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Calcium is a chemical element; it has symbol Ca and atomic number 20. As an 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 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 fossils of early sea life; gypsum, anhydrite, fluorite, and apatite are also sources of calcium. The name comes from Latin calx "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 electrolysis of its oxide by 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 human body. As , calcium ions (Ca2+) play a vital role in the physiology and biochemistry processes of organisms and cells: in signal transduction pathways where they act as a second messenger; in neurotransmitter release from neurons; in contraction of all muscle cell types; as cofactors in many ; and in fertilization. Calcium ions outside cells are important for maintaining the potential difference across excitable , protein synthesis, and bone formation.
Hence, calcium is almost always divalent in its compounds, which are usually ionic compound. 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 MX is much higher than those of the hypothetical MX. This occurs because of the much greater lattice energy afforded by the more highly charged Ca cation compared to the hypothetical Ca cation.
Calcium, strontium, barium, and radium are always considered to be alkaline earth metals; the lighter beryllium and 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 behavior: they behave more like aluminium and zinc respectively and have some of the weaker metallic character of the post-transition metals, which is why the traditional definition of the term "alkaline earth metal" excludes them.
Calcium is harder than lead but can be cut with a knife with effort. While calcium is a poorer conductor of electricity than copper or aluminium by volume, it is a better conductor by mass than both due to its very low density. While calcium is infeasible as a conductor for most terrestrial applications as it reacts quickly with atmospheric oxygen, its use as such in space has been considered.
Besides the simple oxide CaO, calcium peroxide, CaO, 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(O).Calcium hydroxide, Ca(OH), is a strong base, though not as strong as the hydroxides of strontium, barium or the alkali metals. All four dihalides of calcium are known. Calcium carbonate (CaCO) and calcium sulfate (CaSO) are particularly abundant minerals. Like strontium and barium, as well as the alkali metals and the divalent europium and ytterbium, calcium metal dissolves directly in liquid ammonia to give a dark blue solution.
Due to the large size of the calcium ion (Ca), high coordination numbers are common, up to 24 in some intermetallic compounds such as CaZn. Calcium is readily complexed by oxygen such as EDTA and , 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 are involved the trend is reversed.
A dominant theme in molecular organocalcium chemistry is the large radius of calcium, which often leads to high coordination numbers. For example, dimethylcalcium appears to be a 3-dimensional polymer, whereas dimethylmagnesium is a linear polymer with tetrahedral Mg centers. are often required to disfavor polymeric species. For example, calcium dicyclopentadienyl, has a polymeric structure and thus is nonvolatile and insoluble in solvents. Replacing the ligand with the bulkier (pentamethylcyclopentadienyl) gives a soluble complex that sublimes and forms well-defined adducts with ethers. Organocalcium compounds tend to be more similar to organoytterbium compounds due to the similar ionic radii of Yb (102 pm) and Ca (100 pm).
Organocalcium compounds have been well investigated. Some such complexes exhibit catalytic properties, although none have been commercialized.
By far the most common isotope is Ca, which makes up 96.941% of natural calcium. It is produced in the silicon-burning process from fusion of and is the heaviest stable nuclide with equal proton and neutron numbers; its occurrence is also supplemented slowly by the decay of primordial K. Adding another alpha particle leads to unstable Ti, which decays via two successive to stable Ca; this makes up 2.806% of natural calcium and is the second-most common isotope.
The other four natural isotopes, Ca, are significantly rarer, each comprising less than 1% of natural calcium. The four lighter isotopes are mainly products of oxygen-burning and silicon-burning, leaving the two heavier ones to be produced via neutron capture. Ca is mostly produced in a "hot" s-process, as its formation requires a rather high neutron flux to allow short-lived Ca to capture a neutron. Ca is produced by electron capture in the r-process in type Ia supernovae, where high neutron excess and low enough entropy ensures its survival.
Ca and Ca are the first "classically stable" nuclides with a 6-neutron or 8-neutron excess respectively. Though extremely neutron-rich for such a light element, Ca is very stable because it is a doubly magic nucleus, with 20 protons and 28 neutrons arranged in closed shells. Its beta decay to scandium is very hindered by the gross mismatch of nuclear spin: Ca has zero nuclear spin, being even–even, while Sc has spin 6+, so the decay is forbidden by conservation of angular momentum. While two excited states of Sc are available for decay as well, they are also forbidden due to their high spins. As a result, when Ca does decay, it does so by double beta decay to titanium instead, being the lightest nuclide known to undergo double beta decay.
Ca can also theoretically double-beta-decay to Ti, but this has never been observed. The most common isotope Ca is also doubly magic and could undergo double electron capture to argon, 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 of Ca and Ca are 5.9 × 10 years and 2.8 × 10 years respectively.
Excluding Ca, the longest lived radioisotope of calcium is Ca. It decays by electron capture to stable potassium with a half-life of about 10 years. Its existence in the early Solar System as an extinct radionuclide has been inferred from excesses of K. Traces of Ca also still exist today, as it is a cosmogenic nuclide, continuously produced through neutron activation of natural Ca.
Many other calcium radioisotopes are known, ranging from Ca to Ca. They are all much shorter-lived than Ca; the most stable are Ca (half-life 163 days) and Ca (half-life 4.54 days). Isotopes lighter than Ca usually undergo beta plus decay to isotopes of potassium, and those heavier than Ca usually undergo beta minus decay to scandium; though 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 Ca/Ca) in a sample compared to the same ratio in a standard reference material. Ca/Ca varies by about 1–2‰ among organisms on Earth.
About the same time, dehydrated gypsum (CaSO·2HO) was being used in the Great Pyramid of Giza. This material would later be used for the plaster in the tomb of Tutankhamun. The instead used lime mortars made by heating limestone (CaCO). 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 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), and silice (silica, silicon dioxide)). About these "elements", Lavoisier reasoned:
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 of 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.
The major producers of calcium are China (about 10000 to 12000 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 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 U.S. and Canada, calcium is instead produced by reducing lime with aluminium at high temperatures. In this process, powdered high-calcium lime and powdered aluminum are mixed and compacted into for a high degree of contact, which are then placed in a sealed retort which has been Vacuum and heated to ~1200°C. 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 of calcium aluminate. High-purity calcium can be obtained by Distillation low-purity calcium at high temperatures.
The actual reaction is more complicated and involves the bicarbonate ion (HCO) that forms when CO reacts with water at seawater pH:
At seawater pH, most of the dissolved CO is immediately converted back into . The reaction results in a net transport of one molecule of CO from the ocean/atmosphere into the lithosphere. The result is that each Ca ion released by chemical weathering ultimately removes one CO 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 CO from the ocean and air, exerting a strong long-term effect on climate.
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 Saponification 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, vanadium 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 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.
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 Ca/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 osteoporosis.
A similar system exists in seawater, where Ca/Ca tends to rise when the rate of removal of Ca 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 Ca/Ca over geologic time, along with a theoretical explanation of these changes. More recent papers have confirmed this observation, demonstrating that seawater Ca 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 . Calcium lactobionate is a white powder that is used as a suspending agent for pharmaceuticals. In baking, calcium phosphate 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.
Calcium is on the World Health Organization's List of Essential Medicines.
Because of concerns for long-term adverse side effects, including calcification of arteries and , 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 g/day combined intake; for ages 19–50, not to exceed 2.5 g/day; for ages 51 and older, not to exceed 2 g/day. EFSA set the UL for all adults at 2.5 g/day, but decided the information for children and adolescents was not sufficient to determine ULs.
| + Age-adjusted daily calcium recommendations (from U.S. Institute of Medicine RDAs) | |
| 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).
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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.
Chronic hypercalcaemia typically leads to calcification of soft tissue and its serious consequences: for example, calcification can cause loss of elasticity of and disruption of laminar blood flow—and thence to plaque rupture and thrombosis. Conversely, inadequate calcium or vitamin D intakes may result in hypocalcemia, often caused also by inadequate secretion of parathyroid hormone or defective PTH receptors in cells. Symptoms include neuromuscular excitability, which potentially causes tetany and disruption of conductivity in cardiac tissue.
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