Calcium is a chemical element with symbol Ca and atomic number 20. An
alkaline earth metal, calcium is a reactive pale yellow 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 fossilised remnants of early sea life; gypsum,
anhydrite, fluorite, and apatite are also sources of calcium.
The name derives from Latin calx "lime", which was obtained from
heating limestone. Its compounds were known to the ancients, though
their chemistry was unknown until the seventeenth century. It was
Humphry Davy in 1808 via electrolysis of its oxide, who
named the element. While the pure metal does not have many
applications due to its high reactivity, it is often used as an
alloying component in small quantities in steelmaking, and
calcium–lead alloys are sometimes used in automotive batteries.
Calcium compounds on the other hand are very widely used in many
industries: for example, they are used in foods and pharmaceuticals
for calcium supplementation, in the paper industry as bleaches, in
cement, in the manufacture of soaps, and as electrical insulators.
Calcium is the fifth most abundant element in the human body and the
most abundant metal.
Calcium ions play a vital role in the physiology
and biochemistry of organisms and the cell as electrolytes. They play
an important role in signal transduction pathways, where they act as a
second messenger, in neurotransmitter release from neurons, in
contraction of all muscle cell types, and in fertilization. Many
enzymes require calcium ions as a cofactor.
Calcium ions outside cells
are also important for maintaining the potential difference across
excitable cell membranes, as well as proper bone formation.
3 Occurrence and production
3.1 Geochemical cycling
5 Biological and pathological role
7 See also
10 External links
Calcium is a very ductile silvery metal with a pale yellow tint whose
properties are very similar to the heavier elements in its group,
strontium, barium, and radium. A calcium atom has twenty electrons,
arranged in the electron configuration [Ar]4s2. Like the other
elements placed 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. 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.
Calcium is considered to be an alkaline earth metal, along with these
heavier elements and the lighter beryllium and magnesium.
Nevertheless, there are significant differences in chemical and
physical properties between beryllium and magnesium (which behave more
like aluminium and zinc respectively and have some of the weaker
metallic character of the post-transition metals) and the group
members from calcium onwards, which traditionally led to "alkaline
earth metal" only applying to the latter group. This classification
is mostly obsolete in English-language sources, but is still used in
other countries such as Japan. As a result, comparisons with
strontium and barium are more germane to calcium chemistry than
comparisons with magnesium.
Calcium metal melts at 842 °C and boils at 1494 °C, higher
than its adjacent group 2 metals do. It crystallises in the
face-centered cubic arrangement like strontium; above 450 °C, it
changes to an anisotropic hexagonal close-packed arrangement like
magnesium. The density of 1.55 g·cm−3 is the lowest in its
group, with others decreasing towards it.
Calcium can be cut with a
knife with effort, although it is still harder than lead. While
calcium is a poorer conductor of electricity than copper or aluminium
by volume, it is a better conductor than both of them by mass due to
its very low density. Although it is infeasible for terrestrial
applications as it reacts quickly with atmospheric oxygen, its use as
a conductor in space has been considered.
Structure of the polymeric [Ca(H2O)6]2+ center in hydrated calcium
chloride, illustrating the high coordination number typical for
The chemistry of calcium is that of a typical heavy alkaline earth
metal. For example, calcium spontaneously reacts with water more
quickly than magnesium and less quickly than strontium to produce
calcium hydroxide and hydrogen gas. It also reacts with the oxygen and
nitrogen in the air to form a mixture of calcium oxide and calcium
nitride. When finely divided, it spontaneously burns in air to
produce the nitride. In 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.
Besides the simple oxide CaO, the 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. Calcium
hydroxide, Ca(OH)2, is a strong base, though it is not as strong as
the hydroxides of strontium, barium or the alkali metals. All four
dihalides of calcium are known.
Calcium carbonate (CaCO3) and
calcium sulfate (CaSO4) are particularly abundant minerals. 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.
Due to the large size of the Ca2+ ion, high coordination numbers are
common, up to 24 in some intermetallic compounds such as CaZn13.
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.
Although calcium is in the same group as magnesium and organomagnesium
compounds are very commonly used throughout chemistry, organocalcium
compounds are not similarly widespread because they are more difficult
to make and more reactive, although they have recently been
investigated as possible catalysts. Organocalcium
compounds tend to be more similar to organoytterbium compounds due to
the similar ionic radii of Yb2+ (102 pm) and Ca2+ (100 pm).
Most of these compounds can only be prepared at low temperatures;
bulky ligands tend to favor stability. For example, calcium
dicyclopentadienyl, Ca(C5H5)2, must be made by directly reacting
calcium metal with mercurocene or cyclopentadiene itself; replacing
the C5H5 ligand with the bulkier C5(CH3)5 ligand on the other hand
increases the compound's solubility, volatility, and kinetic
Main article: Isotopes of calcium
Natural calcium is a mixture of five stable isotopes (40Ca, 42Ca,
43Ca, 44Ca, and 46Ca) and one isotope with a half-life so long that it
can be considered stable for all practical purposes (48Ca, with a
half-life of about 4.3 × 1019 years).
Calcium is the
first (lightest) element to have six naturally occurring isotopes.
By far the most common isotope of calcium in nature is 40Ca, which
makes up 96.941% of all 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 would lead to unstable 44Ti, which
quickly decays via two successive electron captures to stable 44Ca;
this makes up 2.806% of all natural calcium and is the second-most
common isotope. The other four natural isotopes, 42Ca, 43Ca, 46Ca, and
48Ca, are significantly rarer, each comprising less than 1% of all
natural calcium. The four lighter isotopes are mainly products of the
oxygen-burning and silicon-burning processes, leaving the two heavier
ones to be produced via neutron-capturing processes. 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
46Ca and 48Ca are the first "classically stable" nuclides with a
six-neutron or eight-neutron excess respectively. Although extremely
neutron-rich for such a light element, 48Ca is very stable because it
is a doubly magic nucleus, having 20 protons and 28 neutrons arranged
in closed shells. Its beta decay to 48Sc is very hindered because of
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 the
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. The heavy isotope 46Ca can also
theoretically undergo double beta decay to 46Ti as well, but this has
never been observed; the lightest and 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 to
have 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
Apart from the practically stable 48Ca, the longest lived radioisotope
of calcium is 41Ca. It decays by electron capture to stable 41K with a
half-life of about a hundred thousand 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 reformed through neutron activation
of natural 40Ca. Many other calcium radioisotopes are known,
ranging from 34Ca to 57Ca: they are all much shorter-lived than 41Ca,
the most stable among them being 45Ca (half-life 163 days) and
47Ca (half-life 4.54 days). The isotopes lighter than 42Ca
usually undergo beta plus decay to isotopes of potassium, and those
heavier than 44Ca usually undergo beta minus decay to isotopes of
scandium, although near the nuclear drip lines proton emission and
neutron emission begin to be significant decay modes as well.
Like other elements, a variety of processes alter the relative
abundance of calcium isotopes. 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% among common earth
One of the 'Ain Ghazal Statues, made from lime plaster
Calcium compounds were known for millennia, although their chemical
makeup was not understood until the 17th century. Lime as a
building material and as plaster for statues was used as far back
as around 7000 BC. The first dated lime kiln dates back to
2500 BC and was found in Khafajah, Mesopotamia. At 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 climate of present-day Italy
being warmer than that of Egypt, the ancient Romans instead used lime
mortars made by heating limestone (CaCO3); the name "calcium" itself
derives from the Latin word calx "lime".
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.
Antoine Lavoisier suspected that lime might be an oxide of a
fundamental chemical 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 speculated:
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
Calcium, along with its congeners magnesium, strontium, and barium,
was first isolated by
Humphry Davy in 1808. Following the work of
Jöns Jakob Berzelius
Jöns Jakob Berzelius and Magnus Martin af 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. 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.
Occurrence and production
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. It is also the fourth most abundant element in the lunar
highlands. 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
The major producers of calcium are
China (about 10000 to 12000 tonnes
Russia (about 6000 to 8000 tonnes per year), and the United
States (about 2000 to 4000 tonnes per year).
also 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. In
Russia and China, Davy's method of electrolysis is still used, but is
instead applied to molten calcium chloride. 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. In the United
States and Canada, calcium is instead produced by reducing lime with
aluminium metal at high temperatures.
Main article: Carbonate–silicate cycle
Calcium provides a link between tectonics, climate, and the carbon
cycle. In the simplest terms, uplift of mountains exposes
calcium-bearing rocks to chemical weathering and releases Ca2+ into
surface water. These ions are transported to the ocean where they
react with dissolved CO2 to form limestone, 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).
The actual reaction is more complicated and involves the bicarbonate
3) that forms when CO2 reacts with water at seawater pH:
Ca2+ + 2HCO−
3 → CaCO
3 (limestone) + CO
2 + H
At seawater pH, most of the CO2 is immediately converted back into
3. The reaction results in a net transport of one molecule of CO2 from
the ocean/atmosphere into the lithosphere. 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.
The largest use of calcium is in steelmaking, due to its strong
chemical affinity for 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 became 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. 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.
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
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. 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, 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
Calcium isotope fractionation during mineral formation has led to
several applications of calcium isotopes. In particular, the 1997
observation by Skulan and DePaolo that calcium minerals are
isotopically lighter than the solutions from which the minerals
precipitate is the basis of analogous applications in medicine and in
paleooceanography. In animals with skeletons mineralized with calcium,
the calcium isotopic composition of soft tissues reflects the relative
rate of formation and dissolution of skeletal mineral. In humans,
changes in the calcium isotopic composition of urine have been shown
to be related to changes in bone mineral balance. When the rate of
bone formation exceeds the rate of bone resorption, the 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. 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.
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 monophosphate 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.
Biological and pathological role
Calcium in biology
Age-adjusted daily calcium recommendations (from U.S. Institute of
Calcium is an essential element needed in large quantities. 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; 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
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. 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 also binds to the phospholipid layer
of the cell membrane, anchoring proteins associated with the cell
surface. 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
2.0 mM and the hydroxyapatite of bones in an organic matrix is
tricalcium phosphate at 100 µM.
About three-quarters of dietary calcium is from dairy products and
grains, the rest being accounted for by vegetables, protein-rich
foods, fruits, sugar, fats, and oil.
Calcium supplementation is
controversial, as the bioavailability of calcium is strongly dependent
on the solubility of the salt involved: calcium citrate, malate, and
lactate are highly bioavailable while the oxalate is much less so. The
intestine absorbs about one-third of calcium eaten as the free ion,
and plasma calcium level is then regulated by the kidneys. 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. 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.
Excess intake of calcium may cause hypercalcaemia, but because of the
inefficient absorption of calcium by the intestines a more likely
cause is excessive vitamin D intake or excessive secretion of
parathyroid hormone. It can also occur due to the bone destruction
that occurs when tumours metastasise to bone. This results in
deposition of calcium salts into the heart, the blood vessels, and the
kidneys. Symptoms include anorexia, nausea, vomiting, memory loss,
confusion, muscle weakness, increased urination, dehydration, and
metabolic bone disease. Chronic hypercalcaemia may lead to soft tissue
calcification, which can lead to serious consequences: for example,
calcification of the vascular wall can lead to a loss of elasticity
and the disruption of laminar blood flow, and thence to plaque rupture
and thrombosis. Likewise, inadequate calcium or vitamin D intake
results in hypocalcaemia, often caused by inadequate secretion of
parathyroid hormone or defective receptors to it in cells. Symptoms
include neuromuscular excitability, potentially causing tetany and
defects in cardiac conduction.
As calcium is heavily involved in bone manufacture, many bone diseases
can be traced to problems with the organic matrix or the
hydroxyapatite in molecular structure or organisation. For example,
osteoporosis is a reduction in mineral content of bone per unit
volume, and can be treated by supplementation of calcium, vitamin D,
Calcium supplements may benefit the serum lipids in
women who have passed menopause as well as older men; in
post-menopausal women calcium supplementation also appears to be
inversely correlated with cardiovascular disease. Inadequate amounts
of calcium, vitamin D, or phosphates can lead to the softening of
bones, known as osteomalacia.
Because calcium reacts exothermically with water and acids, contact of
calcium metal with bodily moisture results in severe corrosive
irritation. When swallowed, calcium metal has the same effect on
the mouth, oesophagus, and stomach, and can be fatal. However,
long-term exposure is not known to have distinct adverse effects.
Because of concerns of long-term adverse side effects such as
calcification of arteries and kidney stones, the U.S. Institute of
Medicine (IOM) and the
European Food Safety Authority
European Food Safety Authority (EFSA) both set
Tolerable Upper Intake Levels (ULs) for the combination of dietary and
supplemental calcium. From the IOM, people ages 9–18 years are not
supposed to exceed 3,000 mg/day; for ages 19–50 not to exceed
2,500 mg/day; for ages 51 and older, not to exceed
2,000 mg/day. The EFSA set UL at 2,500 mg/day for adults
but decided the information for children and adolescents was not
sufficient to determine ULs.
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BNF: cb11976285h (data)
Pharmacy and pharmacology port