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 mos
Chemical element with atomic number 20 (Ca) Mg
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Ca
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Sr potassium ← calcium → scandium Atomic number (Z)20Groupgroup 2 (alkaline earth metals)Periodperiod 4Block s-blockElectron configuration[Ar] 4s2Electrons per shell2, 8, 8, 2Physical propertiesPhase at STPsolidMelting 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) Specific heat capacity646.963 J/(kg·K) Vapour pressure Atomic propertiesOxidation 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
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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 the Latin word calx (meaning "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
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 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
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
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 Main article: Organocalcium chemistryIn 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
Main article: Isotopes of calciumNatural 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]
