Calcium Minerals

Calcium is a chemical element with the 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 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.

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.

Characteristics

Classification

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 twenty electrons, arranged in the electron configuration rs². 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.Greenwood and Earnshaw, pp. 112–13

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 behaviour: 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; above 450 °C, it changes to an anisotropic hexagonal close-packed arrangement like magnesium. Its density of 1.55 g/cm³ 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.Hluchan and Pomerantz, p. 484

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 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.C. R. Hammond ''The elements'' (pp. 4–35) in 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.Hluchan and Pomerantz, p. 483

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.Greenwood and Earnshaw, pp. 122–15 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 (Ca²⁺), 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).Greenwood and Earnshaw, pp. 136–37

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 stability.

Isotopes

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 leads 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 capture 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 survival.

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 ⁴⁸ Sc 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 ⁴⁸ Ti 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 ⁴⁰ Ar, 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 respectively.

Apart from the practically stable ⁴⁸Ca, the longest lived radioisotope of calcium is ⁴¹Ca. It decays by electron capture to stable ⁴¹ K 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 ³⁵Ca to ⁶⁰Ca. 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 materials.

History

Calcium compounds were known for millennia, although their chemical makeup was not understood until the 17th century.Greenwood and Earnshaw, p. 108 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 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.

In 1789, Antoine Lavoisier

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