Hubbry Logo
CalciumCalciumMain
Open search
Calcium
Community hub
Calcium
logo
8 pages, 0 posts
0 subscribers
Be the first to start a discussion here.
Be the first to start a discussion here.
Calcium
Calcium
Calcium
View on Wikipedia
from Wikipedia

Calcium, 20Ca
Crystals of silvery calcium metal in a glass vial
Calcium
Appearancedull gray, silver; with a pale yellow tint[1]
Standard atomic weight Ar°(Ca)
Calcium in the periodic table
Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon
Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson
Mg
↑
Ca
↓
Sr
potassium ← calcium → scandium
Atomic number (Z)20
Groupgroup 2 (alkaline earth metals)
Periodperiod 4
Block  s-block
Electron configuration[Ar] 4s2
Electrons per shell2, 8, 8, 2
Physical properties
Phase at STPsolid
Melting point1115 K ​(842 °C, ​1548 °F)
Boiling point1757 K ​(1484 °C, ​2703 °F)
Density (at 20° C)1.526 g/cm3 [4]
when liquid (at m.p.)1.378 g/cm3
Heat of fusion8.54 kJ/mol
Heat of vaporisation154.7 kJ/mol
Molar heat capacity25.929 J/(mol·K)
Vapour pressure
P (Pa) 1 10 100 1 k 10 k 100 k
at T (K) 864 956 1071 1227 1443 1755
Atomic properties
Oxidation statescommon: +2
+1[5]
ElectronegativityPauling scale: 1.00
Ionisation energies
  • 1st: 589.8 kJ/mol
  • 2nd: 1145.4 kJ/mol
  • 3rd: 4912.4 kJ/mol
  • (more)
Atomic radiusempirical: 197 pm
Covalent radius176±10 pm
Van der Waals radius231 pm
Color lines in a spectral range
Spectral lines of calcium
Other properties
Natural occurrenceprimordial
Crystal structure ​face-centred cubic (fcc) (cF4)
Lattice constant
Face-centered cubic crystal structure for calcium
a = 558.8 pm (at 20 °C)[4]
Thermal expansion22.27×10−6/K (at 20 °C)[4]
Thermal conductivity201 W/(m⋅K)
Electrical resistivity33.6 nΩ⋅m (at 20 °C)
Magnetic orderingdiamagnetic
Molar magnetic susceptibility+40.0×10−6 cm3/mol[6]
Young's modulus20 GPa
Shear modulus7.4 GPa
Bulk modulus17 GPa
Speed of sound thin rod3810 m/s (at 20 °C)
Poisson ratio0.31
Mohs hardness1.75
Brinell hardness170–416 MPa
CAS Number7440-70-2
History
Namingfrom the Latin word for lime, calx
Discovery and first isolationHumphry Davy (1808)
Isotopes of calcium
Main isotopes[7] Decay
Isotope abun­dance half-life (t1/2) mode pro­duct
40Ca 96.9% stable
41Ca trace 9.94×104 y ε 41K
42Ca 0.647% stable
43Ca 0.135% stable
44Ca 2.09% stable
45Ca synth 162.61 d β− 45Sc
46Ca 0.004% stable
47Ca synth 4.536 d β− 47Sc
48Ca 0.187% 5.6×1019 y β−β− 48Ti
 Category: Calcium
| references

Calcium is a chemical element; it has symbol Ca and atomic number 20. As an alkaline earth metal, calcium is a reactive metal that forms a dark oxide-nitride layer when exposed to air. Its physical and chemical properties are most similar to its heavier homologues strontium and barium. It is the fifth most abundant element in Earth's crust, and the third most abundant metal, after iron and aluminium. The most common calcium compound on Earth is calcium carbonate, found in limestone and the fossils of early sea life; gypsum, anhydrite, fluorite, and apatite are also sources of calcium. The name comes from Latin calx "lime", which was obtained from heating limestone.

Some calcium compounds were known to the ancients, though their chemistry was unknown until the seventeenth century. Pure calcium was isolated in 1808 via electrolysis of its oxide by Humphry Davy, who named the element. Calcium compounds are widely used in many industries: in foods and pharmaceuticals for calcium supplementation, in the paper industry as bleaches, as components in cement and electrical insulators, and in the manufacture of soaps. On the other hand, the metal in pure form has few applications due to its high reactivity; still, in small quantities it is often used as an alloying component in steelmaking, and sometimes, as a calcium–lead alloy, in making automotive batteries.

Calcium is the most abundant metal and the fifth-most abundant element in the human body.[8] As electrolytes, calcium ions (Ca2+) play a vital role in the physiological and biochemical processes of organisms and cells: in signal transduction pathways where they act as a second messenger; in neurotransmitter release from neurons; in contraction of all muscle cell types; as cofactors in many enzymes; and in fertilization.[8] Calcium ions outside cells are important for maintaining the potential difference across excitable cell membranes, protein synthesis, and bone formation.[8][9]

Characteristics

[edit]

Classification

[edit]
Calcium crystals stored in mineral oil

Calcium is a very ductile silvery metal (sometimes described as pale yellow) whose properties are very similar to the heavier elements in its group, strontium, barium, and radium. A calcium atom has 20 electrons, with electron configuration [Ar]4s2. Like the other elements in group 2 of the periodic table, calcium has two valence electrons in the outermost s-orbital, which are very easily lost in chemical reactions to form a dipositive ion with the stable electron configuration of a noble gas, in this case argon.[10]

Hence, calcium is almost always divalent in its compounds, which are usually ionic. Hypothetical univalent salts of calcium would be stable with respect to their elements, but not to disproportionation to the divalent salts and calcium metal, because the enthalpy of formation of MX2 is much higher than those of the hypothetical MX. This occurs because of the much greater lattice energy afforded by the more highly charged Ca2+ cation compared to the hypothetical Ca+ cation.[10]

Calcium, strontium, barium, and radium are always considered to be alkaline earth metals; the lighter beryllium and magnesium, also in group 2 of the periodic table, are often included as well. Nevertheless, beryllium and magnesium differ significantly from the other members of the group in their physical and chemical behavior: they behave more like aluminium and zinc respectively and have some of the weaker metallic character of the post-transition metals, which is why the traditional definition of the term "alkaline earth metal" excludes them.[11]

Physical properties

[edit]

Calcium metal melts at 842 °C and boils at 1494 °C; these values are higher than those for magnesium and strontium, the neighbouring group 2 metals. It crystallises in the face-centered cubic arrangement like strontium and barium; above 443 °C (716 K), it changes to body-centered cubic.[4][12] Its density of 1.526 g/cm3 (at 20 °C)[4] is the lowest in its group.[10]

Calcium is harder than lead but can be cut with a knife with effort. While calcium is a poorer conductor of electricity than copper or aluminium by volume, it is a better conductor by mass than both due to its very low density.[13] While calcium is infeasible as a conductor for most terrestrial applications as it reacts quickly with atmospheric oxygen, its use as such in space has been considered.[13]

Chemical properties

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

The chemistry of calcium is that of a typical heavy alkaline earth metal. For example, calcium spontaneously reacts with water more quickly than magnesium but less quickly than strontium to produce calcium hydroxide and hydrogen gas. It also reacts with the oxygen and nitrogen in air to form a mixture of calcium oxide and calcium nitride.[14] When finely divided, it spontaneously burns in air to produce the nitride. Bulk calcium is less reactive: it quickly forms a hydration coating in moist air, but below 30% relative humidity it may be stored indefinitely at room temperature.[15]

Besides the simple oxide CaO, calcium peroxide, CaO2, can be made by direct oxidation of calcium metal under a high pressure of oxygen, and there is some evidence for a yellow superoxide Ca(O2)2.[16]Calcium hydroxide, Ca(OH)2, is a strong base, though not as strong as the hydroxides of strontium, barium or the alkali metals.[17] All four dihalides of calcium are known.[18] Calcium carbonate (CaCO3) and calcium sulfate (CaSO4) are particularly abundant minerals.[19] Like strontium and barium, as well as the alkali metals and the divalent lanthanides europium and ytterbium, calcium metal dissolves directly in liquid ammonia to give a dark blue solution.[20]

Due to the large size of the calcium ion (Ca2+), high coordination numbers are common, up to 24 in some intermetallic compounds such as CaZn13.[21] Calcium is readily complexed by oxygen chelates such as EDTA and polyphosphates, which are useful in analytic chemistry and removing calcium ions from hard water. In the absence of steric hindrance, smaller group 2 cations tend to form stronger complexes, but when large polydentate macrocycles are involved the trend is reversed.[19]

Organocalcium compounds

[edit]
Main article: Organocalcium chemistry

In contrast to organomagnesium compounds, organocalcium compounds are not similarly useful, with one major exception, calcium carbide, CaC2. This material, which has historic significance, is prepared by heating calcium oxide with carbon. According to X-ray crystallography, calcium carbide can be described as Ca2+ derivative of acetylide, C22-, although it is not a salt. Several million tons of calcium carbide are produced annually. Hydrolysis gives acetylene, which is used in welding and a chemical precursor. Reaction with nitrogen gas converts calcium carbide to calcium cyanamide.[22]

A dominant theme in molecular organocalcium chemistry is the large radius of calcium, which often leads to high coordination numbers. For example, dimethylcalcium appears to be a 3-dimensional polymer,[23] whereas dimethylmagnesium is a linear polymer with tetrahedral Mg centers. Bulky ligands are often required to disfavor polymeric species. For example, calcium dicyclopentadienyl, Ca(C5H5)2 has a polymeric structure and thus is nonvolatile and insoluble in solvents. Replacing the C5H5 ligand with the bulkier C5(CH3)5 (pentamethylcyclopentadienyl) gives a soluble complex that sublimes and forms well-defined adducts with ethers.[19] Organocalcium compounds tend to be more similar to organoytterbium compounds due to the similar ionic radii of Yb2+ (102 pm) and Ca2+ (100 pm).[24]

Organocalcium compounds have been well investigated. Some such complexes exhibit catalytic properties,[25] although none have been commercialized.

Isotopes

[edit]
Main article: Isotopes of calcium

Natural calcium is a mixture of five stable isotopes—40Ca, 42Ca, 43Ca, 44Ca, and 46Ca—and 48Ca, whose half-life of 4.3 × 1019 years is so long that it can be considered stable for all practical purposes. Calcium is the first (lightest) element to have six naturally occurring isotopes.[14]

By far the most common isotope is 40Ca, which makes up 96.941% of natural calcium. It is produced in the silicon-burning process from fusion of alpha particles and is the heaviest stable nuclide with equal proton and neutron numbers; its occurrence is also supplemented slowly by the decay of primordial 40K. Adding another alpha particle leads to unstable 44Ti, which decays via two successive electron captures to stable 44Ca; this makes up 2.806% of natural calcium and is the second-most common isotope.[26][27]

The other four natural isotopes, 42, 43, 46, 48Ca, are significantly rarer, each comprising less than 1% of natural calcium. The four lighter isotopes are mainly products of oxygen-burning and silicon-burning, leaving the two heavier ones to be produced via neutron capture. 46Ca is mostly produced in a "hot" s-process, as its formation requires a rather high neutron flux to allow short-lived 45Ca to capture a neutron. 48Ca is produced by electron capture in the r-process in type Ia supernovae, where high neutron excess and low enough entropy ensures its survival.[26][27]

46Ca and 48Ca are the first "classically stable" nuclides with a 6-neutron or 8-neutron excess respectively. Though extremely neutron-rich for such a light element, 48Ca is very stable because it is a doubly magic nucleus, with 20 protons and 28 neutrons arranged in closed shells. Its beta decay to 48Sc is very hindered by the gross mismatch of nuclear spin: 48Ca has zero nuclear spin, being even–even, while 48Sc has spin 6+, so the decay is forbidden by conservation of angular momentum. While two excited states of 48Sc are available for decay as well, they are also forbidden due to their high spins. As a result, when 48Ca does decay, it does so by double beta decay to 48Ti instead, being the lightest nuclide known to undergo double beta decay.[28][29]

46Ca can also theoretically double-beta-decay to 46Ti, but this has never been observed. The most common isotope 40Ca is also doubly magic and could undergo double electron capture to 40Ar, but this has likewise never been observed. Calcium is the only element with two primordial doubly magic isotopes. The experimental lower limits for the half-lives of 40Ca and 46Ca are 5.9 × 1021 years and 2.8 × 1015 years respectively.[28]

Excluding 48Ca, the longest lived radioisotope of calcium is 41Ca. It decays by electron capture to stable 41K with a half-life of about 105 years. Its existence in the early Solar System as an extinct radionuclide has been inferred from excesses of 41K. Traces of 41Ca also still exist today, as it is a cosmogenic nuclide, continuously produced through neutron activation of natural 40Ca.[27]

Many other calcium radioisotopes are known, ranging from 35Ca to 60Ca. They are all much shorter-lived than 41Ca; the most stable are 45Ca (half-life 163 days) and 47Ca (half-life 4.54 days). Isotopes lighter than 42Ca usually undergo beta plus decay to isotopes of potassium, and those heavier than 44Ca usually undergo beta minus decay to scandium; though near the nuclear drip lines, proton emission and neutron emission begin to be significant decay modes as well.[28]

Like other elements, a variety of processes alter the relative abundance of calcium isotopes.[30] The best studied of these processes is the mass-dependent fractionation of calcium isotopes that accompanies the precipitation of calcium minerals such as calcite, aragonite and apatite from solution. Lighter isotopes are preferentially incorporated into these minerals, leaving the surrounding solution enriched in heavier isotopes at a magnitude of roughly 0.025% per atomic mass unit (amu) at room temperature. Mass-dependent differences in calcium isotope composition are conventionally expressed by the ratio of two isotopes (usually 44Ca/40Ca) in a sample compared to the same ratio in a standard reference material. 44Ca/40Ca varies by about 1–2‰ among organisms on Earth.[31]

History

[edit]
One of the 'Ain Ghazal Statues, made from lime plaster

Calcium compounds were known for millennia, though their chemical makeup was not understood until the 17th century.[32] Lime as a building material[33] and as plaster for statues was used as far back as around 7000 BC.[34] The first dated lime kiln dates back to 2500 BC and was found in Khafajah, Mesopotamia.[35][36]

About the same time, dehydrated gypsum (CaSO4·2H2O) was being used in the Great Pyramid of Giza. This material would later be used for the plaster in the tomb of Tutankhamun. The ancient Romans instead used lime mortars made by heating limestone (CaCO3). The name "calcium" itself derives from the Latin word calx "lime".[32]

Vitruvius noted that the lime that resulted was lighter than the original limestone, attributing this to the boiling of the water. In 1755, Joseph Black proved that this was due to the loss of carbon dioxide, which as a gas had not been recognised by the ancient Romans.[37]

In 1789, Antoine Lavoisier suspected that lime might be an oxide of an element. In his table of the elements, Lavoisier listed five "salifiable earths" (i.e., ores that could be made to react with acids to produce salts (salis = salt, in Latin): chaux (calcium oxide), magnésie (magnesia, magnesium oxide), baryte (barium sulfate), alumine (alumina, aluminium oxide), and silice (silica, silicon dioxide)). About these "elements", Lavoisier reasoned:

We are probably only acquainted as yet with a part of the metallic substances existing in nature, as all those which have a stronger affinity to oxygen than carbon possesses, are incapable, hitherto, of being reduced to a metallic state, and consequently, being only presented to our observation under the form of oxyds, are confounded with earths. It is extremely probable that barytes, which we have just now arranged with earths, is in this situation; for in many experiments it exhibits properties nearly approaching to those of metallic bodies. It is even possible that all the substances we call earths may be only metallic oxyds, irreducible by any hitherto known process.[38]

Calcium, along with its congeners magnesium, strontium, and barium, was first isolated by Humphry Davy in 1808. Following the work of Jöns Jakob Berzelius and Magnus Martin of Pontin on electrolysis, Davy isolated calcium and magnesium by putting a mixture of the respective metal oxides with mercury(II) oxide on a platinum plate which was used as the anode, the cathode being a platinum wire partially submerged into mercury. Electrolysis then gave calcium–mercury and magnesium–mercury amalgams, and distilling off the mercury gave the metal.[32][39] However, pure calcium cannot be prepared in bulk by this method and a workable commercial process for its production was not found until over a century later.[37]

Occurrence and production

[edit]
Travertine terraces in Pamukkale, Turkey

At 3%, calcium is the fifth most abundant element in the Earth's crust, and the third most abundant metal behind aluminium and iron.[40] It is also the fourth most abundant element in the lunar highlands.[15] Sedimentary calcium carbonate deposits pervade the Earth's surface as fossilised remains of past marine life; they occur in two forms, the rhombohedral calcite (more common) and the orthorhombic aragonite (forming in more temperate seas). Minerals of the first type include limestone, dolomite, marble, chalk, and Iceland spar; aragonite beds make up the Bahamas, the Florida Keys, and the Red Sea basins. Corals, sea shells, and pearls are mostly made up of calcium carbonate. Among the other important minerals of calcium are gypsum (CaSO4·2H2O), anhydrite (CaSO4), fluorite (CaF2), and apatite ([Ca5(PO4)3X], X = OH, Cl, or F)[32]

The major producers of calcium are China (about 10000 to 12000 tonnes per year), Russia (about 6000 to 8000 tonnes per year), and the United States (about 2000 to 4000 tonnes per year). Canada and France are among the minor producers. In 2005, about 24000 tonnes of calcium were produced; about half of the world's extracted calcium is used by the United States, with about 80% of the output used each year.[13]

In Russia and China, Davy's method of electrolysis is still used, but is instead applied to molten calcium chloride.[13] Since calcium is less reactive than strontium or barium, the oxide–nitride coating that results in air is stable and lathe machining and other standard metallurgical techniques are suitable for calcium.[41]

In the U.S. and Canada, calcium is instead produced by reducing lime with aluminium at high temperatures.[13] In this process, powdered high-calcium lime and powdered aluminum are mixed and compacted into briquettes for a high degree of contact, which are then placed in a sealed retort which has been evacuated and heated to ~1200°C.[13] The briquettes release calcium vapor into the vacuum for about 8 hours, which then condenses in the cooled ends of the retorts to form 24-34 kg pieces of calcium metal, as well as some residue of calcium aluminate.[13] High-purity calcium can be obtained by distilling low-purity calcium at high temperatures.[13]

Geochemical cycling

[edit]
Main article: Carbonate–silicate cycle

Calcium cycling provides a link between tectonics, climate, and the carbon cycle. In the simplest terms, mountain-building exposes calcium-bearing rocks such as basalt and granodiorite to chemical weathering and releases Ca2+ into surface water. These ions are transported to the ocean where they react with dissolved CO2 to form limestone (CaCO3), which in turn 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).[42]

The actual reaction is more complicated and involves the bicarbonate ion (HCO−
3) that forms when CO2 reacts with water at seawater pH:

Ca2+ + 2 HCO−3 → CaCO3↓ + CO2 + H2O

At seawater pH, most of the dissolved 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.[43] The result is that each Ca2+ ion released by chemical weathering ultimately removes one CO2 molecule from the surficial system (atmosphere, ocean, soil and living organisms), storing it in carbonate rocks where it is likely to stay for hundreds of millions of years. The weathering of calcium from rocks thus scrubs CO2 from the ocean and air, exerting a strong long-term effect on climate.[42][44]

Applications

[edit]

The largest use of metallic calcium is in steelmaking, due to its strong chemical affinity for chalcogens oxygen and sulfur. Its oxides and sulfides, once formed, give liquid lime aluminate and sulfide inclusions in steel which float out; on treatment, these inclusions disperse throughout the steel and become small and spherical, improving castability, cleanliness and general mechanical properties. Calcium is also used in maintenance-free automotive batteries, in which the use of 0.1% calcium–lead alloys instead of the usual antimony–lead alloys leads to lower water loss and lower self-discharging.[45]

Due to the risk of expansion and cracking, aluminium is sometimes also incorporated into these alloys. These lead–calcium alloys are also used in casting, replacing lead–antimony alloys.[45] Calcium is also used to strengthen aluminium alloys used for bearings, for the control of graphitic carbon in cast iron, and to remove bismuth impurities from lead.[41] 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 (for example, those in hair) that block drains.[46]

Besides metallurgy, the reactivity of calcium is exploited to remove nitrogen from high-purity argon gas and as a getter for oxygen and nitrogen. It is also used as a reducing agent in the production of chromium, zirconium, thorium, vanadium and uranium. It can also be used to store hydrogen gas, as it reacts with hydrogen to form solid calcium hydride, from which the hydrogen can easily be re-extracted.[41]

Calcium isotope fractionation during mineral formation has led to several applications of calcium isotopes. In particular, the 1997 observation by Skulan and DePaolo[47] that calcium minerals are isotopically lighter than the solutions from which the minerals precipitate is the basis of analogous applications in medicine and in paleoceanography. In animals with skeletons mineralised with calcium, the calcium isotopic composition of soft tissues reflects the relative rate of formation and dissolution of skeletal mineral.[48]

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 44Ca/40Ca ratio in soft tissue rises and vice versa. Because of this relationship, calcium isotopic measurements of urine or blood may be useful in the early detection of metabolic bone diseases like osteoporosis.[48]

A similar system exists in seawater, where 44Ca/40Ca tends to rise when the rate of removal of Ca2+ by mineral precipitation exceeds the input of new calcium into the ocean. In 1997, Skulan and DePaolo presented the first evidence of change in seawater 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 is never in a "steady state" with respect to calcium input and output. This has important climatological implications, as the marine calcium cycle is closely tied to the carbon cycle.[49][50]

Many calcium compounds are used in food, as pharmaceuticals, and in medicine, among others. For example, calcium and phosphorus are supplemented in foods through the addition of calcium lactate, calcium diphosphate, and tricalcium phosphate. The last is also used as a polishing agent in toothpaste and in antacids. Calcium lactobionate is a white powder that is used as a suspending agent for pharmaceuticals. In baking, calcium phosphate is used as a leavening agent. Calcium sulfite is used as a bleach in papermaking and as a disinfectant, calcium silicate is used as a reinforcing agent in rubber, and calcium acetate is a component of liming rosin and is used to make metallic soaps and synthetic resins.[45]

Calcium supplement is on the World Health Organization's List of Essential Medicines.[51]

Food sources

[edit]
See also: Calcium supplement

Foods rich in calcium include dairy products such as milk, yogurt, and cheese, as well as sardines, salmon, soy products, kale, and fortified breakfast cereals.[9]

Because of concerns for long-term adverse side effects, including calcification of arteries and kidney stones, both the U.S. Institute of Medicine (IOM) and the European Food Safety Authority (EFSA) set tolerable upper intake levels (ULs) for combined dietary and supplemental calcium. From the IOM, people of ages 9–18 years are not to exceed 3 g/day combined intake; for ages 19–50, not to exceed 2.5 g/day; for ages 51 and older, not to exceed 2 g/day.[52] EFSA set the UL for all adults at 2.5 g/day, but decided the information for children and adolescents was not sufficient to determine ULs.[53]

Biological and pathological role

[edit]
Main article: Calcium in biology
Age-adjusted daily calcium recommendations (from U.S. Institute of Medicine RDAs)[54]
Age Calcium (mg/day)
1–3 years 700
4–8 years 1000
9–18 years 1300
19–50 years 1000
>51 years 1000
Pregnancy 1000
Lactation 1000
Global dietary calcium intake among adults (mg/day).[55]
  <400
  400–500
  500–600
  600–700
  700–800
  800–900
  900–1000
  >1000

Function

[edit]

Calcium is an essential element needed in large quantities.[8][9] The Ca2+ ion acts as an electrolyte and is vital to the health of the muscular, circulatory, and digestive systems; is indispensable to the building of bone in the form of hydroxyapatite; and supports synthesis and function of blood cells. For example, it regulates the contraction of muscles, nerve conduction, and the clotting of blood. As a result, intra- and extracellular calcium levels are tightly regulated by the body. Calcium can play this role because the Ca2+ ion forms stable coordination complexes with many organic compounds, especially proteins; it also forms compounds with a wide range of solubilities, enabling the formation of the skeleton.[8] [56]

Binding

[edit]

Calcium ions may be complexed by proteins through binding the carboxyl groups of glutamic acid or aspartic acid residues; through interacting with phosphorylated serine, tyrosine, or threonine residues; or by being chelated by γ-carboxylated amino acid residues. Trypsin, a digestive enzyme, uses the first method; osteocalcin, a bone matrix protein, uses the third.[57]

Some other bone matrix proteins such as osteopontin and bone sialoprotein use both the first and the second. Direct activation of enzymes by binding calcium is common; some other enzymes are activated by noncovalent association with direct calcium-binding enzymes. Calcium also binds to the phospholipid layer of the cell membrane, anchoring proteins associated with the cell surface.[57]

Solubility

[edit]

As an example of the wide range of solubility of calcium compounds, monocalcium phosphate is very soluble in water, 85% of extracellular calcium is as dicalcium phosphate with a solubility of 2.00 mM, and the hydroxyapatite of bones in an organic matrix is tricalcium phosphate with a solubility of 1000 μM.[57]

Nutrition

[edit]

Calcium is a common constituent of multivitamin dietary supplements,[8] but the composition of calcium complexes in supplements may affect its bioavailability which varies by solubility of the salt involved: calcium citrate, malate, and lactate are highly bioavailable, while the oxalate is less. Other calcium preparations include calcium carbonate, calcium citrate malate, and calcium gluconate.[8] The intestine absorbs about one-third of calcium eaten as the free ion, and plasma calcium level is then regulated by the kidneys.[8]

Hormonal regulation of bone formation and serum levels

[edit]

Parathyroid hormone and vitamin D promote the formation of bone by allowing and enhancing the deposition of calcium ions there, allowing rapid bone turnover without affecting bone mass or mineral content.[8] When plasma calcium levels fall, cell surface receptors are activated and the secretion of parathyroid hormone occurs; it then proceeds to stimulate the entry of calcium into the plasma pool by taking it from targeted kidney, gut, and bone cells, with the bone-forming action of parathyroid hormone being antagonised by calcitonin, whose secretion increases with increasing plasma calcium levels.[57]

Abnormal serum levels

[edit]

Excess intake of calcium may cause hypercalcemia. However, because calcium is absorbed rather inefficiently by the intestines, high serum calcium is more likely caused by excessive secretion of parathyroid hormone (PTH) or possibly by excessive intake of vitamin D, both of which facilitate calcium absorption. All these conditions result in excess calcium salts being deposited in the heart, blood vessels, or kidneys. Symptoms include anorexia, nausea, vomiting, memory loss, confusion, muscle weakness, increased urination, dehydration, and metabolic bone disease.[57]

Chronic hypercalcaemia typically leads to calcification of soft tissue and its serious consequences: for example, calcification can cause loss of elasticity of vascular walls and disruption of laminar blood flow—and thence to plaque rupture and thrombosis. Conversely, inadequate calcium or vitamin D intakes may result in hypocalcemia, often caused also by inadequate secretion of parathyroid hormone or defective PTH receptors in cells. Symptoms include neuromuscular excitability, which potentially causes tetany and disruption of conductivity in cardiac tissue.[57]

Bone disease

[edit]

As calcium is required for bone development, many bone diseases can be traced to the organic matrix or the hydroxyapatite in molecular structure or organization of bone. Osteoporosis is a reduction in mineral content of bone per unit volume, and can be treated by supplementation of calcium, vitamin D, and bisphosphonates.[8][9] Inadequate amounts of calcium, vitamin D, or phosphates can lead to softening of bones, called osteomalacia.[57]

Safety

[edit]

Metallic calcium

[edit]
Calcium
Hazards
GHS labelling:[58]
GHS02: Flammable
Danger
H261
P231+P232
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 0: Exposure under fire conditions would offer no hazard beyond that of ordinary combustible material. E.g. sodium chlorideFlammability 3: Liquids and solids that can be ignited under almost all ambient temperature conditions. Flash point between 23 and 38 °C (73 and 100 °F). E.g. gasolineInstability 1: Normally stable, but can become unstable at elevated temperatures and pressures. E.g. calciumSpecial hazard W: Reacts with water in an unusual or dangerous manner. E.g. sodium, sulfuric acid
0
3
1
W

Because calcium reacts exothermically with water and acids, calcium metal coming into contact with bodily moisture results in severe corrosive irritation.[59] When swallowed, calcium metal has the same effect on the mouth, oesophagus, and stomach, and can be fatal.[46] However, long-term exposure is not known to have distinct adverse effects.[59]

References

[edit]

Bibliography

[edit]
Revisions and contributorsEdit on WikipediaRead on Wikipedia

Calcium

View on Grokipedia
from Grokipedia
Calcium is a chemical element with the symbol Ca and atomic number 20, classified as a soft, silvery-white alkaline earth metal in group 2 of the periodic table. It has an atomic mass of 40.078 u, an electron configuration of [Ar] 4s², and occurs naturally in compounds such as limestone (calcium carbonate) and gypsum (calcium sulfate dihydrate), making it the fifth-most-abundant element by mass in Earth's crust at over 3%. Essential for life, calcium plays critical roles in biomineralization for bones, teeth, and shells, as well as in cellular signaling, muscle contraction, and nerve transmission. Physically, calcium is a solid at with a of 1.54 g/cm³, a of 842°C (1115 K), and a of 1484°C (1757 K). The metal tarnishes in air due to oxidation but is protected by a thin layer, and it reacts vigorously with to produce gas and . Calcium was first isolated in its metallic form in 1808 by English chemist through the of a mixture of lime (CaO) and mercuric , though its compounds like lime had been used since ancient times for and ; the name derives from the Latin calx, meaning lime. In nature, calcium is never found uncombined but constitutes key minerals including (CaF₂) and (calcium phosphate), with significant abundances in (about 4220 ppm by weight) and the (around 1.4% by weight, primarily in bones as ). Biologically, over 99% of bodily calcium is stored in the for and as a reservoir, while the remaining 1% in and tissues regulates processes like blood clotting, secretion, and activation; ionized calcium in serum normally ranges from 8.8 to 10.4 mg/dL. Deficiency can lead to conditions like , underscoring its indispensable role in health. Industrially, calcium metal serves as a in extracting metals like and , and as an alloying agent to strengthen aluminum, , and ; its compounds dominate applications, including lime for production, , and manufacturing. Calcium isotopes, such as ⁴²Ca and ⁴⁶Ca, are employed in nutritional to study absorption and . Ongoing explores calcium in advanced batteries for its abundance and potential for high-energy density storage.

Characteristics

Classification

Calcium is a with 20, positioned in group 2 and period 4 of the periodic table, classifying it as an . Its ground-state is [Ar] 4s², with the two valence electrons occupying the 4s orbital, which contributes to its characteristic chemical behavior within the group. Alkaline earth metals like calcium are highly reactive, owing to their tendency to lose both valence electrons and form divalent cations (M²⁺), resulting in a +2 in most compounds./08:_Chemistry_of_the_Main_Group_Elements/8.05:_Group_2_The_Alkaline_Earth_Metals/8.5.02:_Alkaline_Earth_Metals'_Chemical_Properties) This reactivity is enhanced by their low ; calcium specifically has an electronegativity of 1.00 on the Pauling scale, making it prone to with nonmetals.80142-5) Compared to other group 2 elements, calcium's atomic radius of 197 pm is intermediate, larger than magnesium (160 pm) but smaller than strontium (215 pm), illustrating the trend of increasing size down the group due to additional electron shells. Its ionization energies reflect this position: the first is 589.8 kJ/mol and the second is 1145 kJ/mol, values higher than magnesium's (738 kJ/mol first, 1451 kJ/mol second) but lower than strontium's (549 kJ/mol first, 1064 kJ/mol second), consistent with decreasing ionization energy as atomic size increases.

Physical properties

Calcium is a soft, silvery-white metal that rapidly tarnishes in air to form a grayish-white layer. This malleable and ductile nature allows it to be easily shaped or drawn into wires, characteristic of its position among the alkaline earth metals. The density of elemental calcium is 1.55 g/cm³ at 20 °C. It has a of 842 °C (1115 K) and a of 1484 °C (1757 K). Calcium exhibits a face-centered cubic (FCC) crystal structure with a of 558.8 pm at 20 °C. Elemental calcium demonstrates good thermal conductivity of 201 W/(m·K) and electrical resistivity of 34 nΩ·m at 20 °C, reflecting its and free electron behavior. These properties contribute to its utility in applications requiring heat dissipation or electrical conduction, though its reactivity limits direct use.

Chemical properties

Calcium is a highly reactive that tarnishes rapidly in air, forming a gray-white coating of (CaO) and calcium (Ca₃N₂) due to its affinity for and . Finely divided calcium is pyrophoric and ignites spontaneously in air, burning with an intense white flame. It reacts vigorously with at , liberating gas and forming (Ca(OH)₂), as represented by the equation: \ceCa+2H2O−>Ca(OH)2+H2\ce{Ca + 2H2O -> Ca(OH)2 + H2} This exothermic reaction demonstrates calcium's strong reducing properties. The predominant oxidation state of calcium is +2, resulting from the facile loss of its two 4s valence electrons, enabled by relatively low successive ionization energies (first: 589.8 kJ/mol; second: 1145.3 kJ/mol). This electron configuration promotes ionic bonding in calcium compounds, where the Ca²⁺ cation electrostatically interacts with anions, favoring the formation of ionic lattices over covalent structures. Calcium's position in the reactivity series places it above magnesium but below sodium, underscoring its ability to displace hydrogen from dilute acids, yielding hydrogen gas and soluble calcium salts. Calcium readily forms compounds such as oxides (e.g., CaO), halides (e.g., CaCl₂), and sulfates (e.g., ), with solubility trends governed by and hydration effects. Most calcium salts exhibit good solubility in , but notable exceptions include (CaCO₃), which is insoluble (K_{sp} = 2.8 \times 10^{-9}), and calcium sulfate dihydrate (, ), which has limited (K_{sp} = 2.4 \times 10^{-5} at 25°C). The standard reduction potential for the Ca²⁺/Ca half-cell is -2.87 V versus the standard hydrogen electrode, indicating calcium's potent reducing capability and thermodynamic tendency to oxidize.

Isotopes

Calcium has 25 known isotopes, with mass numbers ranging from 34 to 58. Six of these isotopes are and occur naturally on , while the remainder are radioactive with half-lives spanning from microseconds to over 100,000 years. The isotopes dominate the natural composition of calcium, with variations in abundance reflecting both primordial and minor cosmogenic production. The isotopes of calcium, along with their approximate natural abundances in terrestrial materials, are listed below:
IsotopeNatural Abundance (%)
⁴⁰Ca96.941
⁴²Ca0.647
⁴³Ca0.135
⁴⁴Ca2.086
⁴⁶Ca0.004
⁴⁸Ca0.187
These abundances are based on measurements from standard geological samples and show minor variations due to mass-dependent processes in natural systems. Among the radioactive isotopes, ⁴¹Ca has a half-life of 99,400 years and is produced cosmogenically through neutron capture on ⁴⁰Ca; it serves as a long-lived tracer in studies of calcium metabolism and has applications in paleodietary research by enabling dating of ancient bone samples beyond the range of radiocarbon methods. Another notable example is ⁴⁵Ca, with a half-life of 162.7 days, which is widely used as a radioactive tracer in biological and medical investigations of calcium uptake and transport in living organisms. The other radioactive isotopes decay primarily via beta emission or electron capture, with most having short half-lives that limit their natural occurrence. The stable isotopes of calcium are primarily synthesized through , where explosive stellar events produce neutron-rich nuclei via rapid neutron capture (r-process) and silicon burning, supplemented by contributions from asymptotic giant branch stars through the slow neutron capture (s-process). Calcium isotopes find specialized applications in and . For instance, ⁴⁴Ca is employed in theoretical and experimental studies of , a rare process that could reveal insights into properties if observed. Unlike heavier elements, calcium has no fissile isotopes capable of sustaining a .

History

Discovery and isolation

Lime, or (CaO), has been utilized as a compound since prehistoric times, with evidence of its use in dating back to approximately 4000 BCE in , where it served as a binding mortar and in the building of early pyramids, such as those at around 2600 BCE. This early recognition of lime's properties as a reactive material for mortar highlights its role in foundational , though the elemental nature of calcium remained unrecognized for millennia. Efforts to isolate calcium as a pure element began in the early amid advances in . In 1803, Swedish chemists and Magnus Martin Pontin conducted of lime dissolved in mercury, producing a calcium-mercury amalgam but failing to obtain the pure metal due to the amalgam's stability. Inspired by their work, British chemist pursued similar experiments at the Royal Institution in . Davy successfully isolated metallic calcium in 1808 through electrolysis of a mixture of moist lime (CaO) and mercuric (HgO), using a mercury to form a calcium amalgam, which he then heated in a to distill away the mercury and yield impure calcium metal. This breakthrough, announced in a series of lectures to the Royal Society, marked the first production of calcium in its elemental form and demonstrated the power of for decomposing refractory s, a technique Davy had previously applied to and sodium. The resulting calcium appeared as a soft, silvery-white solid that rapidly tarnished in air, confirming its reactivity and placement among the alkaline earth metals.

Etymology and early recognition

The term "calcium" originates from the Latin word , meaning lime or limestone, reflecting the element's long association with calcium oxide (CaO), commonly known as quicklime. This nomenclature was proposed by British chemist in 1808, following his electrolytic isolation of the metal from a mixture of lime and mercuric oxide, to denote the new element distinct from earlier uses of "calx" for calcined substances like ashes. Prior to this, in alchemical and early chemical traditions, calcium compounds were referred to as "calcareous earth" or "earth of lime," encompassing materials such as limestone and chalk that were manipulated in processes like calcination to produce lime for various applications. Human recognition of calcium-based materials dates back millennia, with evidence of their use in construction predating the element's identification. (CaSO₄·2H₂O), a hydrated calcium sulfate, was employed as plaster in settlements, notably at in modern-day around 7500 BCE, where it was mixed with lime for wall coatings and decorative purposes. By the Roman era, quicklime played a pivotal role in durable formulations; for instance, the Pantheon's massive dome, completed in 126 CE under Emperor , utilized pozzolanic incorporating quicklime, , and aggregates, enabling self-healing properties through hot-mixing techniques that formed lime clasts for crack repair. In the pre-modern scientific framework, calcium was understood as an elemental "earth" rather than a metal. French , in his 1789 treatise Traité Élémentaire de Chimie, classified "terre calcaire" (calcareous earth) among the simple substances, listing it alongside other earths like magnesia and baryta based on their resistance to decomposition and role in forming salts. Swedish , who assisted Davy in early experiments, further refined the terminology in 1810 by adopting and promoting "calcium" in , aligning it with emerging conventions for metallic elements derived from earths. This progression from alchemical "earth of lime" to a recognized element underscored calcium's foundational role in both practical arts and systematic chemistry.

Occurrence and production

Natural occurrence

Calcium is produced primarily through stellar nucleosynthesis in massive stars and is relatively abundant in the universe compared to heavier elements. Its cosmic distribution reflects contributions from Big Bang nucleosynthesis for lighter elements and subsequent stellar processes for heavier ones like calcium. On Earth, calcium ranks as the fifth most abundant element in the crust, comprising approximately 3.6% by mass. In the oceans, it is present at a concentration of about 0.4 g/L primarily as Ca²⁺ ions, dissolved from continental weathering and contributing to marine chemistry. Calcium occurs naturally in a variety of minerals, with (CaCO₃) being the most prevalent, accounting for roughly 4% of the and forming extensive deposits. Other significant minerals include dolomite (CaMg(CO₃)₂), which is common in sedimentary rocks; (CaSO₄·2H₂O), found in deposits; and (CaF₂), a key source of . In the biosphere, calcium is highly concentrated in the exoskeletons, shells, and bones of organisms, where it forms structural components like in marine shells and in vertebrate bones. Marine organisms, such as mollusks, corals, and , extract calcium from to build these structures, facilitating its distribution across ecosystems. Within the solar system, calcium is a major component of chondritic meteorites, representing primitive solar material with typical CaO contents around 1–2%, and lunar rocks, particularly anorthosites in the highlands, where CaO levels can reach up to 10–15%.

Industrial production

The primary industrial method for producing elemental calcium is the aluminothermic reduction of (CaO), a process conducted under at approximately 1200°C. In this reaction, aluminum serves as the , following the equation: 3CaO+2Al→3Ca+Al2O33\text{CaO} + 2\text{Al} \rightarrow 3\text{Ca} + \text{Al}_2\text{O}_3 This yields calcium metal that can be further processed for higher purity. An alternative method, electrolysis of molten (CaCl₂), has been employed since the 1890s and remains in use for producing calcium of about 99% purity. The process involves electrolyzing the molten with a sacrificial , where calcium deposits at the and gas is evolved at the . Global annual production of calcium metal is approximately 10,000 tons as of 2023, with the majority originating from and . Following initial production, calcium is purified by to achieve 99.9% purity, removing volatile impurities and residual oxides. The raw material for these processes is typically derived from the of (CaCO₃), which decomposes to form CaO and releases CO₂ as a : CaCO₃ → CaO + CO₂.

Geochemical cycling

Calcium plays a central role in the carbonate-silicate cycle, a fundamental geochemical process that regulates Earth's long-term and climate stability. During chemical of calcium-rich silicate rocks, such as plagioclase feldspars, derived from atmospheric CO₂ reacts with minerals to release Ca²⁺ ions into solution, which are then transported to the oceans via rivers. In marine settings, these ions combine with (HCO₃⁻) to form (CaCO₃), primarily through biogenic by organisms like coccolithophores and . This effectively sequesters carbon, with global riverine inputs delivering approximately 0.65 Gt of Ca per year to the oceans from both riverine and minor hydrothermal sources. The reverse process occurs through volcanic and metamorphic , releasing CO₂ and enabling renewed . In the oceanic realm, the calcium cycle involves biological uptake, , and physical mixing. Calcifying incorporate Ca²⁺ into CaCO₃ structures, which sink as particulate matter and accumulate as sediments, forming vast deposits over geological time. A significant portion—over 80%—of produced CaCO₃ dissolves in the or upper sediments due to increasing pressure and decreasing with depth, but the net burial balances inputs at roughly 0.6 Gt Ca per year. circulates dissolved calcium back to surface waters, sustaining in nutrient-limited regions. The average of calcium in is approximately 1 million years, reflecting its conservative behavior and well-mixed distribution throughout the global ocean. Volcanic processes contribute to calcium inputs primarily through the formation and subsequent of at mid-ocean ridges and arcs, where CaO constitutes about 10-12% of composition. The from these sources, including hydrothermal alteration, is estimated at around 0.1 Gt Ca per year, providing a steady supply that supports the cycle's continuity. On land, the influences local calcium dynamics: absorb Ca²⁺ from soils for structural roles in cell walls, while herbivores and decomposers recycle it through and decay, with minimal net loss over short timescales but significant observable in isotopic signatures. Human activities have introduced perturbations to this ancient cycle, notably through intensified mining of and the production of , which consumes over 4 Gt of CaCO₃ annually and releases about 1 Gt of CO₂ per year via (CaCO₃ → CaO + CO₂). This process accelerates calcium mobilization and atmospheric CO₂ enrichment, bypassing natural rates and enhancing the reverse flux in the carbonate-silicate cycle, with potential long-term implications for ocean alkalinity and .

Calcium compounds

Inorganic compounds

Calcium oxide (CaO), commonly known as quicklime, is a white, caustic, alkaline solid produced primarily by the of or other sources. The reaction involves heating to temperatures exceeding 900°C, yielding and gas:
\ceCaCO3−>[Δ]CaO+CO2\ce{CaCO3 ->[Δ] CaO + CO2}.
Quicklime exhibits high reactivity with water, undergoing a strongly exothermic hydration to form (slaked lime):
\ceCaO+H2O−>Ca(OH)2\ce{CaO + H2O -> Ca(OH)2}.
It has a of 2,613°C and is widely used as a key ingredient in production due to its binding properties upon hydration. Calcium carbonate (CaCO₃) occurs naturally in several polymorphic forms, including the stable (rhombohedral structure) and the metastable (orthorhombic structure), both of which are principal components of , , and seashells.
It is sparingly soluble in , with a solubility product constant (Ksp) of approximately 3.7 × 10⁻⁹ at 25°C, which governs its precipitation in aqueous environments and contributes to its role in geological formations like stalactites.
Calcium carbonate decomposes thermally above 840°C to form calcium oxide and , mirroring the preparation of quicklime. Among the calcium halides, (CaCl₂) is highly hygroscopic and deliquescent, readily absorbing moisture from the air to form hydrates such as the dihydrate (CaCl₂·2H₂O).
It is typically prepared by reacting with :
\ceCaCO3+2HCl−>CaCl2+H2O+CO2\ce{CaCO3 + 2HCl -> CaCl2 + H2O + CO2}.
has a of 772°C and is soluble in to the extent of about 74 g/100 mL at 20°C, making it useful in de-icing applications due to its exothermic dissolution.
In contrast, (CaF₂), known as or fluorspar, is nearly insoluble in (Ksp ≈ 3.9 × 10⁻¹¹ at 25°C) and adopts a cubic .
It occurs naturally as a and has a high of 1,418°C, contributing to its use in and for its transparency to radiation. Calcium sulfate exists predominantly as the dihydrate (CaSO₄·2H₂O), or , a soft, white with a Mohs of 2, commonly found in sedimentary deposits.
Upon heating to around 150°C, loses three-quarters of its of hydration to form hemihydrate (CaSO₄·0.5H₂O), known as of , which sets rapidly upon rehydration to regenerate the dihydrate structure.
The dihydrate has low in (about 0.21 g/100 mL at 20°C) and is chemically stable under neutral conditions. Calcium nitrate (Ca(NO₃)₂) is a highly water-soluble salt ( ≈ 121 g/100 mL at 20°C), often encountered as the tetrahydrate (Ca(NO₃)₂·4H₂O), which is colorless and deliquescent.
It can be synthesized by dissolving in or via the reaction of with .
The anhydrous form melts at 561°C and decomposes at higher temperatures, while its solutions are neutral to slightly acidic due to partial .

Organocalcium compounds

Organocalcium compounds, featuring direct calcium-carbon bonds, are exceedingly rare owing to the element's pronounced reactivity, rendering most species highly unstable toward air and moisture; consequently, they must be synthesized and manipulated under strict inert atmospheric conditions. These compounds exhibit bonding that is predominantly ionic, reflecting calcium's electropositive nature, yet displays subtle covalent character in the Ca-C interaction, which contributes to their enhanced nucleophilicity compared to analogous magnesium species. In the gas phase, computational and spectroscopic studies reveal that many such compounds favor dimeric structures, stabilized by bridging ligands or metal-metal interactions. Prominent examples include calcium cyclopentadienide (\ceCp2Ca\ce{Cp2Ca}), a metallocene analog of featuring two η5\eta^5-coordinated rings, first synthesized in the late 1960s by the reaction of calcium metal with in , yielding an insoluble polymeric solid. Another key representative is calcium acetylide (\ceCaC2\ce{CaC2}), formed via the high-temperature reaction of elemental calcium with carbon (typically above 2000°C), which serves as the primary industrial organocalcium and is hydrolyzed to generate gas for . The practical applications of organocalcium compounds remain confined largely to academic research, where their strong nucleophilicity enables roles as initiators in anionic polymerization processes, such as the ring-opening polymerization of lactones or the synthesis of polyolefins, often outperforming traditional Grignard reagents due to reduced side reactions under optimized conditions. Advancements in stabilization have yielded notable complexes like bis[tris(trimethylsilyl)methyl]calcium (\ce[CaC(SiMe3)32]\ce{[Ca{C(SiMe3)3}2]}), prepared in the 1990s through metathesis of calcium diiodide with the corresponding lithium alkyl in hydrocarbon solvents; the enormous steric bulk of the \ce−C(SiMe3)3\ce{-C(SiMe3)3} ligands shields the reactive Ca-C bonds, enabling isolation as a two-coordinate, bent monomeric species stable enough for structural characterization and reactivity studies.

Industrial applications

Metallurgical uses

Calcium plays a critical role in production through desulfurization processes, where it is added to molten in amounts of 0.1–0.3% to react with , forming calcium (CaS) inclusions. These spherical CaS inclusions replace elongated inclusions, which can cause and reduce during ; the modification improves the 's and by preventing crack along inclusion boundaries. This treatment is typically performed in the ladle during secondary metallurgy, using calcium wire or cored wire injection to ensure uniform distribution and minimize losses due to calcium's high reactivity with oxygen. In lead alloys for maintenance-free batteries, calcium is incorporated at levels of 0.03–0.15% in combination with lead-antimony (Pb-Sb) to enhance grid performance. The addition reduces rates by forming stable passivation layers on the grid surfaces, minimizing and water loss, which extends battery life and eliminates the need for maintenance. These Pb-Ca-Sb alloys are into grids for automotive and industrial lead-acid batteries, where the calcium stabilizes the alloy against grid growth during charge-discharge cycles. Calcium also contributes to aluminum alloys through the CaAl₂, which acts as a grain refiner to promote finer equiaxed grains during solidification. This refinement enhances the alloy's mechanical strength, fatigue resistance, and castability by increasing sites and restricting growth, particularly in wrought aluminum alloys used for and automotive components. The addition is controlled to avoid excess calcium, which could lead to brittle phases. In rare earth metal production, calcium serves as a to eliminate fluoride impurities during . High-purity calcium is added to rare earth s (e.g., NdF₃ or DyF₃) in a sealed under , where it preferentially reduces the fluorides to volatile CaF₂, allowing of the pure rare earth metal at temperatures around 1000–1400°C. This calciothermic process is essential for achieving metallic purity above 99.5%, as residual fluorides degrade magnetic properties in applications like permanent magnets.

Chemical and construction uses

Calcium compounds play a pivotal role in construction and chemical industries, particularly through the use of (CaO) and its derivatives. In cement production, is calcined to produce CaO, which constitutes 60–67% of clinker and reacts with silicates, aluminates, and ferrites to form the binding phases essential for hydraulic . Global Portland cement output reached 4.1 gigatons in 2023, underscoring its scale in infrastructure development. Lime, primarily as (Ca(OH)₂), is widely employed in for softening and adjusting . The process involves raising the water's to precipitate hardness-causing ions as (CaCO₃), effectively removing temporary hardness while stabilizing for downstream and control. This application reduces scaling in pipes and improves in municipal systems. In fertilizers, (CaCN₂) serves as a slow-release nitrogen source derived from processes, providing both nitrogen and calcium to enhance and suppress weeds or pathogens. Its use has declined relative to synthetic alternatives like due to handling requirements and slower nutrient release, though it remains valued in niche applications for . Calcium oxide is integral to the paper and sugar industries for purification and bleaching. In papermaking, quicklime (CaO) is used to regenerate caustic soda in the and produce for bleaching pulp, aiding removal and brightness enhancement. In sugar refining, lime clarifies raw juice by neutralizing acids, precipitating impurities, and forming calcium saccharate for filtration, typically at rates of 0.25 tons per ton of sugar. Calcium fluoride (CaF₂) finds application in toothpaste formulations as a mild abrasive, polishing tooth surfaces while releasing fluoride ions to support remineralization and caries prevention. This dual functionality makes it suitable for sensitive enamel formulations.

Biological role

Dietary sources and nutrition

Calcium is an essential mineral obtained primarily through dietary sources, with recommended daily intakes varying by age, sex, and physiological status to support bone health and other functions. According to the National Institutes of Health (NIH), adults aged 19 to 50 years require 1,000 mg of calcium per day, while women over 50 and men over 70 need 1,200 mg; for pregnant and lactating individuals, the recommended dietary allowance (RDA) is 1,300 mg for those aged 14 to 18 and 1,000 mg for ages 19 to 50. These guidelines align with broader international recommendations from organizations like the International Osteoporosis Foundation, emphasizing intakes around 1,000 mg for most adults to prevent deficiency. These elevated requirements for older adults support bone strength and help mitigate age-related bone loss, with calcium often paired with vitamin D to enhance absorption and efficacy in preventing osteoporosis. Healthiest calcium sources prioritize high bioavailability, nutrient density, and minimal added sugars, saturated fats, or processing. Dairy products are primary sources for many people and provide highly bioavailable calcium with approximately 30% absorption. Examples include cow's milk (approximately 300 mg per cup), yogurt (approximately 450 mg per cup), and hard cheeses such as cheddar (approximately 200 mg per ounce). Fortified foods achieve similar absorption rates and include calcium-fortified orange juice (approximately 300 mg per 8 oz) and fortified soy milk (200–400 mg per cup). Non-dairy options encompass calcium-set tofu (250–750 mg per 4 oz, depending on processing with calcium salts), canned fish with edible bones such as sardines (approximately 370 mg per 3 oz) and salmon (170–210 mg per 3 oz), and low-oxalate leafy green vegetables such as (around 150 mg per 100 g), broccoli, and bok choy. Other good sources include almonds and fortified cereals. Foods highest in calcium per typical serving include dairy products and fortified foods. Top sources are:
  • Plain nonfat yogurt: 488 mg per 8 oz
  • Plain low-fat yogurt: 448 mg per 8 oz
  • Fortified almond milk: 442 mg per cup
  • Tofu prepared with calcium sulfate: 434 mg per ½ cup
  • Fortified orange juice: 349 mg per cup
  • Sardines canned with bones: 325 mg per 3 oz
  • Milk (nonfat or low-fat): ~300 mg per cup
Other notable sources include cheese (e.g., mozzarella part-skim: 333 mg per 1.5 oz), certain leafy greens like collard greens (268 mg per cooked cup), and fortified plant milks. Dairy and fortified foods generally provide the most bioavailable calcium. Examples of calcium-rich foods with approximate calcium content per serving include:
  • Yogurt (1 cup): 450 mg
  • Milk (1 cup): 300 mg
  • Fortified orange juice (8 oz): 300 mg
  • Fortified soy milk (1 cup): 200–400 mg
  • Tofu (firm, calcium-set, 4 oz): 250–750 mg
  • Hard cheese (e.g., cheddar, 1 oz): 200 mg
  • Sardines, canned with bones (3 oz): 370 mg
  • Salmon, canned with bones (3 oz): 170–210 mg
Calcium content in fortified products varies, so checking nutrition labels is recommended. Dairy and fortified foods generally offer high bioavailability, while absorption from plant-based sources depends on factors such as oxalate and phytate content. The of dietary calcium, or the fraction absorbed in the intestines, typically ranges from 30% to 40% for sources like and fortified products. Absorption is enhanced by factors such as , which promotes intestinal uptake, and in , which aids solubility. Conversely, inhibitors like phytates (found in grains and ) and oxalates (present in and ) can bind calcium, reducing its absorption by up to 50% in high-fiber diets. For individuals unable to meet requirements through diet, such as (CaCO₃) or are commonly used, with effective doses of 500 to 600 mg taken in divided portions to maximize absorption. , which requires for dissolution, is inexpensive but best taken with meals, whereas citrate is more readily absorbed on an empty . The (WHO) includes on its List, recommending 1.5 to 2 g of elemental calcium daily for pregnant women in regions with low dietary intake to prevent complications like . Globally, inadequate calcium intake affects an estimated 3.5 billion people, particularly in low- and middle-income countries where consumption is limited and plant-based diets predominate. This deficiency heightens the risk of , a condition characterized by reduced and increased fracture susceptibility, underscoring the need for targeted nutritional interventions.

Physiological functions

Calcium plays a fundamental structural role in the , with over 99% of total body calcium stored in bones and teeth in the form of , \ceCa10(PO4)6(OH)2\ce{Ca10(PO4)6(OH)2}, which imparts rigidity and strength to these tissues. In an average adult, this amounts to approximately 1–2 kg of calcium, primarily contributing to the mineralization and mechanical support of skeletal structures. crystals integrate with an organic matrix of and other proteins, forming a that withstands compressive forces and maintains overall body posture. As a key signaling molecule, the calcium ion (\ceCa2+\ce{Ca^2+}) serves as a ubiquitous second messenger in cellular processes, particularly in neurons and muscle cells, where it orchestrates responses to stimuli. In neurons, \ceCa2+\ce{Ca^2+} influx during action potentials triggers release at synapses by binding to proteins like synaptotagmin, while in muscle cells, it is released from (ER) stores via channels such as receptors, amplifying signals for contraction and other functions. These transient elevations in cytosolic \ceCa2+\ce{Ca^2+} concentration, often from 100 nM to 1–10 μM, activate downstream pathways including calmodulin-dependent kinases, ensuring precise temporal and spatial control of cellular events. In muscle contraction, \ceCa2+\ce{Ca^2+} binds specifically to , a regulatory subunit of the complex on thin filaments, inducing a conformational change that shifts away from myosin-binding sites on . This exposure allows myosin heads to interact with , forming cross-bridges powered by , which generates the sliding filament mechanism essential for force production in skeletal, cardiac, and smooth muscles. The process is tightly regulated, with \ceCa2+\ce{Ca^2+} dissociation leading to relaxation as re-blocks the sites. Calcium ions are required for the activity of numerous enzymes across various metabolic pathways, either as direct cofactors or through calcium-binding regulatory proteins such as , facilitating their activation. In blood clotting, for instance, \ceCa2+\ce{Ca^2+} is required for the activation of coagulation factors VII and IX, enabling the formation of the tenase complex that amplifies the proteolytic cascade leading to clot formation. This cofactor role extends to other processes, such as and , where \ceCa2+\ce{Ca^2+} modulates by altering substrate affinity or inducing allosteric changes. Finally, extracellular \ceCa2+\ce{Ca^2+} contributes to cell stability by modulating permeability and maintaining the integrity of bilayers. It interacts with proteins and to regulate gating and prevent excessive leakage, thereby preserving the electrochemical gradients vital for cellular excitability and volume control. Intracellularly, controlled \ceCa2+\ce{Ca^2+} levels further support function by influencing cytoskeletal attachments and vesicular trafficking.

Homeostasis and regulation

Calcium homeostasis in humans maintains serum calcium concentrations within a narrow range of 2.2–2.6 mmol/L (8.8–10.4 mg/dL) to support essential physiological processes. In this total serum calcium, approximately 50% circulates as ionized (free) calcium, which is biologically active; 40% is bound to proteins, primarily ; and 10% forms complexes with anions such as , citrate, and . This distribution ensures that only the ionized fraction directly influences cellular functions and signaling, while bound and complexed forms serve as reservoirs. The primary hormonal regulators of calcium are (PTH), calcitonin, and the active form of , 1,25-dihydroxyvitamin D₃ (). PTH, secreted by the parathyroid glands in response to low serum calcium, elevates calcium levels by stimulating osteoclast-mediated bone resorption to release calcium from crystals, enhancing renal reabsorption of calcium in the , and promoting the renal synthesis of . Calcitonin, produced by C-cells, opposes PTH by inhibiting activity and promoting renal calcium excretion, thereby lowering serum calcium, though its role in humans is less prominent under normal conditions. , synthesized in the kidneys from 25-hydroxyvitamin D under PTH stimulation, primarily increases intestinal calcium absorption efficiency from about 10–15% in deficiency states to 30–40% in sufficiency by upregulating the expression of calcium transport proteins in enterocytes. Cellular transport mechanisms are crucial for maintaining calcium balance across tissues. In the intestine, apical entry of calcium into enterocytes occurs primarily through TRPV6 channels, which are highly selective for calcium and upregulated by to facilitate transcellular absorption, particularly in the and . Once inside the cell, calcium is buffered by and extruded basolaterally via plasma membrane Ca²⁺-ATPase (PMCA1b) pumps or the Na⁺/Ca²⁺ exchanger (NCX1), ensuring vectorial transport into the bloodstream. Similar PMCA and NCX mechanisms operate in other tissues, such as the and cells, to fine-tune calcium flux and prevent intracellular overload. Bone serves as the main reservoir for calcium, with homeostasis achieved through continuous remodeling involving osteoclasts and osteoblasts. Osteoclasts dissolve in bone matrix via acidification and proteolytic enzymes, releasing calcium into the circulation to counteract ; this process is stimulated by PTH and (receptor activator of nuclear factor kappa-B ligand) produced by osteoblasts. Osteoblasts, in turn, deposit new to rebuild bone, maintaining structural integrity; they regulate osteoclastogenesis by balancing secretion with (OPG), a receptor that inhibits RANKL-RANK interactions on osteoclast precursors. This RANKL/OPG ratio ensures that and formation remain coupled, preventing net calcium loss under steady-state conditions. Feedback mechanisms tightly control these processes to stabilize serum calcium. Detection of low ionized calcium by calcium-sensing receptors on parathyroid chief cells triggers rapid PTH release, which acts within minutes to hours to restore levels through the aforementioned pathways. Conversely, rising calcium inhibits PTH secretion and stimulates calcitonin release, providing a counter-regulatory loop, while further amplifies intestinal uptake as a longer-term adjustment. These integrated systems—hormonal, transport, and remodeling—collectively maintain calcium despite varying dietary intake or physiological demands.

Pathological conditions

Hypocalcemia is defined as a serum calcium concentration below 2.2 mmol/L (8.8 mg/dL), resulting from impaired calcium mobilization from , reduced intestinal absorption, or increased renal . This condition can lead to neuromuscular irritability, manifesting as , muscle cramps, , and in severe cases, seizures or . In children, chronic hypocalcemia often contributes to , characterized by softened bones and skeletal deformities due to impaired mineralization, while in adults, it may cause , leading to and increased fracture risk. Hypercalcemia occurs when serum calcium exceeds 2.6 mmol/L (10.4 mg/dL), most commonly due to or malignancy-associated factors such as bone metastases or humoral hypercalcemia of malignancy. Symptoms arise from , renal effects, and involvement, including , , , kidney stones, and ; severe cases can progress to cardiac arrhythmias or . Osteoporosis is a systemic skeletal disorder defined by low bone mass and microarchitectural deterioration, diagnosed when bone mineral density (BMD) is more than 2.5 standard deviations below the young adult mean (T-score ≤ -2.5), often linked to chronic calcium dysregulation and deficiency in postmenopausal women, though it also affects older men, particularly after age 70. This condition affects over 200 million people worldwide, with postmenopausal women at highest risk due to accelerated outpacing formation, increasing susceptibility to fragility fractures at the , spine, and . Primary hyperparathyroidism, with a prevalence of approximately 1 in 1000 adults, involves autonomous overproduction of (PTH) by one or more parathyroid glands, typically due to , leading to excessive , hypercalcemia, and increased fracture risk. Elevated PTH promotes activity, releasing calcium from bone into the bloodstream, which can exacerbate renal stone formation and cardiovascular complications if untreated. Calcium channelopathies, genetic disorders affecting voltage-gated calcium channels, represent an emerging area of research in the 2020s, with conditions like Timothy syndrome—caused by mutations in the CACNA1C gene—resulting in gain-of-function alterations that disrupt calcium influx, contributing to neurodevelopmental issues such as autism spectrum disorder, seizures, and cardiac arrhythmias. Recent studies highlight the multisystem impact of these channel defects, emphasizing targeted therapies to modulate channel activity.

Safety

Handling elemental calcium

Elemental calcium is a highly reactive that poses significant hazards during handling due to its flammability and sensitivity to environmental factors. Finely divided calcium can spontaneously ignite in moist air, forming a dark oxide-nitride layer and potentially leading to fire if exposed to humidity levels above trace amounts. It reacts violently with water, liberating gas according to the reaction Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g), which can result in explosions if the gas accumulates in confined spaces or ignites. To prevent such reactions, calcium must be stored and manipulated under an inert atmosphere, such as , or submerged in to exclude oxygen and moisture. Proper storage is essential for maintaining the integrity of elemental calcium, which should be kept in tightly sealed containers made of compatible materials like or certain metals, in a cool, dry, and well-ventilated area isolated from acids, oxidizers, , and any sources of ignition. Grounding and bonding of containers help prevent static discharge that could spark a reaction. When stored under these conditions, elemental calcium maintains its usability for up to 36 months, though regular for surface tarnishing or degradation is recommended. Individuals handling elemental calcium require appropriate (PPE) to minimize risks of burns, irritation, or inhalation hazards. This includes or gloves, chemical-resistant coveralls, safety goggles or a full , and respiratory protection such as an or P3 filter respirator in dusty environments; is advised for high-exposure scenarios exceeding 30 mg/m³. Fire safety measures are critical, with Class D extinguishers containing dry chemicals, , or lime recommended for metal fires—water, , or halogenated agents must be avoided as they intensify the reaction by promoting evolution or . Industrial incidents involving calcium metal are uncommon owing to stringent handling protocols, but when they occur, they often stem from inadvertent exposure leading to rapid or formation, which causes severe burns to or eyes. Regulatory guidelines from the (OSHA) classify calcium dust under particulates not otherwise regulated, mandating an exposure limit of 5 mg/m³ as an 8-hour time-weighted average for the respirable fraction to protect against risks.

Health effects of imbalances

Ingestion of elemental calcium, a highly reactive metal, can lead to the formation of (Ca(OH)₂) upon contact with moisture, resulting in gastrointestinal irritation, , , and due to its alkaline and corrosive nature. Excessive intake or exposure to specifically contributes to , characterized by elevated blood pH, often alongside hypercalcemia and renal impairment in severe cases like . Inhalation of (CaO) or dust, common in industrial settings, irritates the and, with prolonged exposure, can cause , a fibrotic disease marked by scarring and reduced function. Overdose from exceeding the tolerable upper intake level of 2,500 mg/day for adults aged 19–50 (or 2,000 mg/day for those over 50) increases the risk of hypercalcemia, where serum calcium levels rise above 10.5 mg/dL, potentially leading to symptoms like , , and cardiac . Some 2024–2025 studies and reviews have suggested an association between and increased risks of cardiovascular events or in older adults, though evidence remains mixed and dietary calcium is generally safer. Common side effects include and , while chronic excess promotes renal calculi formation by elevating urinary calcium excretion (). In environmental contexts, drinking water hardness from calcium concentrations above 100 mg/L has been associated with reduced (CVD) risk in multiple studies and meta-analyses, potentially due to magnesium co-occurrence or anti-atherogenic effects, though remains unestablished and requires further randomized trials. Post-2010 research on CaO is limited, with most studies focusing on its or stress-alleviating properties rather than human health risks like or from or dermal exposure. As of 2025, updates from health authorities confirm the safety of calcium-fortified foods, with absorption rates around 30% and no widespread adverse effects when consumed within recommended dietary guidelines. Chronic exposure to excess (CaF₂), a fluoride source in some minerals or industrial dusts, is rare but can contribute to through elevated levels during tooth development, manifesting as or discoloration. This condition arises from disrupting enamel mineralization, though calcium itself may mitigate severity by binding ions.

References

  1. https://periodic-table.rsc.org/element/20/calcium
  2. https://www.webelements.com/calcium/geology.html
  3. https://www.ncbi.nlm.nih.gov/books/NBK482128/
  4. https://webbook.nist.gov/cgi/cbook.cgi?ID=C7440702&Units=SI&Mask=FFFF
  5. https://www.webelements.com/calcium/chemistry.html
  6. https://www.webelements.com/calcium/history.html
  7. https://www.webelements.com/calcium/isotopes.html
  8. https://pubchem.ncbi.nlm.nih.gov/element/Calcium
  9. https://www.thoughtco.com/calcium-element-facts-p2-606512
  10. https://www.webelements.com/calcium/atoms.html
  11. https://pubchem.ncbi.nlm.nih.gov/compound/Calcium
  12. https://www.webelements.com/calcium/crystal_structure.html
  13. https://periodictable.com/Elements/020/data.html
  14. https://chem.libretexts.org/Bookshelves/Inorganic_Chemistry/Supplemental_Modules_and_Websites_%28Inorganic_Chemistry%29/Descriptive_Chemistry/Elements_Organized_by_Block/1_s-Block_Elements/Group__2_Elements%253A_The_Alkaline_Earth_Metals/Z020_Chemistry_of_Calcium_%28Z20%29
  15. https://www.lenntech.com/periodic/elements/ca.htm
  16. https://chem.libretexts.org/Ancillary_Materials/Reference/Reference_Tables/Electrochemistry_Tables/P3%253A_Activity_Series_of_Metals
  17. https://owl.oit.umass.edu/departments/Chemistry/appendix/ksp.html
  18. https://www.sciencedirect.com/science/article/abs/pii/S2405829722005207
  19. https://www-nds.iaea.org/relnsd/vcharthtml/VChartHTML.html
  20. https://www.sciencedirect.com/science/article/pii/S0016703712001883
  21. https://www.technologynetworks.com/analysis/news/breaking-down-limits-with-41ca-dating-371962
  22. https://www.pnas.org/doi/10.1073/pnas.1307759110
  23. https://www.sciencedirect.com/science/article/pii/S0009254121003399
  24. https://www.graymont.com/en/markets/construction/building/mortar-applications/history-lime-mortar
  25. https://www.researchgate.net/publication/227658075_Microstructural_Evidence_of_Reconstituted_Limestone_Blocks_in_the_Great_Pyramids_of_Egypt
  26. https://www.sciencelearn.org.nz/resources/1420-history-of-limestone-uses-timeline
  27. https://www.encyclopedia.com/people/science-and-technology/chemistry-biographies/jons-jakob-baron-berzelius
  28. https://www.sciencehistory.org/stories/magazine/science-and-celebrity-humphry-davys-rising-star/
  29. https://www.chemicool.com/elements/calcium.html
  30. https://www.alchemywebsite.com/al_term1.html
  31. https://whc.unesco.org/en/list/1405
  32. https://news.mit.edu/2023/roman-concrete-durability-lime-casts-0106
  33. https://web.lemoyne.edu/giunta/ea/LAVPREFann.HTML
  34. https://www.nndc.bnl.gov/content/elements.html
  35. https://www.geokniga.org/bookfiles/geokniga-manueloriginelements.pdf
  36. http://hyperphysics.phy-astr.gsu.edu/hbase/Tables/elabund.html
  37. https://www.lenntech.com/periodic/water/calcium/calcium-and-water.htm
  38. https://www.sciencedirect.com/topics/physics-and-astronomy/calcite
  39. https://www.sciencelearn.org.nz/resources/468-calcium-carbonate-biomineralisation
  40. https://www.nasa.gov/wp-content/uploads/static/history/alsj/a17/A17SampleCat_1.pdf
  41. https://www.sciencedirect.com/topics/engineering/calcium-metal
  42. https://stacks.cdc.gov/view/cdc/10456/cdc_10456_DS1.pdf
  43. https://www.mdpi.com/2075-4701/8/6/383
  44. https://patents.google.com/patent/WO2013084672A1/en
  45. https://www.sciencedirect.com/science/article/abs/pii/S0025326X25009336
  46. https://www.pnas.org/doi/10.1073/pnas.1918943117
  47. https://www.soest.hawaii.edu/oceanography/faculty/zeebe_files/Publications/RidgwellZeebeEPSL05.pdf
  48. https://agupubs.onlinelibrary.wiley.com/doi/full/10.1029/2009GB003574
  49. https://pubchem.ncbi.nlm.nih.gov/compound/Calcium-Oxide
  50. https://pmc.ncbi.nlm.nih.gov/articles/PMC11313493/
  51. https://pubchem.ncbi.nlm.nih.gov/compound/Calcium-Carbonate
  52. https://etda.libraries.psu.edu/files/final_submissions/4998
  53. https://pubchem.ncbi.nlm.nih.gov/compound/Calcium-Chloride
  54. https://downloads.regulations.gov/EPA-HQ-OW-2009-0819-5450/attachment_10.pdf
  55. https://nvlpubs.nist.gov/nistpubs/jres/070/4/V70.N04.A03.pdf
  56. https://www.ams.usda.gov/sites/default/files/media/CalciumChlorideTR2021_0.pdf
  57. https://pubchem.ncbi.nlm.nih.gov/compound/Calcium-Fluoride
  58. https://cameochemicals.noaa.gov/chemical/8385
  59. https://ntrs.nasa.gov/api/citations/20070035754/downloads/20070035754.pdf
  60. https://pubchem.ncbi.nlm.nih.gov/compound/Calcium-Sulfate
  61. https://commonminerals.esci.umn.edu/minerals-g-m/gypsum
  62. https://peanut.ces.ncsu.edu/wp-content/uploads/2023/06/NSS-16-Gypsum.pdf
  63. https://www.uaex.uada.edu/publications/pdf/FSA2199.pdf
  64. https://gsr.lib.msu.edu/article/walworth-gypsum-5-25-12.pdf
  65. https://pubchem.ncbi.nlm.nih.gov/compound/Calcium-nitrate
  66. https://cameochemicals.noaa.gov/chemical/2785
  67. https://webbook.nist.gov/chemistry/silmarils-solids-hrf-drf/docs/META.MIR.PUR.Calcium-nitrate-tetrahydrate.SLD.SIGMA-ALDRICH.PNNL514811.pdf
  68. https://www.nj.gov/health/eoh/rtkweb/documents/fs/0324.pdf
  69. https://www.pnnl.gov/publications/thermodynamic-properties-aqueous-calcium-nitrate-cano32-373-k-including-new-enthalpy
  70. https://www.researchgate.net/publication/375064835_Modification_of_Non-metallic_Inclusions_in_Steel_by_Calcium_Treatment_A_Review
  71. https://www.mdpi.com/2075-4701/12/5/707
  72. https://www.sciencedirect.com/topics/engineering/calcium-treatment
  73. https://www.sciencedirect.com/science/article/abs/pii/037877539390029Z
  74. https://patents.google.com/patent/US6267923B1/en
  75. https://www.sciencedirect.com/science/article/abs/pii/0378775394019752
  76. https://www.belmontmetals.com/grain-refinement-aluminium-alloys/
  77. https://www.frontiersin.org/journals/materials/articles/10.3389/fmats.2022.1100566/full
  78. https://brite.org/cmi-hub-project-takes-key-step-to-process-rare-earth-metals-in-us/
  79. https://www.osti.gov/servlets/purl/4213369
  80. https://www.researchgate.net/publication/289439848_Calciothermic_reduction_of_neodymium_fluoride
  81. https://www.permutrade.com/what-is-cement-clinker-composition-types-uses
  82. https://drawdown.org/explorer/improve-cement-production
  83. https://actat.wvu.edu/files/d/e8f4ec5e-eda6-4889-983f-dc30342ed958/lime-softening.pdf
  84. https://www.watertechnologies.com/handbook/chapter-07-precipitation-softening
  85. https://www.alzchem.com/en/brands/perlka/effect-calcium-cyanamide/
  86. https://www.researchgate.net/publication/284135320_Calcium_cyanamide_-_A_synoptic_review_of_an_environmentally_benign_fertiliser_which_enhances_soil_health
  87. https://www.lime.org/resource/other-uses-of-lime/
  88. https://www.omri.org/sugar-lime-and-calcium-oxide
  89. https://pmc.ncbi.nlm.nih.gov/articles/PMC7245492/
  90. https://ods.od.nih.gov/factsheets/Calcium-HealthProfessional/
  91. https://ods.od.nih.gov/factsheets/Pregnancy-HealthProfessional/
  92. https://www.osteoporosis.foundation/educational-hub/topic/calcium/calcium-recommendations
  93. https://www.niams.nih.gov/health-topics/calcium-and-vitamin-d-important-bone-health
  94. https://medlineplus.gov/ency/patientinstructions/000490.htm
  95. https://pmc.ncbi.nlm.nih.gov/articles/PMC10721582/
  96. https://www.sciencedirect.com/science/article/pii/S0022316623727922
  97. https://pmc.ncbi.nlm.nih.gov/articles/PMC9306636/
  98. https://www.mayoclinic.org/healthy-lifestyle/nutrition-and-healthy-eating/in-depth/calcium-supplements/art-20047097
  99. https://list.essentialmeds.org/recommendations/333
  100. https://www.nature.com/articles/s41387-024-00321-2
  101. https://e-nrp.org/DOIx.php?id=10.4162/nrp.2025.19.2.155
  102. https://www.ncbi.nlm.nih.gov/books/NBK56060/
  103. https://www.ncbi.nlm.nih.gov/books/NBK109827/
  104. https://lpi.oregonstate.edu/mic/minerals/calcium
  105. https://www.ncbi.nlm.nih.gov/books/NBK27950/
  106. https://www.frontiersin.org/journals/pharmacology/articles/10.3389/fphar.2012.00061/full
  107. https://pmc.ncbi.nlm.nih.gov/articles/PMC1334730/
  108. https://www.nature.com/articles/s41467-019-14008-1
  109. https://www.sciencedirect.com/science/article/pii/S0006349503700411
  110. https://www.ncbi.nlm.nih.gov/books/NBK507850/
  111. https://pubmed.ncbi.nlm.nih.gov/6466679/
  112. https://www.sciencedirect.com/science/article/abs/pii/S1567539421002516
  113. https://academic.oup.com/bmb/article-abstract/42/4/391/292086
  114. https://www.uclahealth.org/medical-services/surgery/endocrine-surgery/patient-resources/patient-education/normal-calcium-levels
  115. https://emedicine.medscape.com/article/2087447-overview
  116. https://www.ncbi.nlm.nih.gov/books/NBK499940/
  117. https://teachmephysiology.com/biochemistry/electrolytes/calcium-regulation/
  118. https://pmc.ncbi.nlm.nih.gov/articles/PMC2669834/
  119. https://www.ncbi.nlm.nih.gov/books/NBK531440/
  120. https://pmc.ncbi.nlm.nih.gov/articles/PMC7327788/
  121. https://journals.physiology.org/doi/full/10.1152/physrev.00003.2004
  122. https://www.ncbi.nlm.nih.gov/books/NBK499863/
  123. https://pmc.ncbi.nlm.nih.gov/articles/PMC2413418/
  124. https://arthritis-research.biomedcentral.com/articles/10.1186/ar2165
  125. https://emedicine.medscape.com/article/874690-overview
  126. https://open.oregonstate.education/anatomy2e/chapter/calcium-homeostasis/
  127. https://www.ncbi.nlm.nih.gov/books/NBK279022/
  128. https://www.ncbi.nlm.nih.gov/books/NBK430912/
  129. https://www.ncbi.nlm.nih.gov/books/NBK430714/
  130. https://www.mayoclinic.org/diseases-conditions/hypercalcemia/symptoms-causes/syc-20355523
  131. https://www.ncbi.nlm.nih.gov/books/NBK441901/
  132. https://www.niams.nih.gov/health-topics/osteoporosis
  133. https://pmc.ncbi.nlm.nih.gov/articles/PMC9165055/
  134. https://www.ncbi.nlm.nih.gov/books/NBK441895/
  135. https://pmc.ncbi.nlm.nih.gov/articles/PMC80211/
  136. https://pmc.ncbi.nlm.nih.gov/articles/PMC8165229/
  137. https://pmc.ncbi.nlm.nih.gov/articles/PMC10576998/
  138. https://www.sigmaaldrich.com/US/en/sds/aldrich/441872
  139. https://pmc.ncbi.nlm.nih.gov/articles/PMC3875933/
  140. https://www.ncbi.nlm.nih.gov/books/NBK499888/
  141. https://pmc.ncbi.nlm.nih.gov/articles/PMC6438615/
  142. https://www.health.harvard.edu/staying-healthy/is-taking-too-much-calcium-unhealthy
  143. https://pmc.ncbi.nlm.nih.gov/articles/PMC3257663/
  144. https://pmc.ncbi.nlm.nih.gov/articles/PMC11879893/
  145. https://ehjournal.biomedcentral.com/articles/10.1186/s12940-025-01166-7
  146. https://www.frontiersin.org/journals/plant-science/articles/10.3389/fpls.2022.843795/full
  147. https://wwwn.cdc.gov/TSP/ToxFAQs/ToxFAQsDetails.aspx?faqid=211&toxid=38
  148. https://pmc.ncbi.nlm.nih.gov/articles/PMC3433161/
Add your contribution
Related Hubs
User Avatar
No comments yet.