Difference between revisions of "Calcium"

Latest revision as of 20:38, 19 September 2010

Calcium (11px /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.^[2]

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.^[3] Once produced, it rapidly forms a gray-white oxide and nitride coating when exposed to air. In bulk-form (typically as chips or "turnings") the metal is somewhat difficult to ignite, more so even than magnesium chips; 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 this salt is soluble, calcium reacts vigorously.

Calcium, with a density of 1.55 g/cm³, is the lightest of the alkali earth metals; magnesium (specific gravity 1.74) and beryllium (1.84) are more dense, although lighter in atomic mass. From strontium onward, the alkali earth metals become more dense with increasing atomic mass.

Calcium has a higher electrical resistivity than copper or aluminium, yet weight-for-weight, due to 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.^[4] 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 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 which 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 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.^[5]

Compounds

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.

A few calcium compounds in the oxidation state +1 have also been investigated recently.^[1]

Nucleosynthesis

Stable Calcium is created in extremely large, extremely hot (over 2.5 billion kelvin) stars. It requires one atom of argon and one atom of helium.^{[citation needed]}

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]}

Isotope fractionation

As with the isotopes of other elements, a variety of processes fractionate, or alter the relative abundance of, calcium isotopes.^[6] 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.^[7]

Calcium isotope fractionation during mineral formation has led to several applications of calcium isotopes. In particular, the 1997 observation by Skulan and DePaolo^[8] 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.^[9]

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.^[10][11] 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 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⁺ + 2HCO−3 → 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.^[12]

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.^[13] 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.

History

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 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.^[14] 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

See also Category: Calcium minerals 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

Applications

Calcium is used^[15]

• as a reducing agent in the extraction of other metals, such as uranium, zirconium, and thorium.
• as a deoxidizer, desulfurizer, or decarbonizer for various ferrous and nonferrous alloys.
• as an alloying agent used in the production of aluminium, beryllium, copper, lead, and magnesium alloys.
• in the making of cements and mortars to be used in construction.
• in the making of cheese, where calcium ions influence the activity of rennin in bringing about the coagulation of milk.

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Calcium compounds

• 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.
• Calcium hydroxide solution (Ca(OH)₂) (also known as limewater) is used to detect the presence of carbon dioxide by being bubbled through a solution. It turns cloudy where CO₂ is present.
• Calcium arsenate (Ca₃(AsO₄)₂) is used in insecticides.
• Calcium carbide (CaC₂) is used to make acetylene gas (for use in acetylene torches for welding) and in the manufacturing of plastics.
• Calcium chloride (CaCl₂) is used in ice removal and dust control on dirt roads, in conditioner for concrete, as an additive in canned tomatoes, and to provide body for automobile tires.
• Calcium cyclamate (Ca(C₆H₁₁NHSO₃)₂) was used as a sweetening agent but is no longer permitted for use because of suspected cancer-causing properties.
• Calcium gluconate (Ca(C₆H₁₁O₇)₂) is used as a food additive and in vitamin pills.
• Calcium hypochlorite (Ca(OCl)₂) is used as a swimming pool disinfectant, as a bleaching agent, as an ingredient in deodorant, and in algaecide and fungicide.
• Calcium permanganate (Ca(MnO₄)₂) is used in liquid rocket propellant, textile production, as a water sterilizing agent and in dental procedures.
• Calcium phosphate (Ca₃(PO₄)₂) is used as a supplement for animal feed, fertilizer, in commercial production for dough and yeast products, in the manufacture of glass, and in dental products.
• Calcium phosphide (Ca₃P₂) is used in fireworks, rodenticide, torpedoes and flares.
• Calcium stearate (Ca(C₁₈H₃₅O₂)₂) is used in the manufacture of wax crayons, cements, certain kinds of plastics and cosmetics, as a food additive, in the production of water resistant materials and in the production of paints.
• Calcium sulfate (CaSO₄·2H₂O) is used as common blackboard chalk, as well as, in its hemihydrate form better known as Plaster of Paris.
• Calcium tungstate (CaWO₄) is used in luminous paints, fluorescent lights and in X-ray studies.
• Hydroxylapatite (Ca₅(PO₄)₃(OH), but is usually written Ca₁₀(PO₄)₆(OH)₂) makes up seventy percent of bone. Also carbonated-calcium deficient hydroxylapatite is the main mineral of which dental enamel and dentin are comprised.

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.^[17] 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.^[18] 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.^[19][20][21] 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.^[22][23][24] 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.^[25]. An overlooked source of calcium is eggshell, which can be ground into a powder and mixed into food or a glass of water.^[26][27][28] Cultivated vegetables generally have less calcium than wild plants.^[29]

The calcium content of most foods can be found in the USDA National Nutrient Database.^[30]

Dietary calcium supplements

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.^[16] 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.^[31]

• The absorption of calcium from most food and commonly used dietary supplements is very similar.^[32] This is contrary to what many calcium supplement manufacturers claim in their promotional materials.
• Milk is an excellent source of dietary calcium for those whose body tolerate it because it has a high concentration of calcium and the calcium in milk is excellently absorbed.^[32]
• Soymilk and other vegetable milks are usually sold with calcium added so that their calcium concentration is as high as in milk
• Also different kind of juices boosted with calcium are widely available.
• Calcium carbonate is the most common and least expensive calcium supplement. It should be taken with food. It depends on low pH levels for proper absorption in the intestine.^[33] Some studies suggests that the absorption of calcium from calcium carbonate is similar to the absorption of calcium from milk.^[34][35] While most people digest calcium carbonate very well, some might develop gastrointestinal discomfort or gas. Taking magnesium with it can help to avoid constipation. Calcium carbonate is 40% elemental calcium. 1000 mg will provide 400 mg of calcium. However, supplement labels will usually indicate how much calcium is present in each serving, not how much calcium carbonate is present.
• Antacids frequently contain calcium carbonate, and are a commonly used, inexpensive calcium supplement
• Coral Calcium is a salt of calcium derived from fossilized coral reefs. Coral calcium is composed of calcium carbonate and trace minerals.
• Calcium citrate can be taken without food and is the supplement of choice for individuals with achlorhydria or who are taking histamine-2 blockers or proton-pump inhibitors.^[36] It is more easily digested and absorbed than calcium carbonate if taken on an empty stomach and less likely to cause constipation and gas than calcium carbonate. It also has a lower risk of contributing to the formation of kidney stones. Calcium citrate is about 21% elemental calcium. 1000 mg will provide 210 mg of calcium. It is more expensive than calcium carbonate and more of it must be taken to get the same amount of calcium.
• Calcium phosphate costs more than calcium carbonate, but less than calcium citrate. It is easily absorbed and is less likely to cause constipation and gas than either.
• Calcium lactate has similar absorption as calcium carbonate^[37], but is more expensive. Calcium lactate and calcium gluconate are less concentrated forms of calcium and are not practical oral supplements.^[36]
• Calcium chelates are synthetic calcium compounds, with calcium bound to an organic molecule, such as malate, aspartate, or fumarate. These forms of calcium may be better absorbed on an empty stomach. However, in general they are absorbed similarly to calcium carbonate and other common calcium supplements when taken with food.^[38] The 'chelate' mimics the action that natural food performs by keeping the calcium soluble in the intestine. Thus, on an empty stomach, in some individuals, chelates might theoretically be absorbed better.
• Microcrystalline hydroxyapatite (MH) is marketed as a calcium supplement, and has in some randomized trials been found to be more effective than calcium carbonate.

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.^[39]

Prevention of fractures due to osteoporosis

Such studies often do not test calcium alone, but rather combinations of calcium and vitamin D. Randomized controlled trials found both positive^[40][41] and negative^{[42][43][44][45]} effects. The different results may be explained by doses of calcium and underlying rates of calcium supplementation in the control groups.^[46] 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]}

Possible cancer prevention

A meta-analysis^[41] by the international Cochrane Collaboration of two randomized controlled trials^[47][48] 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) ^[49].

• A randomized controlled trial found that 1000 mg of elemental calcium and 400 IU of vitamin D₃ had no effect on colorectal cancer^[50]
• A randomized controlled trial found that 1400–1500 mg supplemental calcium and 1100 IU vitamin D₃ reduced aggregated cancers with a relative risk of 0.402.^[51]
• An observational cohort study found that high calcium and vitamin D intake was associated with "lower risk of developing premenopausal breast cancer."^[52]

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 median lethal dose (LD^{50) for rats for calcium carbonate and calcium chloride are 6.45 [53] and 1.4 g/kg,[54] 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.^[55]

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,^[56] 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.^[57]

See also

• Calcium metabolism
• Calcium in biology
• Calcium compounds
• Disorders of calcium metabolism

References

Further reading

• Rebecca J. Donatelle. Health, The Basics. 6th ed. San Francisco: Pearson Education, Inc. 2005.

External links

• WebElements.com — Calcium
• USDA National Nutrient Database, Calcium content of selected foods
• UK Food Standards Agency: Calcium
• Nutrition fact sheet from the National Institutes of Health

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1. ↑ ^{1.0 1.1} 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: 100818110534020. doi:10.1021/ja105534w. 
2. ↑ 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 (PDF). ORNL/CDIAC-74. CS1 maint: Multiple names: authors list (link)
3. ↑ Pauling, Linus (1970). General Chemistry. Dover Publications. p. 627. ISBN 0716701499. 
4. ↑ Tordoff, M. G. (2001). "Calcium: Taste, Intake, and Appetite". Physiological Reviews. 81 (4): 1567. PMID 11581497. 
5. ↑ Staff (1995). "H-K Project". Mount Wilson Observatory. Retrieved 2006-08-10. 
6. ↑ Russell, WA; Papanastassiou, DA; Tombrello, TA (1978). "Ca isotope fractionation on the earth and other solar system materials". Geochim Cosmochim Acta. 42: 1075–90. doi:10.1016/0016-7037(78)90105-9. 
7. ↑ Skulan, J; DePaolo, DJ (1999). "Calcium isotope fractionation between soft and mineralized tissues as a monitor of calcium use in vertebrates". Proc Natl Acad Sci USA. 96 (24): 13709–13. doi:10.1073/pnas.96.24.13709. PMC 24129 Freely accessible. PMID 10570137. 
8. ↑ Skulan, J; DePaolo, D; Anbar, AD; Puzas, JE; Shackelford, L; Leblanc, A; Smith, SM (1997). "Biological control of calcium isotopic abundances in the global calcium cycle". Geochim. Et Cosmochim. Acta. 61 (12): 2505–2510. doi:10.1016/S0016-7037(97)00047-1. PMID 17463176. 
9. ↑ 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". Clinical Chemistry. 653 (6): 1155–1158. doi:10.1373/clinchem.2006.080143. PMID 17463176. 
10. ↑ 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: 2524–2546. doi:10.1016/j.gca.2007.03.006. 
11. ↑ Griffith, Elizabeth M.; Paytan, Adina; Caldeira, Ken; Bullen, Thomas last5= Thomas; Shackelford, Ellen; Leblanc, A; Smith, SM (2008). "A Dynamic marine calcium cycle during the past 28 million years". Science. 322 (12): 1671–1674. doi:10.1126/science.1163614. PMID 17463176. CS1 maint: Missing pipe (link)
12. ↑ Zeebe (2006). "Marine carbonate chemistry". National Council for Science and the Environment. Retrieved 2010-03-13. 
13. ↑ Berner, Robert (2003). "The long-term carbon cycle, fossil fuels and atmospheric composition". Nature. 426 (6964): 323–326. doi:10.1038/nature02131. PMID 14628061. 
14. ↑ 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, vol. 98, pages 333-370.
15. ↑ Lide, D. R., ed. (2005), CRC Handbook of Chemistry and Physics (86th ed.), Boca Raton (FL): CRC Press, ISBN 0-8493-0486-5 
16. ↑ ^{16.0 16.1} "Dietary Supplement Fact Sheet: Calcium". 
17. ↑ "Osteoporosis Prevention - Calcium Recommendations". 
18. ↑ 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 0309064031. CS1 maint: Multiple names: authors list (link)
19. ↑ 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". The New England journal of medicine. 328 (12): 833–8. doi:10.1056/NEJM199303253281203. PMID 8441427. CS1 maint: Multiple names: authors list (link)
20. ↑ 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. 
21. ↑ 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. CS1 maint: Multiple names: authors list (link)
22. ↑ Abelow, BJ; Holford, TR; Insogna, KL (1992). "Cross-cultural association between dietary animal protein and hip fracture: a hypothesis". Calcified tissue international. 50 (1): 14–8. doi:10.1007/BF00297291. PMID 1739864. 
23. ↑ Hegsted, M; Schuette, SA; Zemel, MB; Linkswiler, HM (1981). "Urinary calcium and calcium balance in young men as affected by level of protein and phosphorus intake". The Journal of nutrition. 111 (3): 553–62. PMID 7205408. 
24. ↑ Kerstetter JE and Allen LH (1990). "Dietary protein increases urinary calcium" (PDF). Journal of Nutrition. 120 (1): 134. PMID 2406396. 
25. ↑ Tucker, KL; Chen, H; Hannan, MT; Cupples, LA; Wilson, PW; Felson, D; Kiel, DP (2002). "Bone mineral density and dietary patterns in older adults: the Framingham Osteoporosis Study". The American journal of clinical nutrition. 76 (1): 245–52. PMID 12081842. 
26. ↑ Schaafsma, Anne and 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. CS1 maint: Multiple names: authors list (link)
27. ↑ 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. CS1 maint: Multiple names: authors list (link)
28. ↑ 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. CS1 maint: Multiple names: authors list (link)
29. ↑ "Original Wild Foods vs. Available Foods Today for Instinctos". 
30. ↑ "USDA National Nutrient Database". 
31. ↑ Combs, G (2008). The Vitamins. Academic Press. p. 161. ISBN 0121834905. 
32. ↑ ^{32.0 32.1} Weaver, CM (2006). "Calcium". Present Knowledge in Nutrition, 9th Ed. I. ILSI Press. p. 377. 
33. ↑ Remington, Joseph (2005). Remington: The Science and Practice of Pharmacy. Lippincott Williams & Wilkins. p. 1338. ISBN 0781746736. 
34. ↑ Zhao, Y; Martin, BR; Weaver, CM (2005). "Calcium bioavailability of calcium carbonate fortified soy milk is equivalent to cow's milk in young women". J. Nutr. 135 (10): 2379. PMID 16177199. 
35. ↑ Martini, Ligia and Wood, Richard J (2002). "Relative bioavailability of calcium-rich dietary sources in the elderly". American Journal of Clinical Nutrition. 76 (6): 1345–1350. PMID 12450902. CS1 maint: Multiple names: authors list (link)
36. ↑ ^{36.0 36.1} Straub, D. A. (2007). "Calcium Supplementation in Clinical Practice: A Review of Forms, Doses, and Indications". Nutrition in Clinical Practice. 22 (3): 286. doi:10.1177/0115426507022003286. PMID 17507729. 
37. ↑ 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. doi:10.1021/jf020065g. PMID 12059174. 
38. ↑ Weaver, Connie M.; Martin, BR; Costa, NM; Saleeb, FZ; Huth, PJ (2002). "Absorption of Calcium Fumarate Salts Is Equivalent to Other Calcium Salts When Measured in the Rat Model". Journal of Agricultural and Food Chemistry. 50 (17): 4974. doi:10.1021/jf0200422. PMID 12166992. 
39. ↑ Harding, Anne. "Calcium May Help With Weight Loss". Retrieved 2007-07-10. 
40. ↑ Dawson-Hughes B, Harris SS, Krall EA, Dallal GE (1997). "Effect of calcium and vitamin D supplementation on bone density in men and women 65 years of age or older". N. Engl. J. Med. 337 (10): 670–6. doi:10.1056/NEJM199709043371003. PMID 9278463. CS1 maint: Multiple names: authors list (link)
41. ↑ ^{41.0 41.1} Weingarten MA, Zalmanovici A, Yaphe J (2005). "Dietary calcium supplementation for preventing colorectal cancer, adenomatous polyps and calcium metabolisism disorder". Cochrane database of systematic reviews (Online) (3): CD003548. doi:10.1002/14651858.CD003548.pub3. PMID 16034903. CS1 maint: Multiple names: authors list (link)
42. ↑ Jackson RD, LaCroix AZ, Gass M (2006). "Calcium plus vitamin D supplementation and the risk of fractures". N. Engl. J. Med. 354 (7): 669–83. doi:10.1056/NEJMoa055218. PMID 16481635. CS1 maint: Multiple names: authors list (link)
43. ↑ Grant AM, Avenell A, Campbell MK (2005). "Oral vitamin D3 and calcium for secondary prevention of low-trauma fractures in elderly people (Randomised Evaluation of Calcium Or vitamin D, RECORD): a randomised placebo-controlled trial". Lancet. 365 (9471): 1621–8. doi:10.1016/S0140-6736(05)63013-9. PMID 15885294. CS1 maint: Multiple names: authors list (link)
44. ↑ Porthouse J, Cockayne S, King C (2005). "Randomised controlled trial of calcium and supplementation with cholecalciferol (vitamin D3) for prevention of fractures in primary care". BMJ. 330 (7498): 1003. doi:10.1136/bmj.330.7498.1003. PMC 557150 Freely accessible. PMID 15860827. CS1 maint: Multiple names: authors list (link)
45. ↑ Prince RL, Devine A, Dhaliwal SS, Dick IM (2006). "Effects of calcium supplementation on clinical fracture and bone structure: results of a 5-year, double-blind, placebo-controlled trial in elderly women". Arch. Intern. Med. 166 (8): 869–75. doi:10.1001/archinte.166.8.869. PMID 16636212. CS1 maint: Multiple names: authors list (link)
46. ↑ Fletcher RH (2006). "Calcium plus vitamin D did not prevent hip fracture or colorectal cancer in postmenopausal women" (subscription required). ACP J. Club. 145 (1): 4–5. PMID 16813354. 
47. ↑ Baron JA, Beach M, Mandel JS (1999). "Calcium supplements for the prevention of colorectal adenomas. Calcium Polyp Prevention Study Group". N. Engl. J. Med. 340 (2): 101–7. doi:10.1056/NEJM199901143400204. PMID 9887161. CS1 maint: Multiple names: authors list (link)
48. ↑ Bonithon-Kopp C, Kronborg O, Giacosa A, Räth U, Faivre J (2000). "Calcium and fibre supplementation in prevention of colorectal adenoma recurrence: a randomised intervention trial. European Cancer Prevention Organisation Study Group". Lancet. 356 (9238): 1300–6. doi:10.1016/S0140-6736(00)02813-0. PMID 11073017. CS1 maint: Multiple names: authors list (link)
49. ↑ Lappe, Jm; Travers-Gustafson, D; Davies, Km; Recker, Rr; Heaney, Rp (2007). "Vitamin D and calcium supplementation reduces cancer risk: results of a randomized trial". The American journal of clinical nutrition (Free full text)|format= requires |url= (help). 85 (6): 1586–91. PMID 17556697. CS1 maint: Multiple names: authors list (link)
50. ↑ Wactawski-Wende J, Kotchen JM, Anderson GL (2006). "Calcium plus vitamin D supplementation and the risk of colorectal cancer". N. Engl. J. Med. 354 (7): 684–96. doi:10.1056/NEJMoa055222. PMID 16481636. CS1 maint: Multiple names: authors list (link)
51. ↑ Lappe JM, Travers-Gustafson D, Davies KM, Recker RR, Heaney RP (2007). "Vitamin D and calcium supplementation reduces cancer risk: results of a randomized trial". Am. J. Clin. Nutr. 85 (6): 1586–91. PMID 17556697. CS1 maint: Multiple names: authors list (link)
52. ↑ Lin J, Manson JE, Lee IM, Cook NR, Buring JE, Zhang SM (2007). "Intakes of calcium and vitamin d and breast cancer risk in women". Arch. Intern. Med. 167 (10): 1050–9. doi:10.1001/archinte.167.10.1050. PMID 17533208. CS1 maint: Multiple names: authors list (link)
53. ↑ Lewis, R. J. (1996). Sax's Dangerous Properties of Industrial Materials (9 ed.). New York, NY: Van Nostrand Reinhold. p. 635. ISBN 0471378585. 
54. ↑ ITII. Toxic and Hazardous Industrial Chemicals Safety Manual. Tokyo, Japan: The International Technical Information Institute. 1988. p. 101. 
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56. ↑ 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. CS1 maint: Multiple names: authors list (link)
57. ↑ Picolos MK and 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.