Calcium

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Calcium is a chemical element with symbol Ca and atomic number 20. An alkaline earth metal, calcium is a reactive pale yellow metal that forms a dark oxide-nitride layer when exposed to air. Its physical and chemical properties are most similar to its heavier homologues strontium and barium. It is the fifth most abundant element in Earth's crust and the third most abundant metal, after iron and aluminium. The most common calcium compound on Earth is calcium carbonate, found in limestone and the fossilised remnants of early sea life; gypsum, anhydrite, fluorite, and apatite are also sources of calcium.

The name derives from Latin calx "lime", which was obtained from heating limestone. Its compounds were known to the ancients, though their chemistry was unknown until the seventeenth century. It was isolated by Humphry Davy in 1808 via electrolysis of its oxide, who named the element. While the pure metal does not have many applications due to its high reactivity, it is often used as an alloying component in small quantities in steelmaking, and calcium-lead alloys are sometimes used in automotive batteries. Calcium compounds on the other hand are very widely used in many industries: for example, they are used in foods and pharmaceuticals for calcium supplementation, in the paper industry as bleaches, in cement, in the manufacture of soaps, and as electrical insulators.

Calcium is the fifth most abundant element in the human body and the most abundant metal. Calcium ions play a vital role in the physiology and biochemistry of organisms and the cell as electrolytes. They play an important role in signal transduction pathways, where they act as a second messenger, in neurotransmitter release from neurons, in contraction of all muscle cell types, and in fertilization. Many enzymes require calcium ions as a cofactor. Calcium ions outside cells are also important for maintaining the potential difference across excitable cell membranes, as well as proper bone formation.

Video Calcium

Characteristics

Classification

Calcium is a very ductile silvery metal with a pale yellow tint whose properties are very similar to the heavier elements in its group, strontium, barium, and radium. A calcium atom has twenty electrons, arranged in the electron configuration [Ar]4s². Like the other elements placed in group 2 of the periodic table, calcium has two valence 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 a noble gas, in this case argon. Hence, calcium is almost always divalent in its compounds, which are usually ionic. Hypothetical univalent salts of calcium would be stable with respect to their elements, but not to disproportionation to the divalent salts and calcium metal, because the enthalpy of formation of MX₂ is much higher than those of the hypothetical MX. This occurs because of the much greater lattice energy afforded by the more highly charged Ca²⁺ cation compared to the hypothetical Ca⁺ cation.

Calcium is considered to be an alkaline earth metal, along with these heavier elements and the lighter beryllium and magnesium. Nevertheless, there are significant differences in chemical and physical properties between beryllium and magnesium (which behave more like aluminium and zinc respectively and have some of the weaker metallic character of the post-transition metals) and the group members from calcium onwards, which traditionally led to "alkaline earth metal" only applying to the latter group. 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, higher than its adjacent group 2 metals do. It crystallises in the face-centered cubic arrangement like strontium; above 450 °C, it changes to an anisotropic hexagonal close-packed arrangement like magnesium. The density of 1.55 g·cm^-3 is the lowest in its group, with others decreasing towards it. Calcium can be cut with a knife with effort, although it is still harder than lead. While calcium is a poorer conductor of electricity than copper or aluminium by volume, it is a better conductor than both of them by mass due to its very low density. Although it is infeasible for terrestrial applications as it reacts quickly with atmospheric oxygen, its use as a conductor in space has been considered.

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 calcium hydroxide and hydrogen gas. It also reacts with the oxygen and nitrogen in the air to form a mixture of calcium oxide and calcium nitride. 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% relative humidity it may be stored indefinitely at room temperature.

Besides the simple oxide CaO, the peroxide CaO₂ can be made by direct oxidation of calcium metal under a high pressure of oxygen, and there is some evidence for a yellow superoxide 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. Calcium carbonate (CaCO₃) and calcium sulfate (CaSO₄) are particularly abundant minerals. Like strontium and barium, as well as the alkali metals and the divalent lanthanides europium and ytterbium, calcium metal dissolves directly in liquid ammonia 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 intermetallic compounds such as CaZn₁₃. Calcium is readily complexed by oxygen chelates such as EDTA and polyphosphates, which are useful in analytic chemistry and removing calcium ions from hard water. In the absence of steric hindrance, smaller group 2 cations tend to form stronger complexes, but when large polydentate macrocycles are involved the trend is reversed.

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 catalysts. Organocalcium compounds tend to be more similar to organoytterbium compounds due to the similar ionic radii 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 dicyclopentadienyl, Ca(C₅H₅)₂, must be made by directly reacting calcium metal with mercurocene or cyclopentadiene 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 isotopes (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 (⁴⁸Ca, 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 silicon-burning process from fusion of alpha particles and is the heaviest stable nuclide with equal proton and neutron numbers; its occurrence is also supplemented slowly by the decay of primordial ⁴⁰K. Adding another alpha particle would lead to unstable ⁴⁴Ti, which quickly decays via two successive electron captures 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 oxygen-burning and silicon-burning processes, leaving the two heavier ones to be produced via neutron-capturing processes. ⁴⁶Ca is mostly produced in a "hot" s-process, 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 r-process in type Ia supernovae, 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 doubly magic nucleus, having 20 protons and 28 neutrons arranged in closed shells. Its beta decay to ⁴⁸Sc is very hindered because of the gross mismatch of nuclear spin: ⁴⁸Ca has zero nuclear spin, being even-even, while ⁴⁸Sc has spin 6+, so the decay is forbidden by the conservation of angular momentum. 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 double beta decay to ⁴⁸Ti 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 double electron capture to ⁴⁰Ar, 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 radioisotope of calcium is ⁴¹Ca. It decays by electron capture to stable ⁴¹K with a half-life of about a hundred thousand years. Its existence in the early Solar System as an extinct radionuclide has been inferred from excesses of ⁴¹K: traces of ⁴¹Ca also still exist today, as it is a cosmogenic nuclide, continuously reformed through neutron activation of natural ⁴⁰Ca. Many other calcium radioisotopes are known, ranging from ³⁴Ca to ⁵⁷Ca: 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 beta plus decay to isotopes of potassium, and those heavier than ⁴⁴Ca usually undergo beta minus decay to isotopes of scandium, although near the nuclear drip lines proton emission and neutron emission 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 fractionation of calcium isotopes that accompanies the precipitation of calcium minerals such as calcite, aragonite and 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 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.

Maps Calcium

History

Calcium compounds were known for millennia, although their chemical makeup was not understood until the 17th century. Lime as a building material and as plaster for statues was used as far back as around 7000 BC. The first dated lime kiln dates back to 2500 BC and was found in Khafajah, Mesopotamia. At about the same time, dehydrated gypsum (CaSO₄·2H₂O) was being used in the Great Pyramid of Giza; this material would later be used for the plaster in the tomb of Tutankhamun. The climate of present-day Italy being warmer than that of Egypt, the ancient Romans instead used lime mortars made by heating limestone (CaCO₃); the name "calcium" itself derives from the Latin word calx "lime". 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 of carbon dioxide, which as a gas had not been recognised by the ancient Romans.

In 1787, Antoine Lavoisier suspected that lime might be an oxide of a fundamental 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 speculated:

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.

Calcium, along with its congeners magnesium, strontium, and barium, was first isolated by Humphry Davy in 1808. Following the work of Jöns Jakob Berzelius and Magnus Martin af Pontin on electrolysis, Davy isolated calcium and magnesium by putting a mixture of the respective metal oxides with mercury(II) oxide on a 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. 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.

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Occurrence and production

At 3%, calcium is the fifth most abundant element in the Earth's crust, and the third most abundant metal behind aluminium and iron. It is also the fourth most abundant element in the lunar highlands. Sedimentary calcium carbonate deposits pervade the Earth's surface as fossilised remains of past marine life; they occur in two forms, the rhombohedral calcite (more common) and the orthorhombic aragonite (forming in more temperate seas). Minerals of the first type include limestone, dolomite, marble, chalk, and iceland spar; aragonite beds make up the Bahamas, the Florida Keys, and the Red Sea basins. Corals, sea shells, and pearls are mostly made up of calcium carbonate. Among the other important minerals of calcium are gypsum (CaSO₄·2H₂O), anhydrite (CaSO₄), fluorite (CaF₂), and apatite ([Ca₅(PO₄)₃F]).

The major producers of calcium are China (about 10000 to 12000 tonnes per year), Russia (about 6000 to 8000 tonnes per year), and the United States (about 2000 to 4000 tonnes per year). Canada and 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. In Russia and China, Davy's method of electrolysis is still used, but is instead applied to molten calcium chloride. Since calcium is less reactive than strontium or barium, the oxide-nitride coating that results in air is stable and lathe machining and other standard metallurgical techniques are suitable for calcium. In the United States and Canada, calcium is instead produced by reducing lime with aluminium metal at high temperatures.

Geochemical cycling

Calcium provides a link between tectonics, climate, and the carbon cycle. 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 limestone, which in turn settles to the sea floor where it is incorporated into new rocks. Dissolved CO₂, along with carbonate and bicarbonate ions, are termed "dissolved inorganic carbon" (DIC).

The actual reaction is more complicated and involves the bicarbonate ion (HCO-
3) that forms when CO₂ reacts with water at seawater pH:

Ca2+
+ 2HCO-
3 -> CaCO
3 (limestone) + CO
2 + H
2O

At seawater pH, most of the CO₂ is immediately converted back into HCO-
3. The reaction results in a net transport of one molecule of CO₂ from the ocean/atmosphere into the lithosphere. 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.

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Uses

The largest use of calcium is in steelmaking, due to its strong chemical affinity for oxygen and sulfur. Its oxides and sulfides, once formed, give liquid lime aluminate 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 automotive batteries, in which the use of 0.1% calcium-lead alloys instead of the usual antimony-lead alloys leads to lower water loss and lower self-discharging. 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. Calcium is also used to strengthen aluminium alloys used for bearings, for the control of graphitic carbon in cast iron, and to remove bismuth impurities from lead. Calcium metal is found in some drain cleaners, where it functions to generate heat and calcium hydroxide that saponifies the fats and liquefies the proteins (for example, those in hair) that block drains. Besides metallurgy, the reactivity of calcium is exploited to remove nitrogen from high-purity argon gas and as a getter for oxygen and nitrogen. It is also used as a reducing agent in the production of chromium, zirconium, thorium, and uranium. It can also be used to store hydrogen gas, as it reacts with hydrogen to form solid calcium hydride, 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 osteoporosis. 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 carbon cycle.

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 calcium lactate, calcium diphosphate, and tricalcium phosphate. The last is also used as a polishing agent in toothpaste and in antacids. Calcium lactobionate is a white powder that is used as a suspending agent for pharmaceuticals. In baking, calcium monophosphate is used as a leavening agent. Calcium sulfite is used as a bleach in papermaking and as a disinfectant, calcium silicate is used as a reinforcing agent in rubber, and calcium acetate is a component of liming rosin and is used to make metallic soaps and synthetic resins.

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Biological and pathological role

Calcium is an essential element needed in large quantities. The Ca²⁺ ion acts as an 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 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 coordination complexes with many organic compounds, especially proteins; it also forms compounds with a wide range of solubilities, enabling the formation of skeletons.

Calcium ions may be complexed by proteins through binding the carboxyl groups of glutamic acid or aspartic acid residues; through interacting with phosphorylated serine, tyrosine, or threonine residues; or by being chelated by ?-carboxylated amino acid residues. Trypsin, a digestive enzyme, uses the first method; osteocalcin, a bone matrix protein, uses the third. Some other bone matrix proteins such as osteopontin and 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 phospholipid layer of the cell membrane, anchoring proteins associated with the cell surface. As an example of the wide range of solubility of calcium compounds, monocalcium phosphate is very soluble in water, 85% of extracellular calcium is as dicalcium phosphate with a solubility of 2.0 mM and the hydroxyapatite of bones in an organic matrix is tricalcium phosphate at 100 µM.

About three-quarters of dietary calcium is from dairy products and grains, the rest being accounted for by vegetables, protein-rich foods, fruits, sugar, fats, and oil. Calcium supplementation is controversial, as the bioavailability of calcium is strongly dependent on the solubility of the salt involved: calcium citrate, malate, and lactate are highly bioavailable while the oxalate is much less so. The intestine absorbs about one-third of calcium eaten as the free ion, and plasma calcium level is then regulated by the kidneys. Parathyroid hormone and 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. 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 calcitonin, whose secretion increases with increasing plasma calcium levels.

Excess intake of calcium may cause hypercalcaemia, but because of the inefficient absorption of calcium by the intestines a more likely cause is excessive vitamin D intake or excessive secretion of parathyroid hormone. It can also occur due to the bone destruction that occurs when tumours metastasise to bone. This results in deposition of calcium salts into the heart, the blood vessels, and the kidneys. Symptoms include anorexia, nausea, vomiting, memory loss, confusion, muscle weakness, increased urination, dehydration, and metabolic bone disease. Chronic hypercalcaemia may lead to soft tissue calcification, which can lead to serious consequences: for example, calcification of the vascular wall can lead to a loss of elasticity and the disruption of laminar blood flow, and thence to plaque rupture and thrombosis. Likewise, inadequate calcium or vitamin D intake results in hypocalcaemia, often caused by inadequate secretion of parathyroid hormone or defective receptors to it in cells. Symptoms include neuromuscular excitability, potentially causing tetany and defects in cardiac conduction.

As calcium is heavily involved in bone manufacture, many bone diseases can be traced to problems with the organic matrix or the hydroxyapatite in molecular structure or organisation. For example, osteoporosis is a reduction in mineral content of bone per unit volume, and can be treated by supplementation of calcium, vitamin D, and biphosphates. Calcium supplements may benefit the serum lipids in women who have passed menopause as well as older men; in post-menopausal women calcium supplementation also appears to be inversely correlated with cardiovascular disease. Inadequate amounts of calcium, vitamin D, or phosphates can lead to the softening of bones, known as osteomalacia.

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Safety

Because calcium reacts exothermically with water and acids, contact of calcium metal 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. However, long-term exposure is not known to have distinct adverse effects.

Because of concerns of long-term adverse side effects such as calcification of arteries and kidney stones, the U.S. Institute of Medicine (IOM) and the European Food Safety Authority (EFSA) both set Tolerable Upper Intake Levels (ULs) for the combination of dietary and supplemental calcium. From the IOM, people ages 9-18 years are not supposed to exceed 3,000 mg/day; for ages 19-50 not to exceed 2,500 mg/day; for ages 51 and older, not to exceed 2,000 mg/day. The EFSA set UL at 2,500 mg/day for adults but decided the information for children and adolescents was not sufficient to determine ULs.

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See also

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References

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Bibliography

• Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0-08-037941-9. 
• Hluchan, Stephen E.; Pomerantz, Kenneth (2005), "Calcium and Calcium Alloys", Ullmann's Encyclopedia of Industrial Chemistry, Weinheim: Wiley-VCH, doi:10.1002/14356007.a04_515.pub2 

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External links

• WebElements.com - Calcium
• Calcium at The Periodic Table of Videos (University of Nottingham)
• Nutrition fact sheet from the National Institutes of Health

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