======================================================================
= Calcium =
======================================================================
Introduction
======================================================================
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. 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. Calcium ions outside cells are important for
maintaining the potential difference across excitable cell membranes,
protein synthesis, and bone formation.
Classification
================
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]4s(2). 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.
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
MX is much higher than those of the hypothetical MX. This occurs
because of the much greater lattice energy afforded by the more highly
charged Ca(2+) cation compared to the hypothetical Ca(+) cation.
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.
Physical properties
=====================
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, it changes to body-centered cubic.
Its density of 1.526 g/cm(3) (at 20 °C) is the lowest in its group.
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. 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.
Chemical properties
=====================
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. 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.
Besides the simple oxide CaO, calcium peroxide, CaO, 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(O).Calcium
hydroxide, Ca(OH), is a strong base, though not as strong as the
hydroxides of strontium, barium or the alkali metals. All four
dihalides of calcium are known. Calcium carbonate (CaCO) and calcium
sulfate (CaSO) 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 calcium ion (Ca(2+)), high coordination
numbers are common, up to 24 in some intermetallic compounds such as
CaZn. 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.
Organocalcium compounds
=========================
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.
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,
whereas dimethylmagnesium is a linear polymer with tetrahedral Mg
centers. Bulky ligands are often required to disfavor polymeric
species. For example, calcium dicyclopentadienyl, has a polymeric
structure and thus is nonvolatile and insoluble in solvents. Replacing
the ligand with the bulkier (pentamethylcyclopentadienyl) gives a
soluble complex that sublimes and forms well-defined adducts with
ethers. Organocalcium compounds tend to be more similar to
organoytterbium compounds due to the similar ionic radii of Yb(2+)
(102 pm) and Ca(2+) (100 pm).
Organocalcium compounds have been well investigated. Some such
complexes exhibit catalytic properties, although none have been
commercialized.
Isotopes
==========
Natural calcium is a mixture of five stable isotopes--(40)Ca, (42)Ca,
(43)Ca, (44)Ca, and (46)Ca--and (48)Ca, whose half-life of 4.3 ×
10(19) years is so long that it can be considered stable all practical
purposes stable. Calcium is the first (lightest) element to have six
naturally occurring isotopes.
By far the most common isotope is (40)Ca, 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 (40)K. Adding another alpha particle
leads to unstable (44)Ti, which decays via two successive electron
captures to stable (44)Ca; this makes up 2.806% of natural calcium and
is the second-most common isotope.
The other four natural isotopes, (42, 43, 46, 48)Ca, 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. (46)Ca is mostly produced in a "hot" s-process, as
its formation requires a rather high neutron flux to allow short-lived
(45)Ca to capture a neutron. (48)Ca is produced by electron capture in
the r-process in type Ia supernovae, where high neutron excess and low
enough entropy ensures its survival.
(46)Ca and (48)Ca are the first "classically stable" nuclides with a
6-neutron or 8-neutron excess respectively. Though extremely
neutron-rich for such a light element, (48)Ca is very stable because
it is a doubly magic nucleus, with 20 protons and 28 neutrons arranged
in closed shells. Its beta decay to (48)Sc is very hindered by the
gross mismatch of nuclear spin: (48)Ca has zero nuclear spin, being
even-even, while (48)Sc has spin 6+, so the decay is forbidden by
conservation of angular momentum. While two excited states of (48)Sc
are available for decay as well, they are also forbidden due to their
high spins. As a result, when (48)Ca does decay, it does so by double
beta decay to (48)Ti instead, being the lightest nuclide known to
undergo double beta decay.
(46)Ca can also theoretically double-beta-decay to (46)Ti, but this
has never been observed. The most common isotope (40)Ca is also doubly
magic and could undergo double electron capture to (40)Ar, 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 (40)Ca and (46)Ca are 5.9 × 10(21) years and 2.8 ×
10(15) years respectively.
Excluding (48)Ca, the longest lived radioisotope of calcium is (41)Ca.
It decays by electron capture to stable (41)K with a half-life of
about 10(5) years. Its existence in the early Solar System as an
extinct radionuclide has been inferred from excesses of (41)K. Traces
of (41)Ca also still exist today, as it is a cosmogenic nuclide,
continuously produced through neutron activation of natural (40)Ca.
Many other calcium radioisotopes are known, ranging from (35)Ca to
(60)Ca. They are all much shorter-lived than (41)Ca; the most stable
are (45)Ca (half-life 163 days) and (47)Ca (half-life 4.54 days).
Isotopes lighter than (42)Ca usually undergo beta plus decay to
isotopes of potassium, and those heavier than (44)Ca 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.
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 (44)Ca/(40)Ca) in a sample compared to
the same ratio in a standard reference material. (44)Ca/(40)Ca varies
by about 1-2‰ among organisms on Earth.
History
======================================================================
Calcium compounds were known for millennia, though 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.
About the same time, dehydrated gypsum (CaSO·2HO) 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 (CaCO). 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.
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:
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. 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
======================================================================
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 (CaSO·2HO),
anhydrite (CaSO), fluorite (CaF), and apatite ([Ca(PO)X], X = OH, Cl,
or F)
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.
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 U.S. and Canada, calcium is instead produced by reducing lime
with aluminium at high temperatures. 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. 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. High-purity calcium can be obtained by distilling
low-purity calcium at high temperatures.
Geochemical cycling
=====================
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 Ca(2+) into surface water. These ions are
transported to the ocean where they react with dissolved CO to form
limestone (CaCO), which in turn settles to the sea floor where it is
incorporated into new rocks. Dissolved CO, along with carbonate and
bicarbonate ions, are termed "dissolved inorganic carbon" (DIC).
The actual reaction is more complicated and involves the bicarbonate
ion (HCO) that forms when CO reacts with water at seawater pH:
:
At seawater pH, most of the dissolved CO is immediately converted back
into . The reaction results in a net transport of one molecule of CO
from the ocean/atmosphere into the lithosphere. The result is that
each Ca(2+) ion released by chemical weathering ultimately removes one
CO 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 CO from the ocean and air, exerting a strong
long-term effect on climate.
Applications
======================================================================
The largest use of metallic 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 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.
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
lead. 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, 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.
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
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.
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
(44)Ca/(40)Ca 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 (44)Ca/(40)Ca tends to rise
when the rate of removal of Ca(2+) 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 (44)Ca/(40)Ca over
geologic time, along with a theoretical explanation of these changes.
More recent papers have confirmed this observation, demonstrating that
seawater Ca(2+) 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 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.
Calcium is on the World Health Organization's List of Essential
Medicines.
Food sources
======================================================================
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.
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. 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.
Biological and pathological role
======================================================================
Age-adjusted daily calcium recommendations (from U.S. Institute of
Medicine RDAs)
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
Function
==========
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 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.
Sosa Torres, Martha; Kroneck, Peter M.H; "Introduction: From Rocks to
Living Cells" pp. 1-32 in "Metals, Microbes and Minerals: The
Biogeochemical Side of Life" (2021) pp. xiv + 341. Walter de Gruyter,
Berlin. Editors Kroneck, Peter M.H. and Sosa Torres, Martha.
Binding
=========
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-binding enzymes.
Calcium also binds to the phospholipid layer of the cell membrane,
anchoring proteins associated with the cell surface.
Solubility
============
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.
Nutrition
===========
Calcium is a common constituent of multivitamin dietary supplements,
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. The
intestine absorbs about one-third of calcium eaten as the free ion,
and plasma calcium level is then regulated by the kidneys.
Hormonal regulation of bone formation and serum levels
========================================================
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.
Abnormal serum levels
=======================
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.
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.
Bone disease
==============
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. Inadequate
amounts of calcium, vitamin D, or phosphates can lead to softening of
bones, called osteomalacia.
Metallic calcium
==================
{{Chembox
|container_only = yes
|Section7=
}}
Because calcium reacts exothermically with water and acids, calcium
metal coming into contact 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.
License
=========
All content on Gopherpedia comes from Wikipedia, and is licensed under CC-BY-SA
License URL:
http://creativecommons.org/licenses/by-sa/3.0/
Original Article:
http://en.wikipedia.org/wiki/Calcium
