</sub>)<sub>3</sub>X], X = OH, Cl, or F).<ref name="Greenwood108" />

The major producers of calcium are {{wipe|China}} (about 10000 to 12000 [[Tonne|tonne]]s per year), {{wipe|Russia}} (about 6000 to 8000 tonnes per year), and the {{wipe|United States}} (about 2000 to 4000 tonnes per year). {{wipe|Canada}} and {{wipe|France}} 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.<ref name="Ullmann484" />

In Russia and China, Davy's method of electrolysis is still used, but is instead applied to molten [[Chemistry:Calcium chloride|calcium chloride]].<ref name="Ullmann484" /> Since calcium is less reactive than strontium or barium, the oxide–nitride coating that results in air is stable and [[Engineering:Lathe|lathe]] machining and other standard metallurgical techniques are suitable for calcium.<ref name="Greenwood110">Greenwood and Earnshaw, p. 110</ref> In the United States and Canada, calcium is instead produced by reducing lime with aluminium at high temperatures.<ref name="Ullmann484" />

===Geochemical cycling===
{{Main|Chemistry:Carbonate–silicate cycle}}
[[Biology:Calcium cycle|Calcium cycling]] provides a link between [[Earth:Tectonics|tectonics]], [[Earth:Climate|climate]], and the [[Earth:Carbon cycle|carbon cycle]]. In the simplest terms, uplift of mountains exposes calcium-bearing rocks such as some [[Chemistry:Granite|granite]]s to chemical weathering and releases Ca<sup>2+</sup> into surface water. These ions are transported to the ocean where they react with dissolved CO<sub>2</sub> to form [[Earth:Limestone|limestone]] ({{chem|CaCO|3}}), which in turn settles to the sea floor where it is incorporated into new rocks. Dissolved CO<sub>2</sub>, along with [[Chemistry:Carbonate|carbonate]] and [[Chemistry:Bicarbonate|bicarbonate]] ions, are termed "[[Chemistry:Total inorganic carbon|dissolved inorganic carbon]]" (DIC).<ref name="Berner" />

The actual reaction is more complicated and involves the bicarbonate ion (HCO{{su|b=3|p=−}}) that forms when CO<sub>2</sub> reacts with water at seawater [[Chemistry:PH|pH]]:
:<chem>Ca^2+ + 2 HCO3- -> CaCO3_v + CO2 + H2O</chem>
At seawater pH, most of the CO<sub>2</sub> is immediately converted back into {{chem|HCO|3|-}}. The reaction results in a net transport of one molecule of CO<sub>2</sub> from the ocean/atmosphere into the [[Earth:Lithosphere|lithosphere]].<ref>{{cite web|author= Zeebe|date=2006|url = http://www.eoearth.org/article/Marine_carbonate_chemistry|title = Marine carbonate chemistry|publisher = National Council for Science and the Environment|access-date = 2010-03-13}}</ref> The result is that each Ca<sup>2+</sup> ion released by chemical weathering ultimately removes one CO<sub>2</sub> 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 CO<sub>2</sub> from the ocean and atmosphere, exerting a strong long-term effect on climate.<ref name="Berner">{{Cite journal|last1=Berner|first1=Robert|title= The long-term carbon cycle, fossil fuels and atmospheric composition |journal=Nature|date=2003|volume=426|pages= 323–26|doi=10.1038/nature02131|pmid=14628061|issue=6964|bibcode = 2003Natur.426..323B |s2cid=4420185}}</ref><ref>{{Cite journal|title = A negative feedback mechanism for the long-term stabilization of Earth's surface temperature|journal = Journal of Geophysical Research: Oceans|date = 1981-10-20|pages = 9776–82|volume = 86|issue = C10|doi = 10.1029/JC086iC10p09776|first1 = James C. G.|last1 = Walker|first2 = P. B.|last2 = Hays|first3 = J. F.|last3 = Kasting|bibcode=1981JGR....86.9776W}}</ref>

==Applications==
The largest use of metallic calcium is in [[Engineering:Steelmaking|steelmaking]], due to its strong [[Chemistry:Chemical affinity|chemical affinity]] for oxygen and [[Chemistry:Sulfur|sulfur]]. Its oxides and sulfides, once formed, give liquid lime [[Chemistry:Aluminate|aluminate]] and sulfide inclusions in steel which float out; on treatment, these inclusions disperse throughout the steel and become small and spherical, improving castability, cleanliness and general mechanical properties. Calcium is also used in maintenance-free [[Engineering:Automotive battery|automotive batteries]], in which the use of 0.1% calcium–[[Chemistry:Lead|lead]] alloys instead of the usual [[Chemistry:Antimony|antimony]]–lead alloys leads to lower water loss and lower self-discharging.<ref name="Ullmann485"/>

Due to the risk of expansion and cracking, [[Chemistry:Aluminium|aluminium]] is sometimes also incorporated into these alloys. These lead–calcium alloys are also used in casting, replacing lead–antimony alloys.<ref name="Ullmann485">Hluchan and Pomerantz, pp. 485–87</ref> Calcium is also used to strengthen aluminium alloys used for bearings, for the control of graphitic [[Chemistry:Carbon|carbon]] in [[Chemistry:Cast iron|cast iron]], and to remove [[Chemistry:Bismuth|bismuth]] impurities from lead.<ref name="Greenwood110" /> Calcium metal is found in some drain cleaners, where it functions to generate heat and [[Chemistry:Calcium hydroxide|calcium hydroxide]] that [[Chemistry:Saponification|saponifies]] the fats and liquefies the proteins (for example, those in hair) that block drains.<ref name="Rumack" />

Besides metallurgy, the reactivity of calcium is exploited to remove [[Chemistry:Nitrogen|nitrogen]] from high-purity [[Chemistry:Argon|argon]] gas and as a [[Engineering:Getter|getter]] for oxygen and nitrogen. It is also used as a reducing agent in the production of [[Chemistry:Chromium|chromium]], [[Chemistry:Zirconium|zirconium]], [[Chemistry:Thorium|thorium]], [[Chemistry:Vanadium|vanadium]] and [[Chemistry:Uranium|uranium]]. It can also be used to store hydrogen gas, as it reacts with hydrogen to form solid [[Chemistry:Calcium hydride|calcium hydride]], from which the hydrogen can easily be re-extracted.<ref name="Greenwood110" />

Calcium isotope fractionation during mineral formation has led to several applications of calcium isotopes. In particular, the 1997 observation by Skulan and DePaolo<ref>{{Cite journal|last1=Skulan|first1=J.|title=Biological control of calcium isotopic abundances in the global calcium cycle|journal=Geochimica et Cosmochimica Acta |date=June 1997 | volume=61|pages=2505–10|last2=Depaolo|first2=D. J. | first3 =T. L.| last3= Owens |issue=12|doi=10.1016/S0016-7037(97)00047-1 | bibcode = 1997GeCoA..61.2505S }}</ref> that calcium minerals are isotopically lighter than the solutions from which the minerals precipitate is the basis of analogous applications in medicine and in paleoceanography. 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.<ref name="ReferenceA">{{Cite journal|last1=Skulan|first1=J.|title=Natural calcium isotopic composition of urine as a marker of bone mineral balance|pmid=17463176|journal=Clinical Chemistry|date=2007|volume=53|pages=1155–58|last2=Bullen|first2=T.|last3=Anbar|first3=A. D.|last4=Puzas|first4=J. E.|last5=Shackelford|first5=L.|last6=Leblanc|first6=A.|last7=Smith|first7=S. M.|issue=6|doi=10.1373/clinchem.2006.080143|doi-access=free}}</ref>

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 <sup>44</sup>Ca/<sup>40</sup>Ca 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 [[Medicine:Osteoporosis|osteoporosis]].<ref name="ReferenceA"/>

A similar system exists in seawater, where <sup>44</sup>Ca/<sup>40</sup>Ca tends to rise when the rate of removal of Ca<sup>2+</sup> 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 <sup>44</sup>Ca/<sup>40</sup>Ca over geologic time, along with a theoretical explanation of these changes. More recent papers have confirmed this observation, demonstrating that seawater Ca<sup>2+</sup> 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 [[Earth:Carbon cycle|carbon cycle]].<ref>{{Cite journal|last1=Fantle|first1=M.|last2=Depaolo|first2=D.|title=Ca isotopes in carbonate sediment and pore fluid from ODP Site 807A: The Ca<sup>2+</sup>(aq)–calcite equilibrium fractionation factor and calcite recrystallization rates in Pleistocene sediments|journal=Geochim Cosmochim Acta|date=2007|volume=71|pages=2524–46|doi=10.1016/j.gca.2007.03.006|issue=10|bibcode=2007GeCoA..71.2524F}}</ref><ref>{{Cite journal|last1=Griffith|first1=Elizabeth M.|title=A Dynamic marine calcium cycle during the past 28 million years|pmid=19074345|journal=Science|date=2008|volume=322|pages=1671–74|last2= Paytan |first2= Adina |last3= Caldeira |first3= Ken|last4= Bullen |first4= Thomas |last5= Thomas |first5= Ellen |author5-link=Ellen Thomas (scientist) |s2cid=206515318|issue=12|doi=10.1126/science.1163614 |bibcode = 2008Sci...322.1671G }}</ref>

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 [[Chemistry:Calcium lactate|calcium lactate]], calcium diphosphate, and [[Chemistry:Tricalcium phosphate|tricalcium phosphate]]. The last is also used as a polishing agent in [[Medicine:Toothpaste|toothpaste]] and in [[Chemistry:Antacid|antacid]]s. Calcium lactobionate is a white powder that is used as a suspending agent for pharmaceuticals. In baking, [[Chemistry:Calcium phosphate|calcium phosphate]] is used as a leavening agent. [[Chemistry:Calcium sulfite|Calcium sulfite]] is used as a bleach in papermaking and as a disinfectant, [[Chemistry:Calcium silicate|calcium silicate]] is used as a reinforcing agent in rubber, and [[Chemistry:Calcium acetate|calcium acetate]] is a component of liming rosin and is used to make metallic soaps and synthetic resins.<ref name="Ullmann485" />

Calcium is on the [[Medicine:WHO Model List of Essential Medicines|World Health Organization's List of Essential Medicines]].<ref name="WHO21st">{{cite book | vauthors = ((World Health Organization)) | title = World Health Organization model list of essential medicines: 21st list 2019 | year = 2019 | hdl = 10665/325771 |  publisher = World Health Organization | location = Geneva | id = WHO/MVP/EMP/IAU/2019.06. License: CC BY-NC-SA 3.0 IGO | hdl-access=free }}</ref>

==Food sources==
Foods rich in calcium include dairy products, such as yogurt and cheese, [[Biology:Sardine|sardine]]s, salmon, soy products, kale, and [[Chemistry:Food fortification|fortified]] [[Engineering:Breakfast cereal|breakfast cereal]]s.<ref name="ods" />

Because of concerns for long-term adverse side effects, including calcification of arteries and kidney stones, both the U.S. Institute of Medicine (IOM) and the European Food Safety Authority (EFSA) set Tolerable Upper Intake Levels (ULs) for combined dietary and supplemental calcium. From the IOM, people of ages 9–18 years are not to exceed 3&nbsp;g/day combined intake; for ages 19–50, not to exceed 2.5&nbsp;g/day; for ages 51 and older, not to exceed 2&nbsp;g/day.<ref>{{cite book | title = Dietary Reference Intakes for Calcium and Vitamin D|chapter =ch 6. Tolerable Upper Intake Levels |pages =403–56 | publisher = National Academies Press | location = Washington, D.C | year = 2011 | isbn = 978-0-309-16394-1 | url = https://www.nap.edu/read/13050/chapter/8| doi = 10.17226/13050 | pmid = 21796828 | author1 = Institute of Medicine (US) Committee to Review Dietary Reference Intakes for Vitamin D Calcium | last2 = Ross | first2 = A. C. | last3 = Taylor | first3 = C. L. | last4 = Yaktine | first4 = A. L. | last5 = Del Valle | first5 = H. B. |s2cid =58721779 }}</ref> EFSA set the UL for all adults at 2.5&nbsp;g/day, but decided the information for children and adolescents was not sufficient to determine ULs.<ref>{{citation| title = Tolerable Upper Intake Levels For Vitamins And Minerals| publisher = European Food Safety Authority| year = 2006| url = http://www.efsa.europa.eu/sites/default/files/efsa_rep/blobserver_assets/ndatolerableuil.pdf}}</ref>

==Biological and pathological role==
{{main|Biology:Calcium in biology}}
{| class="wikitable" style="float:right;"
|+ Age-adjusted daily calcium recommendations (from U.S. Institute of Medicine RDAs)<ref>{{cite book | title = Dietary Reference Intakes for Calcium and Vitamin D|chapter =ch. 5. Dietary Reference Intakes |pages =345–402 | publisher = National Academies Press | location = Washington, D.C | year = 2011 | isbn = 978-0-309-16394-1 | url = https://www.nap.edu/read/13050/chapter/7| doi = 10.17226/13050 | pmid = 21796828 | author1 = Institute of Medicine (US) Committee to Review Dietary Reference Intakes for Vitamin D Calcium | last2 = Ross | first2 = A. C. | last3 = Taylor | first3 = C. L. | last4 = Yaktine | first4 = A. L. | last5 = Del Valle | first5 = H. B. |s2cid =58721779 }}</ref>
|-
! Age
! Calcium (mg/day)
|-
| 1–3 years
| 700
|-
| 4–8 years
| 1000
|-
| 9–18 years
| 1300
|-
| 19–50 years
| 1000
|-
| >51 years
| 1000
|-
| Pregnancy
| 1000
|-
| Lactation
| 1000
|}

[[File:Calcium_intake_world_map.svg|thumb|upright=1.4|Global dietary calcium intake among adults (mg/day).<ref>{{cite journal | vauthors = Balk EM, Adam GP, Langberg VN, Earley A, Clark P, Ebeling PR, Mithal A, Rizzoli R, Zerbini CA, Pierroz DD, Dawson-Hughes B | title = Global dietary calcium intake among adults: a systematic review | journal = Osteoporosis International | volume = 28 | issue = 12 | pages = 3315–24 | date = December 2017 | pmid = 29026938 | pmc = 5684325 | doi = 10.1007/s00198-017-4230-x }}</ref>
{{Div col|small=yes|colwidth=10em}}
{{legend|#f51d1e|<400}}{{legend|#aa2b2c|400–500}}{{legend|#fa7929|500–600}}{{legend|#a46e2c|600–700}}{{legend|#fee940|700–800}}{{legend|#829d29|800–900}}{{legend|#84fc5a|900–1000}}{{legend|#2d621d|>1000}}
{{div col end}}
]]

===Function===
Calcium is an essential element needed in large quantities.<ref name="lpi" /><ref name="ods" /> The Ca<sup>2+</sup> ion acts as an [[Physics:Electrolyte|electrolyte]] 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 [[Biology:Muscle contraction|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 Ca<sup>2+</sup> ion forms stable [[Chemistry:Coordination complex|coordination complex]]es with many organic compounds, especially [[Biology:Protein|protein]]s; it also forms compounds with a wide range of solubilities, enabling the formation of the [[Biology:Skeleton|skeleton]].<ref name="lpi" />
<ref>
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}}</ref>

===Binding===
Calcium ions may be complexed by proteins through binding the carboxyl groups of [[Chemistry:Glutamic acid|glutamic acid]] or [[Chemistry:Aspartic acid|aspartic acid]] residues; through interacting with [[Biology:Phosphorylation|phosphorylated]] [[Chemistry:Serine|serine]], [[Chemistry:Tyrosine|tyrosine]], or [[Chemistry:Threonine|threonine]] residues; or by being [[Chemistry:Chelation|chelated]] by γ-carboxylated amino acid residues. [[Biology:Trypsin|Trypsin]], a digestive enzyme, uses the first method; [[Biology:Osteocalcin|osteocalcin]], a bone matrix protein, uses the third.<ref name="Ullmann489"/>

Some other bone matrix proteins such as osteopontin and [[Biology:Bone sialoprotein|bone sialoprotein]] 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 [[Chemistry:Phospholipid|phospholipid]] layer of the [[Biology:Cell membrane|cell membrane]], anchoring proteins associated with the cell surface.<ref name="Ullmann489">Hluchan and Pomerantz, pp. 489–94</ref>

===Solubility===
As an example of the wide range of solubility of calcium compounds, [[Chemistry:Monocalcium phosphate|monocalcium phosphate]] is very soluble in water, 85% of extracellular calcium is as [[Chemistry:Dicalcium phosphate|dicalcium phosphate]] with a solubility of 2.00&nbsp;[[Physics:Molar concentration|mM]], and the [[Chemistry:Hydroxyapatite|hydroxyapatite]] of bones in an organic matrix is [[Chemistry:Tricalcium phosphate|tricalcium phosphate]] with a solubility of 1000&nbsp;μM.<ref name="Ullmann489" />

=== Nutrition ===
Calcium is a common constituent of [[Chemistry:Multivitamin|multivitamin]] dietary supplements,<ref name="lpi" /> but the composition of calcium complexes in supplements may affect its [[Chemistry:Bioavailability|bioavailability]] which varies by solubility of the salt involved: [[Chemistry:Calcium citrate|calcium citrate]], [[Chemistry:Calcium malate|malate]], and [[Chemistry:Calcium lactate|lactate]] are highly bioavailable, while the [[Chemistry:Calcium oxalate|oxalate]] is less. Other calcium preparations include [[Chemistry:Calcium carbonate|calcium carbonate]], [[Chemistry:Calcium citrate malate|calcium citrate malate]], and [[Chemistry:Calcium gluconate|calcium gluconate]].<ref name="lpi" /> The intestine absorbs about one-third of calcium eaten as the [[Physics:Radical ion|free ion]], and plasma calcium level is then regulated by the kidneys.<ref name="lpi" />

===Hormonal regulation of bone formation and serum levels===
[[Biology:Parathyroid hormone|Parathyroid hormone]] and [[Chemistry:Vitamin D|vitamin D]] 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.<ref name="lpi" /> 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 [[Biology:Calcitonin|calcitonin]], whose secretion increases with increasing plasma calcium levels.<ref name="Ullmann489" />

===Abnormal serum levels===
Excess intake of calcium may cause hypercalcemia. 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.<ref name="Ullmann489" />

Chronic hypercalcaemia typically leads to [[Biology:Calcification|calcification]] of soft tissue and its serious consequences: for example, calcification can cause loss of elasticity of vascular walls and disruption of laminar blood flow—and thence to [[Biology:Vulnerable plaque|plaque rupture]] and [[Medicine:Thrombosis|thrombosis]]. Conversely, inadequate calcium or vitamin D intakes may result in [[Medicine:Hypocalcemia|hypocalcemia]], often caused also by inadequate secretion of parathyroid hormone or defective PTH receptors in cells. Symptoms include neuromuscular excitability, which potentially causes [[Medicine:Tetany|tetany]] and disruption of conductivity in cardiac tissue.<ref name="Ullmann489" />

===Bone disease===
As calcium is required for bone development, many bone diseases can be traced to the organic matrix or the [[Chemistry:Hydroxyapatite|hydroxyapatite]] in molecular structure or organization of bone. [[Medicine:Osteoporosis|Osteoporosis]] is a reduction in mineral content of bone per unit volume, and can be treated by supplementation of calcium, vitamin D, and [[Chemistry:Bisphosphonate|bisphosphonate]]s.<ref name="lpi" /><ref name="ods" /> Inadequate amounts of calcium, vitamin D, or phosphates can lead to softening of bones, called [[Medicine:Osteomalacia|osteomalacia]].<ref name="Ullmann489" />

==Safety==
===Metallic calcium===
{{Chembox
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| GHSPictograms = {{GHS02}}
| GHSSignalWord = Danger
| HPhrases = {{H-phrases|H261}}
| PPhrases = {{P-phrases|P231+P232}}
| GHS_ref  = <ref>{{cite web|url=https://www.sigmaaldrich.com/catalog/product/aldrich/215414?lang=en&region=US |title=Calcium turnings, 99% trace metals basis |publisher=Sigma-Aldrich |date=2021-02-24 |access-date=2021-12-22}}</ref>
| NFPA-H = 0
| NFPA-F = 3
| NFPA-R = 1
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}}
Because calcium reacts exothermically with water and acids, calcium metal coming into contact with bodily moisture results in severe corrosive irritation.<ref name="Ullmann487" /> When swallowed, calcium metal has the same effect on the mouth, oesophagus, and stomach, and can be fatal.<ref name="Rumack">Rumack BH. POISINDEX. Information System Micromedex, Inc., Englewood, CO, 2010; CCIS Volume 143. Hall AH and Rumack BH (Eds)</ref> However, long-term exposure is not known to have distinct adverse effects.<ref name="Ullmann487">Hluchan and Pomerantz, pp. 487–89</ref>

==References==
{{Reflist|30em}}

==Bibliography==
* {{Ullmann | first1=Stephen E. |last1=Hluchan |first2=Kenneth |last2=Pomerantz | title = Calcium and Calcium Alloys | doi = 10.1002/14356007.a04_515.pub2}}

{{Periodic table (navbox)}}
{{Calcium compounds}}
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{{Subject bar
|portal=Chemistry
|book1=Calcium
|book2=Period 4 elements
|book3=Alkaline earth metals
|book4=Chemical elements (sorted&nbsp;alphabetically)
|book5=Chemical elements (sorted by number)
|commons=y
|wikt=y
|wikt-search=calcium
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[[Category:Calcium| ]]
[[Category:Chemical elements]]
[[Category:Alkaline earth metals]]
[[Category:Dietary minerals]]
[[Category:Reducing agents]]
[[Category:Sodium channel blockers]]
[[Category:World Health Organization essential medicines]]
[[Category:Chemical elements with face-centered cubic structure]]

{{Sourceattribution|Calcium}}