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Calcium (pronounced /ˈkælsiəm/, KAL-see-əm) is the chemical element with the symbol Ca and atomic number 20. It has an atomic mass of 40.078 amu. Calcium is a soft gray alkaline earth metal, and is the fifth most abundant element by mass in the Earth's crust. Calcium is also the fifth most abundant dissolved ion in seawater by both molarity and mass, after sodium, chloride, magnesium, and sulfate.[1] Calcium is essential for living organisms, particularly in cell physiology, where movement of the calcium ion Ca2+ into and out of the cytoplasm functions as a signal for many cellular processes. As a major material used in mineralization of bones and shells, calcium is the most abundant metal by mass in many animals. Notable characteristics Chemically calcium is reactive and soft for a metal (though harder than lead, it can be cut with a knife with difficulty). It is a silvery metallic element that must be extracted by electrolysis from a fused salt like calcium chloride.[2] Once produced, it rapidly forms a gray-white oxide and nitride coating when exposed to air. It is somewhat difficult to ignite, unlike magnesium, but when lit, the metal burns in air with a brilliant high-intensity red light. Calcium metal reacts with water, evolving hydrogen gas at a rate rapid enough to be noticeable, but not fast enough at room temperature to generate much heat. In powdered form, however, the reaction with water is extremely rapid, as the increased surface area of the powder accelerates the reaction with the water. Part of the slowness of the calcium-water reaction results from the metal being partly protected by insoluble white calcium hydroxide. In water solutions of acids where the salt is water soluble, calcium reacts vigorously. Calcium, with a specific mass of 1.55 g/cm3, is the lightest of the alkali earth metals; magnesium (1.74) and beryllium (1.84) are heavier although they are lighter in atomic mass. From strontium on, the alkali earth metals get heavier along with the atomic mass. Calcium has a higher electrical resistivity than copper or aluminium. Yet, weight for weight, allowing for its much lower density, it is a rather better conductor than either. However, its use in terrestrial applications is usually limited by its high reactivity with air. Calcium salts are colorless from any contribution of the calcium, and ionic solutions of calcium (Ca2+) are colorless as well. Many calcium salts are not soluble in water. When in solution, the calcium ion to the human taste varies remarkably, being reported as mildly salty, sour, "mineral like" or even "soothing." It is apparent that many animals can taste, or develop a taste, for calcium, and use this sense to detect the mineral in salt licks or other sources.[3] In human nutrition, soluble calcium salts may be added to tart juices without much effect to the average palate. Calcium is the fifth most abundant element by mass in the human body, where it is a common cellular ionic messenger with many functions, and serves also as a structural element in bone. It is the relatively high atomic-numbered calcium in the skeleton which causes bone to be radio-opaque. Of the human body's solid components after drying (as for example, after cremation), about a third of the total mass is the approximately one kilogram of calcium which composes the average skeleton (the remainder being mostly phosphorus and oxygen). Visible spectra of many stars, including the Sun, exhibit strong absorption lines of singly-ionized calcium. Prominent among these are the H-line at 3968.5 Å and the K line at 3933.7 Å of singly-ionized calcium, or Ca II. For the Sun and stars with low temperatures, the prominence of the H and K lines can be an indication of strong magnetic activity in the chromosphere. Measurement of periodic variations of these active regions can also be used to deduce the rotation periods of these stars.[4] Calcium, combined with phosphate to form hydroxylapatite, is the mineral portion of human and animal bones and teeth. The mineral portion of some corals can also be transformed into hydroxylapatite. Calcium hydroxide (slaked lime) is used in many chemical refinery processes and is made by heating limestone at high temperature (above 825 °C) and then carefully adding water to it. When lime is mixed with sand, it hardens into a mortar and is turned into plaster by carbon dioxide uptake. Mixed with other compounds, lime forms an important part of Portland cement. Calcium carbonate (CaCO3) is one of the common compounds of calcium. It is heated to form quicklime (CaO), which is then added to water (H2O). This forms another material known as slaked lime (Ca(OH)2), which is an inexpensive base material used throughout the chemical industry. Chalk, marble, and limestone are all forms of calcium carbonate. When water percolates through limestone or other soluble carbonate rocks, it partially dissolves the rock and causes cave formation and characteristic stalactites and stalagmites and also forms hard water. Other important calcium compounds are calcium nitrate, calcium sulfide, calcium chloride, calcium carbide, calcium cyanamide and calcium hypochlorite. Isotopes Calcium has four stable isotopes (40Ca and 42Ca through 44Ca), plus two more isotopes (46Ca and 48Ca) that have such long half-lives that for all practical purposes they can be considered stable. The 20% range in relative mass among naturally-occurring calcium isotopes is greater than for any element except hydrogen and helium. Calcium also has a cosmogenic isotope, radioactive 41Ca, which has a half-life of 103,000 years. Unlike cosmogenic isotopes that are produced in the atmosphere, 41Ca is produced by neutron activation of 40Ca. Most of its production is in the upper metre or so of the soil column, where the cosmogenic neutron flux is still sufficiently strong. 41Ca has received much attention in stellar studies because it decays to 41K, a critical indicator of solar-system anomalies. 97% of naturally occurring calcium is in the form of 40Ca. 40Ca is one of the daughter products of 40K decay, along with 40Ar. While K-Ar dating has been used extensively in the geological sciences, the prevalence of 40Ca in nature has impeded its use in dating. Techniques using mass spectrometry and a double spike isotope dilution have been used for K-Ca age dating. The most abundant isotope, 40Ca, has a nucleus of 20 protons and 20 neutrons. This is the heaviest stable isotope of any element which has equal numbers of protons and neutrons. In supernova explosions, calcium is formed from the reaction of carbon with various numbers of alpha particles (helium nuclei), until the most common calcium isotope (containing 10 helium nuclei) has been synthesized.[citation needed] As with the isotopes of other elements, a variety of processes fractionate, or alter the relative abundance of, calcium isotopes.[5] 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. Isotopically light calcium is preferentially incorporated into minerals, leaving the solution from which the mineral precipitated enriched in isotopically heavy calcium. At room temperature the magnitude of this fractionation is roughly 0.25‰ (0.025%) per atomic mass unit (AMU). Mass-dependant differences in calcium isotope composition conventionally are expressed 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.[6] Calcium isotope fractionation during mineral formation has led to several applications of calcium isotopes. In particular, the 1997 observation by Skulan and DePaolo[7] 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, soft tissue 44Ca/40Ca rises. Soft tissue 44Ca/40Ca falls when bone resorption exceeds bone formation. Because of this relationship, calcium isotopic measurements of urine or blood may be useful in the early detection of metabolic bone diseases like osteoporosis.[8] A similar system exists in the ocean, where seawater 44Ca/40Ca tends to rise when the rate of removal of Ca2+ from seawater by mineral precipitation exceeds the input of new calcium into the ocean, and fall when calcium input exceeds mineral precipitation. It follows that rising 44Ca/40Ca corresponds to falling seawater Ca2+ concentration, and falling 44Ca/40Ca corresponds to rising seawater Ca2+ concentration. 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 probably never is in “steady state” with respect to its calcium input and output.[9][10] This has important climatological implications, as the marine calcium cycle is closely tied to the carbon cycle (see below). Calcium provides an important link between tectonics, climate and the carbon cycle. In the simplest terms, uplift of mountains exposes Ca-bearing rocks to chemical weathering and releases Ca2+ into surface water. This Ca2+ eventually is transported to the ocean where it reacts with dissolved CO2 to form limestone. Some of this limestone settles to the sea floor where it is incorporated into new rocks. Dissolved CO2, along with carbonate and bicarbonate ions, are referred to as dissolved inorganic carbon (DIC). The actual reaction is more complicated and involves the bicarbonate ion (HCO3-) that forms when CO2 reacts with water at seawater pH: Ca2+ + 2HCO3 → CaCO3 (limestone) + CO2 + H2O Note that at ocean pH most of the CO2 produced in this reaction is immediately converted back into HCO3-. The reaction results in a net transport of one molecule of CO2 from the ocean/atmosphere into the lithosphere.[11] 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.[12] Analogous cycles involving magnesium, and to a much smaller extent strontium and barium, have the same effect. As the weathering of limestone (CaCO3) liberates equimolar amounts of Ca2+ and CO2, it has no net effect on the CO2 content of the atmosphere and ocean. The weathering of silicate rocks like granite, on the other hand, is a net CO2 sink because it produces abundant Ca2+ very little CO2. Calcium (Latin word calcis meaning "lime") was known as early as the first century when the Ancient Romans prepared lime as calcium oxide. Literature dating back to 975 AD notes that plaster of paris (calcium sulphate), is useful for setting broken bones. It was not isolated until 1808 in England when Sir Humphry Davy electrolyzed a mixture of lime and mercuric oxide. Davy was trying to isolate calcium; when he heard that Swedish chemist Jöns Jakob Berzelius and Pontin prepared calcium amalgam by electrolyzing lime in mercury, he tried it himself. He worked with electrolysis throughout his life and also discovered/isolated sodium, potassium, magnesium, boron and barium. Calcium metal was not available in large scale until the beginning of the 20th century. Occurrence Calcium is not naturally found in its elemental state. Calcium occurs most commonly in sedimentary rocks in the minerals calcite, dolomite and gypsum. It also occurs in igneous and metamorphic rocks chiefly in the silicate minerals: plagioclase, amphiboles, pyroxenes and garnets Calcium is used[13] * as a reducing agent in the extraction of other metals, such as uranium, zirconium, and thorium. This article does not cite any references or sources. * Calcium carbonate (CaCO3) is used in manufacturing cement and mortar, lime, limestone (usually used in the steel industry) and aids in production in the glass industry. It also has chemical and optical uses as mineral specimens in toothpastes, for example. Nutrition Calcium is an important component of a healthy diet and a mineral necessary for life. The National Osteoporosis Foundation says, "Calcium plays an important role in building stronger, denser bones early in life and keeping bones strong and healthy later in life." Approximately ninety-nine percent of the body's calcium is stored in the bones and teeth.[15] The rest of the calcium in the body has other important uses, such as some exocytosis, especially neurotransmitter release, and muscle contraction. In the electrical conduction system of the heart, calcium replaces sodium as the mineral that depolarizes the cell, proliferating the action potential. In cardiac muscle, sodium influx commences an action potential, but during potassium efflux, the cardiac myocyte experiences calcium influx, prolonging the action potential and creating a plateau phase of dynamic equilibrium. Long-term calcium deficiency can lead to rickets and poor blood clotting and in case of a menopausal woman, it can lead to osteoporosis, in which the bone deteriorates and there is an increased risk of fractures. While a lifelong deficit can affect bone and tooth formation, over-retention can cause hypercalcemia (elevated levels of calcium in the blood), impaired kidney function and decreased absorption of other minerals.[16] High calcium intakes or high calcium absorption were previously thought to contribute to the development of kidney stones. However, a high calcium intake has been associated with a lower risk for kidney stones in more recent research.[17][18][19] Vitamin D is needed to absorb calcium. Dairy products, such as milk and cheese, are a well-known source of calcium. Some individuals are allergic to dairy products and even more people, particularly those of non Indo-European descent, are lactose-intolerant, leaving them unable to consume non-fermented dairy products in quantities larger than about half a liter per serving. Others, such as vegans, avoid dairy products for ethical and health reasons. Fortunately, many good sources of calcium exist. These include seaweeds such as kelp, wakame and hijiki; nuts and seeds (like almonds and sesame); blackstrap molasses; beans; oranges; figs; quinoa; amaranth; collard greens; okra; rutabaga; broccoli; dandelion leaves; kale; and fortified products such as orange juice and soy milk. Research has found an association between diets high in animal protein and increased urinary calcium loss from the bones.[20][21][22] A diet high in fruit, vegetables and cereals was demonstrated to result in greater femoral bone mineral density in older men, in comparison to a range of other diets. Diets high in candy were found to result in lower bone density in both men and women.[23]. An overlooked source of calcium is eggshell, which can be ground into a powder and mixed into food or a glass of water.[24][25][26] Cultivated vegetables generally have less calcium than wild plants.[27] The calcium content of most foods can be found in the USDA National Nutrient Database.[28] Calcium supplements are used to prevent and to treat calcium deficiencies. Most experts recommend that supplements be taken with food and that no more than 600 mg should be taken at a time because the percent of calcium absorbed decreases as the amount of calcium in the supplement increases.[14] It is recommended to spread doses throughout the day. Recommended daily calcium intake for adults ranges from 1000 to 1500 mg. It is recommended to take supplements with food to aid in absorption. Vitamin D is added to some calcium supplements. Proper vitamin D status is important because vitamin D is converted to a hormone in the body which then induces the synthesis of intestinal proteins responsible for calcium absorption.[29] * The absorption of calcium from most food and commonly-used dietary supplements is very similar.[30] This is contrary to what many calcium supplement manufacturers claim in their promotional materials. In July 2006, a report citing research from Fred Hutchinson Cancer Research Center in Seattle, Washington claimed that women in their 50s gained 5 pounds less in a period of 10 years by taking more than 500 mg of calcium supplements than those who did not. However, the doctor in charge of the study, Dr. Alejandro J. Gonzalez also noted it would be "going out on a limb" to suggest calcium supplements as a weight-limiting aid.[37] Such studies often do not test calcium alone, but rather combinations of calcium and vitamin D. Randomized controlled trials found both positive[38][39] and negative[40][41][42][43] effects. The different results may be explained by doses of calcium and underlying rates of calcium supplementation in the control groups.[44] However, it is clear that increasing the intake of calcium promotes deposition of calcium in the bones, where it is of more benefit in preventing the compression fractures resulting from the osteoporotic thinning of the dendritic web of the bodies of the vertebrae, than it is at preventing the more serious cortical bone fractures which happen at hip and wrist.[citation needed] A meta-analysis[39] by the international Cochrane Collaboration of two randomized controlled trials[45][46] found that calcium "might contribute to a moderate degree to the prevention of adenomatous colonic polyps". More recent studies were conflicting, and one which was positive for effect (Lappe, et al.) did control for a possible anti-carcinogenic effect of vitamin D, which was found to be an independent positive influence from calcium-alone on cancer risk (see second study below) [47]. * A randomized controlled trial found that 1000 mg of elemental calcium and 400 IU of vitamin D3 had no effect on colorectal cancer[48] Hazards and Toxicity Compared to other metals, the calcium ion and most calcium compounds have low toxicity. This is not surprising given the very high natural abundance of calcium compounds in the environment and in organisms. Calcium poses few, if any, serious environmental problems. Acute calcium poisoning is rare, and difficult to achieve unless calcium compounds are administered intravenously. For example, the oral LD50 for rats for calcium carbonate and calcium chloride are 6450 [51] and 1400 mg/kg,[52] respectively. Calcium metal is hazardous because of its sometimes violent reactions with water and acids. 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 (e.g., hair) that block drains. When swallowed calcium metal has the same effect on the mouth, esophagus and stomach, and can be fatal.[53] Excessive consumption of calcium carbonate antacids/dietary supplements (such as Tums) over a period of weeks or months can cause milk-alkali syndrome, with symptoms ranging from hypercalcemia to potentially fatal renal failure. What constitutes “excessive” consumption is not well known and probably varies a great deal from person to person. Persons who consume more than 10 grams/day of CaCO3 (=4 g Ca) are at risk of developing milk-alkali syndrome,[54] but the condition has been reported in at least one person consuming only 2.5 grams/day of CaCO3 (=1 g Ca), an amount usually considered moderate and safe.[55] Calcium: the molecular basis of calcium action in biology and medicine, Roland Pochet Calcium in human health, Connie Weaver, Robert Proulx Heaney Calcium signalling and disease: molecular pathology of calcium, Ernesto Carafoli, Marisa Brini Calcium phosphates in biological and industrial systems, Zahid Amjad See also * Calcium metabolism References 1. ^ Dickson, A. G. and Goyet, C. (1994). "5". Handbook of method for the analysis of the various parameters of the carbon dioxide system in sea water, version 2. ORNL/CDIAC-74. http://cdiac.esd.ornl.gov/ftp/cdiac74/chapter5.pdf. Further reading * Rebecca J. Donatelle. Health, The Basics. 6th ed. San Francisco: Pearson Education, Inc. 2005. External links * WebElements.com — Calcium
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