Calcium is achemical element; it hassymbolCa andatomic number 20. As analkaline 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 homologuesstrontium andbarium. It is the fifth most abundant element in Earth's crust, and the third most abundant metal, afteriron andaluminium. The most common calcium compound on Earth iscalcium carbonate, found inlimestone and the fossils of early sea life;gypsum,anhydrite,fluorite, andapatite are also sources of calcium. The name comes fromLatincalx "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 viaelectrolysis of its oxide byHumphry Davy, who named the element. Calcium compounds are widely used in many industries: in foods and pharmaceuticals forcalcium 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 a very ductile silvery metal (sometimes described as pale yellow) whose properties are very similar to the heavier elements in its group,strontium,barium, andradium. A calcium atom has 20 electrons, withelectron configuration [Ar]4s2. Like the other elements in group 2 of the periodic table, calcium has twovalence electrons in the outermost s-orbital, which are very easily lost in chemical reactions to form a dipositive ion with the stable electron configuration of anoble gas, in this caseargon.[10]
Hence, calcium is almost alwaysdivalent in its compounds, which are usuallyionic. Hypothetical univalent salts of calcium would be stable with respect to their elements, but not todisproportionation to the divalent salts and calcium metal, because theenthalpy of formation of MX2 is much higher than those of the hypothetical MX. This occurs because of the much greaterlattice energy afforded by the more highly charged Ca2+ cation compared to the hypothetical Ca+ cation.[10]
Calcium, strontium, barium, and radium are always considered to bealkaline earth metals; the lighterberyllium andmagnesium, 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 behavior: they behave more likealuminium andzinc respectively and have some of the weaker metallic character of thepost-transition metals, which is why the traditional definition of the term "alkaline earth metal" excludes them.[11]
Physical properties
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 theface-centered cubic arrangement like strontium and barium; above 443 °C (716 K), it changes tobody-centered cubic.[4][12] Its density of 1.526 g/cm3 (at 20 °C)[4] is the lowest in its group.[10]
Calcium is harder thanlead but can be cut with a knife with effort. While calcium is a poorer conductor of electricity thancopper oraluminium by volume, it is a better conductor by mass than both due to its very low density.[13] 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.[13]
Chemical properties
Structure of the polymeric [Ca(H2O)6]2+ 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 but less quickly than strontium to producecalcium hydroxide and hydrogen gas. It also reacts with theoxygen andnitrogen in air to form a mixture ofcalcium oxide andcalcium nitride.[14] When finely divided, it spontaneously burns in air to produce the nitride. Bulk calcium is less reactive: it quickly forms a hydration coating in moist air, but below 30%relative humidity it may be stored indefinitely at room temperature.[15]
Besides the simple oxide CaO,calcium peroxide, CaO2, can be made by direct oxidation of calcium metal under a high pressure of oxygen, and there is some evidence for a yellowsuperoxide Ca(O2)2.[16]Calcium hydroxide, Ca(OH)2, is a strong base, though not as strong as the hydroxides of strontium, barium or the alkali metals.[17] All four dihalides of calcium are known.[18]Calcium carbonate (CaCO3) andcalcium sulfate (CaSO4) are particularly abundant minerals.[19] Like strontium and barium, as well as the alkali metals and the divalentlanthanideseuropium andytterbium, calcium metal dissolves directly in liquidammonia to give a dark blue solution.[20]
In contrast toorganomagnesium compounds, organocalcium compounds are not similarly useful, with one major exception,calcium carbide, CaC2. This material, which has historic significance, is prepared by heating calcium oxide with carbon. According toX-ray crystallography, calcium carbide can be described as Ca2+ derivative of acetylide, C22-, although it is not a salt. Several million tons of calcium carbide are produced annually. Hydrolysis givesacetylene, which is used in welding and a chemical precursor. Reaction with nitrogen gas converts calcium carbide tocalcium cyanamide.[22]
A dominant theme in molecular organocalcium chemistry is the large radius of calcium, which often leads to highcoordination numbers. For example, dimethylcalcium appears to be a 3-dimensional polymer,[23] whereasdimethylmagnesium is a linear polymer with tetrahedral Mg centers.Bulky ligands are often required to disfavor polymeric species. For example, calcium dicyclopentadienyl,Ca(C5H5)2 has a polymeric structure and thus is nonvolatile and insoluble in solvents. Replacing theC5H5 ligand with the bulkierC5(CH3)5 (pentamethylcyclopentadienyl) gives a soluble complex that sublimes and forms well-defined adducts with ethers.[19] Organocalcium compounds tend to be more similar toorganoytterbium compounds due to the similarionic radii of Yb2+ (102 pm) and Ca2+ (100 pm).[24]
Organocalcium compounds have been well investigated. Some such complexes exhibit catalytic properties,[25] although none have been commercialized.
Natural calcium is a mixture of five stableisotopes—40Ca,42Ca,43Ca,44Ca, and46Ca—and48Ca, whose half-life of 4.3 × 1019 years is so long that it can be considered stable for all practical purposes. Calcium is the first (lightest) element to have six naturally occurring isotopes.[14]
By far the most common isotope is40Ca, which makes up 96.941% of natural calcium. It is produced in thesilicon-burning process from fusion ofalpha particles and is the heaviest stable nuclide with equal proton and neutron numbers; its occurrence is also supplemented slowly by the decay ofprimordial40K. Adding another alpha particle leads to unstable44Ti, which decays via two successiveelectron captures to stable44Ca; this makes up 2.806% of natural calcium and is the second-most common isotope.[26][27]
The other four natural isotopes,42, 43, 46, 48Ca, are significantly rarer, each comprising less than 1% of natural calcium. The four lighter isotopes are mainly products ofoxygen-burning and silicon-burning, leaving the two heavier ones to be produced vianeutron capture.46Ca is mostly produced in a "hot"s-process, as its formation requires a rather high neutron flux to allow short-lived45Ca to capture a neutron.48Ca is produced by electron capture in ther-process intype Ia supernovae, where high neutron excess and low enough entropy ensures its survival.[26][27]
46Ca and48Ca are the first "classically stable" nuclides with a 6-neutron or 8-neutron excess respectively. Though extremely neutron-rich for such a light element,48Ca is very stable because it is adoubly magic nucleus, with 20 protons and 28 neutrons arranged in closed shells. Itsbeta decay to48Sc is very hindered by the gross mismatch ofnuclear spin:48Ca has zero nuclear spin, beingeven–even, while48Sc has spin 6+, so the decay isforbidden by conservation ofangular momentum. While two excited states of48Sc are available for decay as well, they are also forbidden due to their high spins. As a result, when48Ca does decay, it does so bydouble beta decay to48Ti instead, being the lightest nuclide known to undergo double beta decay.[28][29]
46Ca can also theoretically double-beta-decay to46Ti, but this has never been observed. The most common isotope40Ca is also doubly magic and could undergodouble electron capture to40Ar, but this has likewise never been observed. Calcium is the only element with two primordial doubly magic isotopes. The experimental lower limits for the half-lives of40Ca and46Ca are 5.9 × 1021 years and 2.8 × 1015 years respectively.[28]
Excluding48Ca, the longest livedradioisotope of calcium is41Ca. It decays by electron capture to stable41K with a half-life of about 105 years. Its existence in the early Solar System as anextinct radionuclide has been inferred from excesses of41K. Traces of41Ca also still exist today, as it is acosmogenic nuclide, continuously produced throughneutron activation of natural40Ca.[27]
Many other calcium radioisotopes are known, ranging from35Ca to60Ca. They are all much shorter-lived than41Ca; the most stable are45Ca (half-life 163 days) and47Ca (half-life 4.54 days). Isotopes lighter than42Ca usually undergobeta plus decay to isotopes of potassium, and those heavier than44Ca usually undergobeta minus decay toscandium; though near thenuclear drip lines,proton emission andneutron emission begin to be significant decay modes as well.[28]
Like other elements, a variety of processes alter the relative abundance of calcium isotopes.[30] The best studied of these processes is the mass-dependentfractionation of calcium isotopes that accompanies the precipitation of calcium minerals such ascalcite,aragonite andapatite 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 (usually44Ca/40Ca) in a sample compared to the same ratio in a standard reference material.44Ca/40Ca varies by about 1–2‰ among organisms on Earth.[31]
About the same time, dehydratedgypsum (CaSO4·2H2O) was being used in theGreat Pyramid of Giza. This material would later be used for the plaster in the tomb ofTutankhamun. Theancient Romans instead used lime mortars made by heatinglimestone (CaCO3). The name "calcium" itself derives from the Latin wordcalx "lime".[32]
Vitruvius noted that the lime that resulted was lighter than the original limestone, attributing this to the boiling of the water. In 1755,Joseph Black proved that this was due to the loss ofcarbon dioxide, which as a gas had not been recognised by the ancient Romans.[37]
In 1789,Antoine Lavoisier suspected that lime might be an oxide of an 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), andsilice (silica, silicon dioxide)). About these "elements", Lavoisier reasoned:
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.[38]
Calcium, along with its congeners magnesium, strontium, and barium, was first isolated byHumphry Davy in 1808. Following the work ofJöns Jakob Berzelius andMagnus Martin of Pontin onelectrolysis, Davy isolated calcium and magnesium by putting a mixture of the respective metal oxides withmercury(II) oxide on aplatinum 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.[32][39] 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.[37]
The major producers of calcium areChina (about 10000 to 12000tonnes per year),Russia (about 6000 to 8000 tonnes per year), and theUnited States (about 2000 to 4000 tonnes per year).Canada andFrance are 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.[13]
In Russia and China, Davy's method of electrolysis is still used, but is instead applied to moltencalcium chloride.[13] Since calcium is less reactive than strontium or barium, the oxide–nitride coating that results in air is stable andlathe machining and other standard metallurgical techniques are suitable for calcium.[41]
In the U.S. and Canada, calcium is instead produced by reducinglime with aluminium at high temperatures.[13] In this process, powdered high-calcium lime and powdered aluminum are mixed and compacted intobriquettes for a high degree of contact, which are then placed in a sealedretort which has beenevacuated and heated to ~1200°C.[13] The briquettes release calcium vapor into the vacuum for about 8 hours, which then condenses in the cooled ends of the retorts to form 24-34 kg pieces of calcium metal, as well as some residue ofcalcium aluminate.[13] High-purity calcium can be obtained bydistilling low-purity calcium at high temperatures.[13]
Calcium cycling provides a link betweentectonics,climate, and thecarbon cycle. In the simplest terms, mountain-building exposes calcium-bearing rocks such asbasalt andgranodiorite to chemical weathering and releases Ca2+ into surface water. These ions are transported to the ocean where they react with dissolved CO2 to formlimestone (CaCO3), which in turn settles to the sea floor where it is incorporated into new rocks. Dissolved CO2, along withcarbonate andbicarbonate ions, are termed "dissolved inorganic carbon" (DIC).[42]
The actual reaction is more complicated and involves the bicarbonate ion (HCO− 3) that forms when CO2 reacts with water at seawaterpH:
Ca2+ + 2 HCO−3 → CaCO3↓ + CO2 + H2O
At seawater pH, most of the dissolved CO2 is immediately converted back intoHCO− 3. The reaction results in a net transport of one molecule of CO2 from the ocean/atmosphere into thelithosphere.[43] The result is that each Ca2+ ion released by chemical weathering ultimately removes one CO2 molecule from the surficial system (atmosphere, ocean, soil 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 air, exerting a strong long-term effect on climate.[42][44]
Applications
The largest use of metallic calcium is insteelmaking, due to its strongchemical affinity for chalcogens oxygen andsulfur. Its oxides and sulfides, once formed, give liquid limealuminate 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-freeautomotive batteries, in which the use of 0.1% calcium–lead alloys instead of the usualantimony–lead alloys leads to lower water loss and lower self-discharging.[45]
Due to the risk of expansion and cracking,aluminium is sometimes also incorporated into these alloys. These lead–calcium alloys are also used in casting, replacing lead–antimony alloys.[45] Calcium is also used to strengthen aluminium alloys used for bearings, for the control of graphiticcarbon incast iron, and to removebismuth impurities from lead.[41] Calcium metal is found in some drain cleaners, where it functions to generate heat andcalcium hydroxide thatsaponifies the fats and liquefies the proteins (for example, those in hair) that block drains.[46]
Besides metallurgy, the reactivity of calcium is exploited to removenitrogen from high-purityargon gas and as agetter for oxygen and nitrogen. It is also used as a reducing agent in the production ofchromium,zirconium,thorium,vanadium anduranium. It can also be used to store hydrogen gas, as it reacts with hydrogen to form solidcalcium hydride, from which the hydrogen can easily be re-extracted.[41]
Calcium isotope fractionation during mineral formation has led to several applications of calcium isotopes. In particular, the 1997 observation by Skulan and DePaolo[47] 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 mineralised with calcium, the calcium isotopic composition of soft tissues reflects the relative rate of formation and dissolution of skeletal mineral.[48]
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, the44Ca/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 likeosteoporosis.[48]
A similar system exists in seawater, where44Ca/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 seawater44Ca/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 thecarbon cycle.[49][50]
Because of concerns for long-term adverse side effects, including calcification of arteries andkidney stones, both the U.S.Institute of Medicine (IOM) and theEuropean Food Safety Authority (EFSA) settolerable upper intake levels (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.[52] 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.[53]
Age-adjusted daily calcium recommendations (from U.S. Institute of Medicine RDAs)[54]
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
Global dietary calcium intake among adults (mg/day).[55]
<400
400–500
500–600
600–700
700–800
800–900
900–1000
>1000
Function
Calcium is anessential element needed in large quantities.[8][9] The Ca2+ ion acts as anelectrolyte and is vital to the health of the muscular, circulatory, and digestive systems; is indispensable to the building of bone in the form ofhydroxyapatite; and supports synthesis and function of blood cells. For example, it regulates thecontraction 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 stablecoordination complexes with many organic compounds, especiallyproteins; it also forms compounds with a wide range of solubilities, enabling the formation of theskeleton.[8][56]
Some other bone matrix proteins such asosteopontin andbone 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 thephospholipid layer of thecell membrane, anchoring proteins associated with the cell surface.[57]
Hormonal regulation of bone formation and serum levels
Parathyroid hormone andvitamin 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.[8] 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 bycalcitonin, whose secretion increases with increasing plasma calcium levels.[57]
Abnormal serum levels
Excess intake of calcium may causehypercalcemia. 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.[57]
Chronic hypercalcaemia typically leads tocalcification of soft tissue and its serious consequences: for example, calcification can cause loss of elasticity ofvascular walls and disruption of laminar blood flow—and thence toplaque rupture andthrombosis. Conversely, inadequate calcium or vitamin D intakes may result inhypocalcemia, often caused also by inadequate secretion of parathyroid hormone or defective PTH receptors in cells. Symptoms include neuromuscular excitability, which potentially causestetany and disruption of conductivity in cardiac tissue.[57]
Bone disease
As calcium is required for bone development, many bone diseases can be traced to the organic matrix or thehydroxyapatite in molecular structure or organization of bone.Osteoporosis is a reduction in mineral content of bone per unit volume, and can be treated by supplementation of calcium, vitamin D, andbisphosphonates.[8][9] Inadequate amounts of calcium, vitamin D, or phosphates can lead to softening of bones, calledosteomalacia.[57]
Because calcium reacts exothermically with water and acids, calcium metal coming into contact with bodily moisture results in severe corrosive irritation.[59] When swallowed, calcium metal has the same effect on the mouth, oesophagus, and stomach, and can be fatal.[46] However, long-term exposure is not known to have distinct adverse effects.[59]
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