Calcium

potassium ← calcium → scandium
Mg
↑
Ca
↓
Sr
Appearance
Dull gray, silver


Spectral lines of Calcium
General properties
Name, symbol, number calcium, Ca, 20
Pronunciation /ˈkælsiəm/ KAL-see-əm
Element category alkaline earth metal
Group, period, block 2, 4, s
Standard atomic weight 40.078g·mol−1
Electron configuration [Ar] 4s2
Electrons per shell 2, 8, 8, 2 (Image)
Physical properties
Phase solid
Density (near r.t.) 1.55 g·cm−3
Liquid density at m.p. 1.378 g·cm−3
Melting point 1115 K, 842 °C, 1548 °F
Boiling point 1757 K, 1484 °C, 2703 °F
Heat of fusion 8.54 kJ·mol−1
Heat of vaporization 154.7 kJ·mol−1
Specific heat capacity (25 °C) 25.929 J·mol−1·K−1
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[1]
(strongly basic oxide)
Electronegativity 1.00 (Pauling scale)
Ionization energies
(more)		 1st: 589.8 kJ·mol−1
2nd: 1145.4 kJ·mol−1
3rd: 4912.4 kJ·mol−1
Atomic radius 197 pm
Covalent radius 176±10 pm
Van der Waals radius 231 pm
Miscellanea
Crystal structure face-centered cubic
Magnetic ordering diamagnetic
Electrical resistivity (20 °C) 33.6 nΩ·m
Thermal conductivity (300 K) 201 W·m−1·K−1
Thermal expansion (25 °C) 22.3 µm·m−1·K−1
Speed of sound (thin rod) (20 °C) 3810 m/s
Young's modulus 20 GPa
Shear modulus 7.4 GPa
Bulk modulus 17 GPa
Poisson ratio 0.31
Mohs hardness 1.75
Brinell hardness 167 MPa
CAS registry number 7440-70-2
Most stable isotopes
Main article: Isotopes of calcium
iso NA half-life DM DE (MeV) DP
40Ca 96.941% 40Ca is stable with 20 neutrons
41Ca trace 1.03×105 y ε - 41K
42Ca 0.647% 42Ca is stable with 22 neutrons
43Ca 0.135% 43Ca is stable with 23 neutrons
44Ca 2.086% 44Ca is stable with 24 neutrons
45Ca syn 162.7 d β− 0.258 45Sc
46Ca 0.004% >2.8×1015 y β−β−  ? 46Ti
47Ca syn 4.536 d β− 0.694, 1.99 47Sc
γ 1.297 -
48Ca 0.187% >4×1019 y β−β−  ? 48Ti

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.[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.

Contents

Notable characteristics

Flame test. Brick red color originates from calcium.

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.

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.

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.

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).

Travertine terraces Pamukkale, Turkey

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]

Calcium compounds

Nutrition

Recommended adequate intake by the IOM for calcium:[16]
Age Calcium (mg/day)
0–6 months 210
7–12 months 270
1–3 years 500
4–8 years 800
9–18 years 1300
19–50 years 1000
51+ years 1200

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

500 milligram calcium supplements made from calcium carbonate

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]

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.

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].

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 (LD50) 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

References

  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 132 (35): 100818110534020. doi:10.1021/ja105534w. PMID 20718434. 
  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. ORNL/CDIAC-74. http://cdiac.esd.ornl.gov/ftp/cdiac74/chapter5.pdf. 
  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. http://physrev.physiology.org/cgi/content/full/81/4/1567. 
  5. ↑ Staff (1995). "H-K Project". Mount Wilson Observatory. http://www.mtwilson.edu/hk/. 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. 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. 
  12. ↑ Zeebe (2006). "Marine carbonate chemistry". National Council for Science and the Environment. http://www.eoearth.org/article/Marine_carbonate_chemistry. 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". http://ods.od.nih.gov/factsheets/calcium.asp. 
  17. ↑ "Osteoporosis Prevention - Calcium Recommendations". http://www.nof.org/prevention/calcium2.htm. 
  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. 
  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. 
  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. 
  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". Journal of Nutrition 120 (1): 134. PMID 2406396. http://jn.nutrition.org/cgi/reprint/120/1/134.pdf. 
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  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. 
  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. 
  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. 
  29. ↑ "Original Wild Foods vs. Available Foods Today for Instinctos". http://www.beyondveg.com/nieft-k/instincto-guide/instincto-guide1e.shtml. 
  30. ↑ "USDA National Nutrient Database". http://www.nal.usda.gov/fnic/foodcomp/search. 
  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. pp. 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. http://jn.nutrition.org/cgi/content/full/135/10/2379. 
  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. http://www.ajcn.org/cgi/content/abstract/76/6/1345. 
  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". http://www.rxalternativemedicine.com/headlines_news.php#headline77. 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. 
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  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. 
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