Calcium

This article is about the metallic element. For the place, see Calcium, New York.

Calcium,  20Ca

Spectral lines of calcium
General properties
Name, symbol calcium, Ca
Pronunciation /ˈkælsiəm/
KAL-see-əm
Appearance dull gray, silver; with a pale yellow tint[1]
Calcium in the periodic table
Mg
↑
Ca
↓
Sr
potassium ← calcium → scandium
Atomic number (Z) 20
Group, block group 2 (alkaline earth metals), s-block
Period period 4
Element category   alkaline earth metal
Standard atomic weight (±) (Ar) 40.078(4)[2]
Electron configuration [Ar] 4s2
per shell
 2, 8, 8, 2
Physical properties
Phase solid
Melting point 1115 K ​(842 °C, ​1548 °F)
Boiling point 1757 K ​(1484 °C, ​2703 °F)
Density near r.t. 1.55 g/cm3
when liquid, at m.p. 1.378 g/cm3
Heat of fusion 8.54 kJ/mol
Heat of vaporization 154.7 kJ/mol
Molar heat capacity 25.929 J/(mol·K)

vapor pressure

P (Pa) 1 10 100 1 k 10 k 100 k
at T (K) 864 956 1071 1227 1443 1755
Atomic properties
Oxidation states +2, +1[3] ​(a strongly basic oxide)
Electronegativity Pauling scale: 1.00
Ionization energies 1st: 589.8 kJ/mol
2nd: 1145.4 kJ/mol
3rd: 4912.4 kJ/mol
(more)
Atomic radius empirical: 197 pm
Covalent radius 176±10 pm
Van der Waals radius 231 pm
Miscellanea
Crystal structure

​face-centered cubic (fcc)

Face-centered cubic crystal structure for calcium
Speed of sound thin rod 3810 m/s (at 20 °C)
Thermal expansion 22.3 µm/(m·K) (at 25 °C)
Thermal conductivity 201 W/(m·K)
Electrical resistivity 33.6 nΩ·m (at 20 °C)
Magnetic ordering diamagnetic
Young's modulus 20 GPa
Shear modulus 7.4 GPa
Bulk modulus 17 GPa
Poisson ratio 0.31
Mohs hardness 1.75
Brinell hardness 170–416 MPa
CAS Number 7440-70-2
History
Discovery and first isolation Humphry Davy (1808)
Most stable isotopes of calcium
iso NA half-life DM DE (MeV) DP
40Ca 96.941% is stable with 20 neutrons
41Ca trace 1.03×105 y ε – 41K
42Ca 0.647% is stable with 22 neutrons
43Ca 0.135% is stable with 23 neutrons
44Ca 2.086% is stable with 24 neutrons
45Ca syn 162.7 d β− 0.258 45Sc
46Ca 0.004% is stable with 26 neutrons
47Ca syn 4.536 d β− 0.694, 1.99 47Sc
γ 1.297 –
48Ca 0.187% 4.3×1019 y β−β− 4.274 48Ti

Calcium is a chemical element with symbol Ca and atomic number 20. Calcium is a soft gray Group 2 alkaline earth metal, fifth-most-abundant element by mass in the Earth's crust. The ion Ca2+ is also the fifth-most-abundant dissolved ion in seawater by both molarity and mass, after sodium, chloride, magnesium, and sulfate.[4] Free calcium metal is too reactive to occur in nature. Calcium is produced in supernova nucleosynthesis.

Calcium is essential for living organisms, particularly in cell physiology where movement of the calcium ion into and out of the cytoplasm functions as a signal for many cellular processes. As a major material used in mineralization of bone, teeth and shells, calcium is the most abundant metal by mass in many animals.

Uses

Calcium metal find use as a deoxidizer, desulfurizer, and decarbonizer for production of some ferrous and nonferrous alloys. In steelmaking and production of iron, Ca reacts with oxygen and sulfur-containing impurities.[5] Almost all applications are associated with calcium compounds and salts.[6]

Food

In solution, the calcium ion varies remarkably to the human taste, 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.[8] In human nutrition, soluble calcium salts may be added to tart juices without much effect to the average palate.

Recommended adequate intake by the IOM for calcium:[9][10]
Age Calcium (mg/day)
0–6 months 200
7–12 months 260
1–3 years 700
4–8 years 1000
9–18 years 1300
19–50 years 1000
51–70 years (male) 1000
51–70 years (female) 1200
71+ years 1200

Calcium is an important component of a healthy diet and a mineral necessary for life. The National Osteoporosis Foundation states, "Calcium plays an important role in building stronger, denser bones early in life and keeping bones strong and healthy later in life." Approximately 99 percent of the calcium in the human body is in the bones and teeth.[11] The rest of the calcium in the body has other important uses, such as some exocytosis, especially neurotransmitter release, and muscle contraction. Intracellular calcium overload may lead some kind of cells to oxidative stress and apoptosis, and produces several diseases.[12] 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; in menopausal women, deficiency can lead to osteoporosis, a condition in which the bone deteriorates and fractures more readily. 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.[13][14] Several sources suggest a correlation between high calcium intake (2000 mg per day, or twice the U.S. recommended daily allowance, equivalent to six or more glasses of milk per day) and prostate cancer.[15] 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, in particular 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 ethical and/or health vegans, voluntarily avoid dairy products.

Many good vegetable sources of calcium exist:

Numerous vegetables, notably spinach, chard, and rhubarb have a high calcium content, but they may also contain varying amounts of oxalic acid that binds calcium and reduces its absorption. The same problem may affect the absorption of calcium from amaranth, collard greens, and chicory greens. This process may also be related to the generation of calcium oxalate.

An overlooked source of calcium is eggshell, which can be ground into a powder and mixed into food or a glass of water.[17][18][19]

The calcium content of most foods can be found in the USDA National Nutrient Database.[20]

Dietary supplements

500 milligram calcium supplements made from calcium carbonate

Calcium supplements are used to prevent and to treat calcium deficiencies. However, it has been found that the taking of calcium supplements by people with a history of stroke or with white matter lesions greatly increased their chances of developing dementia.[21][22] The Office of Dietary Supplements (National Institutes of Health) recommends that no more than 600 mg of supplement should be taken at a time because the percent of calcium absorbed decreases as the amount of calcium in the supplement increases.[9] It is therefore recommended to spread doses throughout the day.[23] Recommended daily calcium intake for adults ranges from 1000 to 1300 mg.[23] Calcium supplements may have side effects such as bloating and constipation in some people. It is suggested that taking the supplements with food may aid in nullifying these side effects.[23]

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.[24]

Bone health

Dietary calcium supplement is generally not necessary for maintaining bone mineral density, and carries risks that outweigh any benefits.[31] Calcium intake is not significantly associated with hip fracture risk in either men nor women.[32] The U.S. Preventive Service Task Force therefore recommends against a daily supplement of calcium or Vitamin D.[33]

Cardiovascular impact

A study investigating the effects of personal calcium supplement on cardiovascular risk in the Women’s Health Initiative Calcium/Vitamin D Supplementation Study (WHI CaD Study) found a modestly increased risk of cardiovascular events, particularly myocardial infarction in postmenopausal women. A broad recommendation of calcium/vitamin D supplements is therefore not warranted.[34] In contrast, the authors of a 2013 literature review concluded that the benefits of calcium supplementation in some studies, such as on bone health, appear to outweigh any risk calcium supplementation may pose to the cardiovascular health.[35]

Cancer

Overall, there is no clear evidence of the effect of calcium supplements in cancer prevention: some studies suggest it might decrease the risk, but others suggest it might increase the risk. Consequently, the National Cancer Institute, part of the National Institutes of Health, does not recommend the use of calcium supplements for this purpose.[36]

There is weak evidence calcium supplementation might have a preventative effect against developing colorectal adenomatous polyps, but the evidence is not sufficient to recommend such supplementation.[37]

Side effects

Compared with 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 serious environmental problems.

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.[38][39][40]

Acute calcium poisoning is rare, and difficult to achieve without administering calcium intravenously. For example, the oral median lethal dose (LD50) for rats for calcium carbonate and calcium chloride are 6.45[41] and 1.4 g/kg,[42] 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.[43]

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, it is presumed, varies a great deal from person to person. Persons consuming more than 10 grams/day of CaCO3 (=4 g Ca) are at risk of developing milk-alkali syndrome,[44] 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.[45]

Oral calcium supplements diminish the absorption of thyroxine when taken within four to six hours of each other.[46] Thus, people taking both calcium and thyroxine run the risk of inadequate thyroid hormone replacement and thence hypothyroidism if they take them simultaneously or near-simultaneously.[47]

Although some studies have suggested that excessive intake of calcium in the diet or as supplements could be associated with increased cardiovascular mortality,[48][49] other studies found no risk,[50] leading a review to conclude that any risk could only be ascertained with specific further research.[51]

Notable characteristics

Flame test. Brick-red color originates from calcium.

Calcium is reactive and relatively soft for a metal. Although harder than lead, it can be cut with a knife with difficulty. It is a silvery metallic element that can be extracted by electrolysis from a fused salt like calcium chloride.[52] When exposed to the air, it rapidly forms a gray-white coating of calcium oxide and calcium nitride. In bulk form (typically as chips or "turnings"), the metal is somewhat difficult to ignite, more difficult even than magnesium chips; but, when lit, the metal burns in air with a brilliant high-intensity orange-red light. Calcium metal reacts with water, producing hydrogen gas at a moderate rate without generating much heat, making it useful for generating hydrogen.[53] In powdered form, however, the reaction with water is extremely rapid, as the increased surface area of the powder accelerates the reaction. Part of the reason for the slowness of the calcium–water reaction is a partial passivation (chemically protective coating) of insoluble white calcium hydroxide; in acidic solutions, where this compound is more soluble, calcium reacts vigorously.

With a density of 1.54 g/cm3,[54] calcium is the lightest of the alkaline earth metals; magnesium (specific gravity 1.74) and beryllium (1.84) are denser though lighter in atomic mass. From strontium onward, the alkali earth metals become denser with increasing atomic mass. Calcium has two allotropes.[55]

Calcium metal has a higher electrical resistivity than copper or aluminium, yet weight-for-weight, due to its much lower density, it is a better conductor than either. Its use as such in terrestrial applications is usually limited by its high reactivity with air; however, it has potential for use as wiring in off-world applications.[56]

Calcium is the fifth-most-abundant element by mass in the human body, where it is an important cellular ionic messenger with many functions. Calcium also serves as a structural element in bone. It is the relatively high atomic number of calcium that causes bone to be radio-opaque. Of the human body's solid components after drying and burning of organics (as for example, after cremation), about a third of the total "mineral" mass remaining is the approximately one kilogram of calcium that composes the average skeleton (the remainder being mostly phosphorus and oxygen).

H and K lines

Visible spectra of many stars, including the Sun, exhibit strong emission 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. In the Sun or other stars with low temperatures, the prominence of the H and K lines in the visible spectra can be an indication of strong magnetic activity in the chromosphere. Periodic variations of these active regions can indicate the rotation periods of these stars.[57]

Compounds

Calcium chemistry is almost exclusively that of Ca2+ salts.[58] Ca2+ is a "hard cation", that is, it characteristically favors oxide ligands. Hence the abundance of carbonates, nitrates, phosphates, and sulfates in the mineral kingdom. Many of these species crystallize with water. Because it is generally nontoxic and abundant, calcium is found in many foods and useful materials. Most calcium salts are colorless. As with magnesium salts and other alkaline earth metal salts, the halides are soluble in water.

Combined with phosphate, calcium forms hydroxylapatite (Ca5(PO4)3(OH)), the mineral portion of animal bones, teeth, and some corals.[59] Large-scale chemical processes are involved in the conversion of calcium phosphate minerals into fertilizer.

Calcium is the main problematic ion in hard water: it forms insoluble deposits of calcium carbonate that are problematic in plumbing. It also reacts with soap to form soap scum. Calcium carbonate occurs naturally as limestone and chalk. When water percolates through limestone or other calcium-containing rocks, it partially dissolves the rock. The slow re-precipitation of minerals derived from dissolved calcium leads to formation of stalactites and stalagmites.

When heated above 825 °C, calcium carbonate converts calcium oxide (CaO), also known as quicklime:

CaCO3 → CaO + CO2

When added to water, quicklime vigorously reacts (hence its name) to form calcium hydroxide:

CaO + H2O → Ca(OH)2

Also known as slaked lime, Ca(OH)2 is an inexpensive base material used throughout the chemical industry. When 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.

Combined with sulfate, calcium forms the mineral gypsum. When heated to about 300 °F (150 °C), it undergoes partial dehydration:[60]

CaSO4·2H2O → CaSO4·0.5H2O + 1.5 H2O

The resulting powder, when mixed with water, forms a stiff but workable paste that hardens to give Plaster of Paris.

Organocalcium compounds, those containing Ca-C bonds are known, but generally of specialized interest in the research laboratory. One major exception is calcium carbide, which arises from heating calcium compounds with coal or other carbon-rich reducing agents.

CaO + 3 C → CaC2 + CO

It was historically important precursor to acetylene.

CaC2 + 2 H2O → C2H2 + Ca(OH)2

Other important calcium compounds are calcium nitrate used in fertilizers, calcium chloride used as for deicing roads, calcium cyanamide, and calcium hypochlorite, used for bleaching.

Focusing on chemical structure, Ca2+ is a relatively large ion that tends to adopt a high coordination number. In CaF2, the mineral fluorite, each Ca2+ ion is surrounded by eight F− ligands.

Structure of the polymeric [Ca(H2O)6]2+ center in hydrated calcium chloride, illustrating the high coordination number typical for calcium complexes.

Isotopes

Main article: Isotopes of calcium

Calcium has five stable isotopes (40Ca, 42Ca, 43Ca, 44Ca and 46Ca), plus one more (48Ca) that has such a long half-life, it can also be considered stable for all practical purposes. The 20% range in relative mass among naturally occurring calcium isotopes is greater than for any element other than hydrogen and helium. Calcium also has a cosmogenic isotope, radioactive 41Ca, which has a half-life of 103,000 years. Unlike cosmogenic isotopes produced in the atmosphere, 41Ca is produced by neutron activation of 40Ca, primarily in the top metre 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.

Ninety-seven percent of naturally occurring calcium is in the form of 40Ca, 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.

40Ca has a nucleus of 20 protons and 20 neutrons and is the heaviest stable isotope of any element that 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.

Isotope fractionation

As with the isotopes of other elements, a variety of processes fractionate, or alter the relative abundance of, calcium isotopes.[61] 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-dependent differences in calcium isotope composition conventionally are 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.[62]

Calcium isotope fractionation during mineral formation has led to several applications of calcium isotopes. In particular, the 1997 observation by Skulan and DePaolo[63] 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 ration 44Ca/40Ca in soft tissue 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.[64]

A similar system exists in the ocean, where 44Ca/40Ca in seawater 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.[65][66] This has important climatological implications, as the marine calcium cycle is closely tied to the carbon cycle (see below).

Geochemical cycling

Calcium provides an important 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. 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 termed "dissolved inorganic carbon" (DIC).

Travertine terraces Pamukkale, Turkey

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

Note that 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.[67]

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.[68] 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+ but very little CO2.

History

'Ain Ghazal figure

Lime as building material was used since prehistoric times going as far back as 7000 to 14000 BC.[69] Significant statues made from lime plaster date back into the 7 millennia BC were found in 'Ain Ghazal.[70] The first dated lime kiln dates back to 2500 BC and was found in Khafajah mesopotamia.[71][72] Calcium (from Latin calx, genitive calcis, meaning "lime")[73] 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 sulfate), 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.[74] 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: plagioclases, amphiboles, pyroxenes, and garnets.

Cell signaling

The release of calcium ions from the sarcoplasmic reticulum into the cytoplasm is an essential intracellular signal, important in many cellular functions and processes, including muscle contraction, neuronal transmission as in an excitatory synapse, cellular motility (including the movement of flagella and cilia), fertilisation, cell growth or proliferation, learning, memory (as with synaptic plasticity), and secretion of saliva.[75] Calcium signalling can be studied by loading a cell's cytoplasm with a calcium-sensitive fluorescent dye such as Fura-2.[76][77] Many of these dyes were developed by Roger Y. Tsien.[78]

References

  1. ↑ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. p. 112. ISBN 0-08-037941-9.
  2. ↑ Standard Atomic Weights 2013. Commission on Isotopic Abundances and Atomic Weights
  3. ↑ Krieck, Sven; Görls, Helmar; Westerhausen, Matthias (2010). "Mechanistic Elucidation of the Formation of the Inverse Ca(I) Sandwich Complex [(thf)3Ca(μ-C6H3-1,3,5-Ph3)Ca(thf)3] and Stability of Aryl-Substituted Phenylcalcium Complexes". Journal of the American Chemical Society. 132 (35): 12492–12501. doi:10.1021/ja105534w. PMID 20718434.
  4. ↑ Dickson, A. G. & Goyet, C. (1994). "Chapter 5". Handbook of method for the analysis of the various parameters of the carbon dioxide system in sea water, version 2 (PDF). ORNL/CDIAC-74.
  5. ↑ Stephen E. Hluchan and Kenneth Pomerantz "Calcium and Calcium Alloys" in Ullmann's Encyclopedia of Industrial Chemistry, Wiley-VCH, Weinheim, 2006. doi:10.1002/14356007.a04_515.pub2
  6. ↑ Lide, D. R., ed. (2005). CRC Handbook of Chemistry and Physics (86th ed.). Boca Raton (FL): CRC Press. ISBN 0-8493-0486-5.
  7. ↑ Newton, David E (2009-01-01). Food Chemistry. pp. 73–77. ISBN 9781438109756.
  8. ↑ Tordoff, M. G. (2001). "Calcium: Taste, Intake, and Appetite" (PDF). Physiological Reviews. 81 (4): 1567–97. PMID 11581497.
  9. 1 2 "Dietary Supplement Fact Sheet: Calcium". Retrieved 2011-03-08.
  10. ↑ "Dietary Reference Intakes for Calcium and Vitamin D" (PDF). November 2010. Archived from the original (PDF) on 2011-05-19.
  11. ↑ "Dietary Supplement Fact Sheet: Calcium". Office of Dietary Supplements, NIH. Retrieved 2011-03-31.
  12. ↑ Espino J, Bejarano I, Paredes SD, González D, Barriga C, Reiter RJ, Pariente JA, Rodríguez AB (January 2010). "Melatonin Counteracts Altrations in Oxidative Metabolism and Cell Viability Induced by Intracellular Calcium Overload in Human Leucocytes: Changes with Age". Basic & Clinical Pharmacology & Toxicology 107: 590-597. doi:[10.1111/j.1742-7843.2010.00546.x]
  13. ↑ Standing Committee on the Scientific Evaluation of Dietary Reference Intakes, Food and Nutrition Board, Institute of Medicine (1997). Dietary Reference Intakes for Calcium, Phosphorus, Magnesium, Vitamin D and fluoride. Washington DC: The National Academies Press. ISBN 0-309-06403-1.
  14. ↑ Committee to Review Dietary Reference Intakes for Vitamin D and Calcium; Institute of Medicine (2011). A. Catharine Ross; Christine L. Taylor; Ann L. Yaktine; Heather B. Del Valle, eds. Dietary Reference Intakes for Calcium and Vitamin D. ISBN 978-0-309-16394-1.
  15. ↑ Giovannucci E; Rimm EB; Wolk A; et al. (February 1998). "Calcium and fructose intake in relation to risk of prostate cancer". Cancer Research. 58 (3): 442–7. PMID 9458087.
  16. ↑ "Calcium in Diet". Retrieved 14 September 2014.
  17. ↑ Schaafsma, Anne & Beelen, Gerard M (1999). "Eggshell powder, a comparable or better source of calcium than purified calcium carbonate: piglet studies". Journal of the Science of Food and Agriculture. 79 (12): 1596–1600. doi:10.1002/(SICI)1097-0010(199909)79:12<1596::AID-JSFA406>3.0.CO;2-A.
  18. ↑ Schaafsma A; van Doormaal JJ; Muskiet FA; Hofstede GJ; Pakan I; van der Veer E (2002). "Positive effects of a chicken eggshell powder-enriched vitamin-mineral supplement on femoral neck bone mineral density in healthy late post-menopausal Dutch women". Br. J. Nutr. 87 (3): 267–75. doi:10.1079/BJNBJN2001515. PMID 12064336.
  19. ↑ Rovenský J; Stancíková M; Masaryk P; Svík K; Istok R (2003). "Eggshell calcium in the prevention and treatment of osteoporosis". Int J Clin Pharmacol Res. 23 (2–3): 83–92. PMID 15018022.
  20. ↑ "USDA National Nutrient Database". Archived from the original on March 3, 2015.
  21. ↑ Jürgen Kern; et al. (Aug 2016). "Calcium supplementation and risk of dementia in women with cerebrovascular disease". Neurology. doi:10.1212/WNL.0000000000003111.
  22. ↑ "Calcium dementia link is reminder of the dangers of supplements". New Scientist. Aug 27, 2016.
  23. 1 2 3 "Dietary Supplement Fact Sheet: Calcium – Quick Facts — Retrieved 2013-08-23
  24. ↑ Combs, G (2008). The Vitamins. Academic Press. p. 161. ISBN 0-12-183490-5.
  25. 1 2 Weaver, CM (2006). "29 Calcium". In Barbara A. Bowman; Robert M. Russell. Present Knowledge in Nutrition. I (9th ed.). ILSI Press. p. 377. ISBN 1-57881-198-8.
  26. ↑ Remington, Joseph (2005). Remington: The Science and Practice of Pharmacy. Lippincott Williams & Wilkins. p. 1338. ISBN 0-7817-4673-6.
  27. ↑ Zhao, Y; Martin, BR; Weaver, CM (2005). "Calcium bioavailability of calcium carbonate fortified soy milk is equivalent to cow's milk in young women" (PDF). J. Nutr. 135 (10): 2379–82. PMID 16177199.
  28. ↑ Martini, Ligia & Wood, Richard J (2002). "Relative bioavailability of calcium-rich dietary sources in the elderly" (PDF). American Journal of Clinical Nutrition. 76 (6): 1345–1350. PMID 12450902.
  29. 1 2 Straub, D. A. (2007). "Calcium Supplementation in Clinical Practice: A Review of Forms, Doses, and Indications". Nutrition in Clinical Practice. 22 (3): 286–96. doi:10.1177/0115426507022003286. PMID 17507729.
  30. ↑ Martin, Berdine R.; Weaver, CM; Heaney, RP; Packard, PT; Smith, DL (2002). "Calcium Absorption from Three Salts and CaSO4-Fortified Bread in Premenopausal Women". Journal of Agricultural and Food Chemistry. 50 (13): 3874–6. doi:10.1021/jf020065g. PMID 12059174.
  31. ↑ Reid IR, Bristow SM, Bolland MJ (2015). "Calcium supplements: benefits and risks". J. Intern. Med. (Review). 278: 354–68. doi:10.1111/joim.12394. PMID 26174589.
  32. ↑ Bischoff-Ferrari, Heike A.; Dawson-Hughes, Bess; Baron, John A.; Burckhardt, Peter; Li, Ruifeng; Spiegelman, Donna; Specker, Bonny; Orav, John E.; Wong, John B. (2007-12-01). "Calcium intake and hip fracture risk in men and women: a meta-analysis of prospective cohort studies and randomized controlled trials". The American Journal of Clinical Nutrition. 86 (6): 1780–1790. ISSN 0002-9165. PMID 18065599.
  33. ↑ Moyer, Virginia A. (2013-05-07). "Vitamin D and calcium supplementation to prevent fractures in adults: U.S. Preventive Services Task Force recommendation statement". Annals of Internal Medicine. 158 (9): 691–696. doi:10.7326/0003-4819-158-9-201305070-00603. ISSN 1539-3704. PMID 23440163.
  34. ↑ M. J. Bolland, A. Grey, A. Avenell, G. D. Gamble, I. R. Reid. (2011). "Calcium supplements with or without vitamin D and risk of cardiovascular events: reanalysis of the Women's Health Initiative limited access dataset and meta-analysis." BMJ;342:d2040 DOI: 10.1136/bmj.d2040
  35. ↑ Downing L; Islam MA (July 2013). "Influence of calcium supplements on the occurrence of cardiovascular events". Am J Health Syst Pharm (Review). 70 (13): 1132–9. doi:10.2146/ajhp120421. PMID 23784160.
  36. ↑ "Calcium and Cancer Prevention: Strengths and Limits of the Evidence". National Cancer Institute. 4 May 2009. Retrieved November 2, 2014.
  37. ↑ Weingarten MA; Zalmanovici A; Yaphe J (2008). "Dietary calcium supplementation for preventing colorectal cancer and adenomatous polyps". Cochrane Database Syst Rev (Review) (1): CD003548. doi:10.1002/14651858.CD003548.pub4. PMID 18254022.
  38. ↑ Curhan, GC; Willett, WC; Rimm, EB; Stampfer, MJ (1993). "A prospective study of dietary calcium and other nutrients and the risk of symptomatic kidney stones" (PDF). The New England Journal of Medicine. 328 (12): 833–8. doi:10.1056/NEJM199303253281203. PMID 8441427.
  39. ↑ Bihl G; Meyers A. (2001). "Recurrent renal stone disease-advances in pathogenesis and clinical management". Lancet. 358 (9282): 651–656. doi:10.1016/S0140-6736(01)05782-8. PMID 11530173.
  40. ↑ Hall WD; Pettinger M; Oberman A (2001). "Risk factors for kidney stones in older women in the Southern United States". Am J Med Sci. 322 (1): 12–18. doi:10.1097/00000441-200107000-00003. PMID 11465241.
  41. ↑ Lewis, R. J. (1996). Sax's Dangerous Properties of Industrial Materials (9th ed.). New York, NY: Van Nostrand Reinhold. p. 635. ISBN 0-471-37858-5.
  42. ↑ ITII. Toxic and Hazardous Industrial Chemicals Safety Manual. Tokyo, Japan: The International Technical Information Institute. 1988. p. 101.
  43. ↑ Rumack BH. POISINDEX. Information System Micromedex, Inc., Englewood, CO, 2010; CCIS Volume 143. Hall AH and Rumack BH (Eds)
  44. ↑ Beall DP; Henslee HB; Webb HR; Scofield RH (2006). "Milk-alkali syndrome: a historical review and description of the modern version of the syndrome". Am. J. Med. Sci. 331 (5): 233–42. doi:10.1097/00000441-200605000-00001. PMID 16702792.
  45. ↑ Picolos MK & Orlander PR (2005). "Calcium carbonate toxicity: The updated milk-alkali syndrome; report of 3 cases and review of the literature". Endocrine Practice. 4 (11): 272–80. doi:10.4158/EP.11.4.272. PMID 16006300.
  46. ↑ Singh, N.; Singh, P.; Hershman, J. (2000). "Effect of calcium carbonate on the absorption of levothyroxine". JAMA: the Journal of the American Medical Association. 283 (21): 2822–2825. doi:10.1001/jama.283.21.2822. PMID 10838651.
  47. ↑ Harding, Anne. "Calcium May Help With Weight Loss". Archived from the original on 2007-07-03. Retrieved 2007-07-10.
  48. ↑ Michaëlsson K; Melhus H; Warensjö Lemming E; Wolk A; Byberg L (2013). "Long term calcium intake and rates of all cause and cardiovascular mortality: community based prospective longitudinal cohort study". BMJ. 346: f228. doi:10.1136/bmj.f228. PMC 3571949Freely accessible. PMID 23403980.
  49. ↑ Xiao Q; Murphy RA; Houston DK; Harris TB; Chow WH; Park Y (2013). "Dietary and supplemental calcium intake and cardiovascular disease mortality: the National Institutes of Health-AARP diet and health study". JAMA Internal Medicine. 173 (8): 639–646. doi:10.1001/jamainternmed.2013.3283. PMC 3756477Freely accessible. PMID 23381719.
  50. ↑ Paik JM, Curhan GC, Sun Q, et al. (2014). "Calcium supplement intake and risk of cardiovascular disease in women". Osteoporos Int. 25 (8): 2047–56. doi:10.1007/s00198-014-2732-3. PMC 4102630Freely accessible. PMID 24803331.
  51. ↑ Waldman T, Sarbaziha R, Merz CN, Shufelt C (2015). "Calcium Supplements and Cardiovascular Disease: A Review". Am J Lifestyle Med. 9 (4): 298–307. doi:10.1177/1559827613512593. PMC 4560115Freely accessible. PMID 26345134.
  52. ↑ Pauling, Linus (1970). General Chemistry. Dover Publications. p. 627. ISBN 0-7167-0149-9.
  53. ↑ Theodore Gray. The Elements. Page 55
  54. ↑ http://education.jlab.org/itselemental/ele020.html
  55. ↑ Bale, C. W.; Pelton, A. D. (1987). "The Ca−Li (Calcium-Lithium) system". Journal of Phase Equilibria. 8 (2): 125–127. doi:10.1007/BF02873196.
  56. ↑ Geoffrey Landis (2014). "Resource Production on the Moon" (PDF). NASA. Retrieved 2015-05-20.
  57. ↑ Staff (1995). "H-K Project". Mount Wilson Observatory. Archived from the original on 2006-08-31. Retrieved 2006-08-10.
  58. ↑ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0-08-037941-9.
  59. ↑ Hu, J.; Russell, J. J.; Ben-Nissan, B.; Vago, R. (2001). "Production and analysis of hydroxyapatite from Australian corals via hydrothermal process". Journal of Materials Science Letters. 20 (1): 85–87. doi:10.1023/A:1006735319725. ISSN 0261-8028.
  60. ↑ Franz Wirsching "Calcium Sulfate" in Ullmann's Encyclopedia of Industrial Chemistry, 2012 Wiley-VCH, Weinheim. doi:10.1002/14356007.a04_555
  61. ↑ Russell, WA; Papanastassiou, DA; Tombrello, TA (1978). "Ca isotope fractionation on the earth and other solar system materials". Geochim Cosmochim Acta. 42 (8): 1075–90. Bibcode:1978GeCoA..42.1075R. doi:10.1016/0016-7037(78)90105-9.
  62. ↑ Skulan, J; Depaolo, DJ (1999). "Calcium isotope fractionation between soft and mineralized tissues as a monitor of calcium use in vertebrates" (PDF). Proc Natl Acad Sci USA. 96 (24): 13709–13. Bibcode:1999PNAS...9613709S. doi:10.1073/pnas.96.24.13709. PMC 24129Freely accessible. PMID 10570137.
  63. ↑ Skulan, J; Depaolo, DJ; Owens, TL (June 1997). "Biological control of calcium isotopic abundances in the global calcium cycle". Geochimica et Cosmochimica Acta. 61 (12): 2505–10. Bibcode:1997GeCoA..61.2505S. doi:10.1016/S0016-7037(97)00047-1.
  64. ↑ Skulan, J; Bullen, T; Anbar, AD; Puzas, JE; Shackelford, L; Leblanc, A; Smith, SM (2007). "Natural calcium isotopic composition of urine as a marker of bone mineral balance" (PDF). Clinical Chemistry. 653 (6): 1155–1158. doi:10.1373/clinchem.2006.080143. PMID 17463176.
  65. ↑ Fantle, M; Depaolo, D (2007). "Ca isotopes in carbonate sediment and pore fluid from ODP Site 807A: The Ca2+(aq)–calcite equilibrium fractionation factor and calcite recrystallization rates in Pleistocene sediments". Geochim Cosmochim Acta. 71 (10): 2524–2546. Bibcode:2007GeCoA..71.2524F. doi:10.1016/j.gca.2007.03.006.
  66. ↑ Griffith, Elizabeth M.; Paytan, Adina; Caldeira, Ken; Bullen, Thomas; Thomas, Ellen (2008). "A Dynamic marine calcium cycle during the past 28 million years". Science. 322 (12): 1671–1674. Bibcode:2008Sci...322.1671G. doi:10.1126/science.1163614. PMID 19074345.
  67. ↑ Zeebe (2006). "Marine carbonate chemistry". National Council for Science and the Environment. Retrieved 2010-03-13.
  68. ↑ Berner, Robert (2003). "The long-term carbon cycle, fossil fuels and atmospheric composition". Nature. 426 (6964): 323–326. Bibcode:2003Natur.426..323B. doi:10.1038/nature02131. PMID 14628061.
  69. ↑ Miller, M. Michael. "Commodity report:Lime" (PDF). United States Geological Survey. Retrieved 2012-03-06.
  70. ↑ Garfinkel, Yosef (1987). "Burnt Lime Products and Social Implications in the Pre-Pottery Neolithic B Villages of the Near East". Paléorient. 13 (1): 69–76. doi:10.3406/paleo.1987.4417. JSTOR 41492234.
  71. ↑ Williams, Richard (2004). Lime Kilns and Lime Burning. p. 4. ISBN 978-0-7478-0596-0.
  72. ↑ Oates, J. A. H (2008-07-01). Lime and Limestone: Chemistry and Technology, Production and Uses. ISBN 978-3-527-61201-7.
  73. ↑ calx. Charlton T. Lewis and Charles Short. A Latin Dictionary on Perseus Project.
  74. ↑ Davy H (1808). "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". Philosophical Transactions of the Royal Society of London. 98: 333–370. Bibcode:1808RSPT...98..333D. doi:10.1098/rstl.1808.0023.
  75. ↑ Santulli, Gaetano; Marks, Andrew (2015). "Essential Roles of Intracellular Calcium Release Channels in Muscle, Brain, Metabolism, and Aging". Current Molecular Pharmacology. 8 (2): 206–222. doi:10.2174/1874467208666150507105105. ISSN 1874-4672.
  76. ↑ D. González, I. Bejarano, C. Barriga, A.B. Rodríguez, J.A. Pariente (2010). "Oxidative Stress-Induced Caspases are Regulated in Human Myeloid HL-60 Cells by Calcium Signal". Current Signal Transduction Therapy 5: 181-186. doi:[10.2174/157436210791112172]
  77. ↑ Bejarano I, Espino J, González-Flores D, Casado JG, Redondo PC, Rosado JA, Barriga C, Pariente JA, Rodríguez AB (2009). "Role of Calcium Signals on Hydrogen Peroxide-Induced Apoptosis in Human Myeloid HL-60 Cells". International Journal of Biomedical science 5(3): 246-256.
  78. ↑ Palmer, Amy E; Tsien, Roger Y (2006). "Measuring calcium signaling using genetically targetable fluorescent indicators". Nature Protocols. 1 (3): 1057–1065. doi:10.1038/nprot.2006.172. ISSN 1754-2189.

Further reading

Periodic table (Large cells)
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
1 H He
2 Li Be B C N O F Ne
3 Na Mg Al Si P S Cl Ar
4 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
5 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe
6 Cs Ba La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
7 Fr Ra Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Rf Db Sg Bh Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Og
Alkali metal Alkaline earth metal Lan­thanide Actinide Transition metal Post-​transition metal Metalloid Polyatomic nonmetal Diatomic nonmetal Noble gas Unknown
chemical
properties
This article is issued from Wikipedia - version of the 12/3/2016. The text is available under the Creative Commons Attribution/Share Alike but additional terms may apply for the media files.