Calcium is a chemical element with a symbol of Ca and an atomic number of 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 the heavier strontium and barium homologues. It is the fifth most abundant element in the earth's crust and the third most abundant metal, after iron and aluminum. The most common calcium compounds on Earth are calcium carbonate, found in limestone and fossil remains of early marine life; gypsum, anhydrite, fluorite, and apatite are also sources of calcium.
The name is derived from the Latin calx "lime", obtained from limestone heating. The compounds were known to the ancestors, although their chemistry was not known until the seventeenth century. It was isolated by Humphry Davy in 1808 through oxide electrolysis, which was named element. While pure metal does not have many applications because of its high reactivity, it is often used as a small alloy component in steelmaking, and calcium-lead alloys are sometimes used in automotive batteries. Calcium compounds on the other hand are widely used in many industries: for example, they are used in food and medicine for calcium supplementation, in the paper industry as bleach, in cement, in soap making, and as electrical insulators.
Calcium is the fifth most abundant element in the human body and the most abundant metal. Calcium ions play an important role in the physiology and biochemistry of organisms and cells as electrolytes. They play an important role in the signal transduction pathway, where they act as the second messenger, in the neurotransmitter release of neurons, the contraction of all types of muscle cells, and in conception. Many enzymes require calcium ions as cofactors. External calcium calcium ions are also important to maintain potential differences across cell membranes, as well as proper bone formation.
Video Calcium
Characteristics
Classification
Calcium is a very brittle silver metal with a pale yellow color that is very similar to the heavier elements in the group, strontium, barium, and radium. A calcium atom has twenty electrons, arranged in electron configuration [Ar] 4s 2 . Like other elements placed in group 2 of the periodic table, calcium has two valence electrons in the outer s-orbitals, which are very easily lost in chemical reactions to form positively charged ions with the stable electron configuration of the noble gases, in this case Argon. Therefore, calcium is almost always divalent in its compounds, which are usually ionic. The hypothetical hypoxic univalent salt will be stable with respect to its elements, but not to disproportionate to divalent salts and calcium metals, since the enthalpy of MX formation 2 is much higher than that of hypothetical MX. This is because the much larger lattice energy given by Ca 2 cation is higher than that of Ca hypothetical cation.
Calcium is regarded as an alkaline earth metal, along with heavier elements and beryllium and lighter magnesium. Nevertheless, there is a significant difference in the chemical and physical properties between beryllium and magnesium (which behave more like aluminum and zinc respectively and have some weaker metal characters than post-transition metals) and group members of the onward calcium, which traditionally led to "alkaline earth metals" applies only to the last group. This classification is largely 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 associated with calcium chemistry than comparisons with magnesium.
Physical
The metal of calcium melts at 842 Â ° C and boils at 1494 Â ° C, higher than the adjacent group 2 metals do. It crystallizes in a face-centered cubic arrangement such as strontium; above 450 ° C, it turns into a closed hexagonal anisotropic arrangement such as magnesium. The density of 1.55 g/cm 3 is the lowest in the group, with the other descending towards it. Calcium can be cut with a knife by effort, although it is still harder than lead. While calcium is an electrical conductor worse than copper or aluminum by volume, it is a better conductor than both because of its mass because of its very low density. Although not feasible for terrestrial applications because it reacts quickly with atmospheric oxygen, its use as an inner conductor has been considered.
Chemistry
The chemical properties of calcium are typical heavy alkaline earth metals. For example, calcium spontaneously reacts with water faster than magnesium and is less rapid than strontium to produce calcium hydroxide and hydrogen gas. It also reacts with oxygen and nitrogen in the air to form a mixture of calcium oxide and calcium nitride. When it is finely divided, it spontaneously burns in air to produce nitrides. In large quantities, calcium is less reactive: it quickly forms a hydration layer in moist air, but below 30% relative humidity can be stored indefinitely at room temperature.
In addition to simple CaO oxide, CaO peroxide 2 can be prepared by direct oxidation of the calcium metal under high oxygen pressure, and there is some evidence for yellow superoxide (O 2 ) sub sub > 2 . Calcium hydroxide, Ca (OH) 2 , is a strong base, though not as strong as hydroxides of strontium, barium or alkali metals. All four calcium dihalides are known. Calcium carbonate (CaCO 3 ) and calcium sulphate (CaSO 4 ) are very abundant minerals. Like strontium and barium, as well as alkaline and bivalve metal divalen europium and ytterbium, calcium metal directly dissolves in liquid ammonia to produce a dark blue solution.
Because of the large size of Ca 2 ions, high coordination rates are common, up to 24 in some intermetallic compounds such as CaZn 13 . Calcium is ready to be complexed by oxygen chelates such as EDTA and polyphosphates, which are useful in analytical chemistry and removing calcium ions from hard water. In the absence of steric hindrance, the 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 that are very commonly used throughout chemistry, organocalcium compounds are not as widespread as 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 because of the same ionic radii of Yb 2 (102Ã, pm) and Ca 2 (100Ã, pm). Most of these compounds can only be prepared at low temperatures; Large ligands tend to support stability. For example, diciclopentadienyl calcium, Ca (C 5 H 5 ) 2 , should be done by directly reacting the calcium metal with mercurocene or cyclopentadiene alone; replacing C 5 H 5 ligand with bulkier C 5 (CH 3 ) 5 ligands on the other hand increase solubility, volatility, and kinetic stability of the compound.
Isotope
Natural calcium is a mixture of five stable isotopes ( 40 Ca, 42 Ca, 43 Ca, 44 Ca, and 46 Ca) and one isotope with half-life so long that it can be considered stable for all practical purposes ( 48 Ca, with a half-life of about 4.3 Ã, ÃÆ' â € "10 < soup> 19 Ã, year). Calcium is the first (lightest) element to have six natural isotopes.
By far the most common calcium isotope in nature is 40 Ca, which makes up 96.941% of all natural calcium. It is produced in the silicon combustion process of alpha particle fusion and is the heaviest stable nuclide with the same number of protons and neutrons; the incidence is also added slowly by the decay of primordial 40 K. Adding other alpha particles will cause unstable <44 soup, which rapidly decays through two successive electron catches to a stable Ca 44 ; this means 2.806% of all natural calcium and is the second most common isotope. Four other natural isotopes, 42 Ca, 43 Ca, 46 Ca, and 48 Ca, significantly less , each containing less than 1% of all natural calcium. The four lighter isotopes, especially the products of the oxygen-burning and silicon-burning processes, leave two heavier isotopes to be produced through neutron capture. 46 Ca is mostly produced in "hot" s-processes, since its formation requires a rather high flux of neutrons to allow short life of Ca to capture neutrons. 48 Ca is generated by electron capture in the r-process in Ia-type supernovae, where high neutron excess and low entropy guarantee its survival.
46 Ca and 48 Ca is the first "classical stable" nuclear with six neutrons or eight excess neutrons. Although very rich in neutrons for such light elements, the 48 Ca is very stable because it is a double magical nucleus, has 20 protons and 28 neutrons arranged in a sealed shell. The decay of its beta to 48 Sc was severely hindered by the gross sponge incompatibility: 48 Ca has zero nuclear rotation, becomes even, while 48 Sc has round 6, so its decay is forbidden by conservation of angular momentum. While the two happy conditions 48 Sc are available for decay as well, they are also prohibited because of their high spins. Consequently, when the 48 Ca decays, it does so with a double beta decay to a 48 Ti instead, being the lightest nuclide known to have double beta decay. Heavy isotope 46 Ca also theoretically can experience double beta decay to 46 Ti too, but this is never observed; the lightest and most common isotope 40 Ca is also double magic and can have double electron capture to 40 Ar, but this is also never observed. Calcium is the only element that has two dual primordial magical isotopes. The lower limit of experiments for half-time 40 Ca and 46 Ca were 5,9Ã,ÃÆ'â € "10 21 Ã, year and 2.8Ã , ÃÆ' â € "Ã, 10 15 Ã, successive years.
Apart from the 48 Ca which is practically stable, the longest living radioisotope of calcium is 41 Ca. It decays by capturing the electrons to a stable 41 K with a half-life of about a hundred thousand years. Its presence in the early solar system as extinct radionuclides has been inferred from the excess 41 K: traces of 41 Ca also still present, as cosmogenic nuclides, are constantly reformed through neutron activation natural 40 Ca. Many other known calcium radioisotopes, ranging from 34 Ca to 57 Ca: they are all much shorter-alive than the 41 Ca, the most stable of which are 45 Ca (half-time 163 days) and 47 Ca (half-life of 4.54 days). Isotopes lighter than 42 Ca usually suffers from beta plus decay of potassium isotope, and the heavier than 44 Ca usually undergoes beta decay reduced to the scandium isotope, although near the nuclear drip line proton emissions and neutron emissions begin to become significant decay modes as well.
Like other elements, various processes change the relative abundance of calcium isotopes. The most learned of this process is fractionation depending on the mass of calcium isotopes that accompany the deposition of calcium minerals such as calcite, aragonite and apatite from the solution. Lighter isotopes are typically incorporated into these minerals, leaving behind a solution enriched with heavier isotopes with a magnitude of about 0.025% per atom mass unit (amu) at room temperature. Mass dependent differences in calcium isotope compositions are conventionally expressed by the ratio of two isotopes (usually 44 Ca/ 40 Ca) in the sample compared to the same ratio in the reference material standard. 44 Ca/ 40 Ca varies about 1% between ordinary earth materials.
Maps Calcium
History
Calcium compounds are known for thousands of years, although their chemical composition was not understood until the 17th century. Lime as a building material and as a plaster for sculpture used as far as about 7000 BC. The first lime luster dates back to 2500 BC and is found in Khafajah, Mesopotamia. At the same time, dehydration gypsum (CaSO 4 Ã, Â · 2H 2 O) is used in Great Pyramid of Giza; this material will be used for plaster in Tutankhamun's tomb. The current Italian climate is warmer than Egypt, the ancient Romans instead used lime lime made by heating limestone (CaCO 3 ); the name "calcium" itself comes from the Latin calx "lime". Vitruvius notes that the resulting lime is lighter than the original limestone, linking it with boiling water; in 1755, Joseph Black proved that this was due to the loss of carbon dioxide, which as gas has not been recognized by the ancient Romans.
In 1787, Antoine Lavoisier suspected that lime might be oxides of a fundamental chemical element. In the table of its elements, Lavoisier lists five "salable soils" (ie, ores that can be made to react with acids to produce saline (salt = salt, in Latin): chaux (calcium oxide), magnesium (magnesia, magnesium oxide), baryte (barium sulfate), alumine ( alumina, aluminum oxide), and silica (silica, silicon dioxide).About this "element", Lavoisier speculates:
We may be acquainted with only part of the metallic substances present in nature, since all that have a stronger affinity for oxygen than carbon can not, until now, be reduced to a metallic state, and consequently, only presented for our observations below forms of oxyds, confused with the earth. It is possible that barite-barite, which we just set with the earth, is in this situation; because in many experiments, it shows properties almost near metal objects. There is even the possibility that all the substances we call the earth may be just metal oxides, can not be reduced by the processes that are now known.
Calcium, together with magnesium, strontium, and barium, was first isolated by Humphry Davy in 1808. After working JÃÆ'¶ns Jakob Berzelius and Magnus Martin af Pontin on electrolysis, Davy isolated calcium and magnesium by placing a mixture of each metal oxide with mercury (II) oxide on a platinum plate used as an anode, the cathode being a platinum wire partially submerged into mercury. Electrolysis then gives the amalgam of calcium-mercury and magnesium-mercury, and refining mercury gives the metal. However, pure calcium can not be mass-prepared by this method and commercially viable processes for production were not discovered until more than a century later.
Genesis and production
At 3%, calcium is the fifth most element in the earth's crust, and the third most metals behind aluminum and iron. It is also the fourth most common element in the lunar plateau. Sedimentary calcium carbonate deposits include the Earth's surface as the remains of fossils of marine life in the past; they occur in two forms, rhombohedral calcite (more common) and orthorhombic aragonite (formed in more temperate seas). First types of minerals include limestone, dolomite, marble, lime, and Iceland; aragonite beds form the Bahamas, the Florida Keys, and the Red Sea basin. Coral, sea shells, and pearls are mostly made of calcium carbonate. Among other important minerals of calcium are gypsum (CaSO 4 Ã, Â · 2H 2 O), anhydride (CaSO 4 ), fluorite CaF 2 ), and apatite ([Ca 5 (PO 4 ) 3 F]).
The main producers of calcium are China (about 10,000 to 12,000 tons per year), Russia (about 6000 to 8,000 tons per year), and the United States (about 2000 to 4000 tons per year). Canada and France also include small producers. In 2005, about 24,000 tons 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 annually. In Russia and China, Davy's electrolysis method is still used, but is applied to liquid calcium chloride. Because calcium is less reactive than strontium or barium, the oxide-nitride coating that produces stable air and lathe machining and other standard metallurgical techniques is suitable for calcium. In the United States and Canada, calcium is produced by reducing lime with aluminum metal at high temperatures.
Cycling geochemistry
Calcium provides a link between tectonic, climatic, and carbon cycles. In the simplest terms, lifting the mountains exposes calcium-containing stones to chemical weathering and releases Ca 2 into surface water. These ions are transported into the ocean where they react with dissolved CO 2 to form limestone ( CaCO
3 ), which in turn settles to the seafloor where it is inserted into the new stone. The dissolved CO 2 , together with carbonate and bicarbonate ions, is called "dissolved inorganic carbon" (DIC).
Reactions are actually more complicated and involve the bicarbonate ions (HCO - 3 ) formed when CO 2 reacts with water at sea water pH:
- Ca 2 2 HCO -
3 -> CaCO
3 (s) CO
style = "font-size: inherit; line-height: inherit; vertical-align: baseline"> 2 H
2 O
At sea water pH, most CO 2 is immediately converted back to HCO -
3 . The reaction results in the net transport of one CO 2 molecule from the ocean/atmosphere to the lithosphere. The result is that every Ca 2 ion released by chemical weathering ultimately removes one CO 2 molecule from the surficial system (atmosphere, oceans, soil and living organisms), deposits it in rock carbonate where it is likely to last for hundreds of millions of years. Calcium weathering from rocks thus rubs CO 2 from the ocean and atmosphere, exerting a strong long-term effect on climate.
Usage
The greatest use of calcium in steelmaking, because of its strong chemical affinity for oxygen and sulfur. The oxide and sulphide, once formed, give the alumina the liquid chalk and sulfide inclusions in the steel floating out; on treatment, these inclusions spread throughout the steel and become small and round, increasing castability, cleanliness and general mechanical properties. Calcium is also used in maintenance-free automotive batteries, where the use of a 0.1% calcium-lead alloy in lieu of lead-antimonium alloys leads to lower water loss and lower self-use. Due to the risk of expansion and cracking, aluminum is sometimes also incorporated into this alloy. The tin-calcium alloy is also used in casting, replacing the tin antimony alloy. Calcium is also used to strengthen the aluminum alloy used for bearings, to control graphite carbon in cast iron, and to remove dirt bismuth from lead. Calcium metals are found in some duct cleaners, where it functions to produce heat and calcium hydroxide that paralyze the fat and melt proteins (eg, those with hair) that clog the canal. In addition to metallurgy, calcium reactivity is exploited to remove nitrogen from high purity argon gas and as a picker 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, because it reacts with hydrogen to form a solid calcium hydride, from which hydrogen can be easily re-extracted.
Calcium isotopic fraction during mineral formation has led to several applications of calcium isotope. In particular, the 1997 observations by Skulan and DePaolo that isotope mineral calcium is lighter than the solution from which the mineral settles is the basis of analog applications in medicine and paleooceanography. In animals with mineralized skeletons with calcium, soft tissue isotope calcium compositions reflect the relative degree of formation and dissolution of skeletal minerals. In humans, changes in urinary calcium isotope composition have been shown to be associated with changes in bone mineral balance. When the rate of bone formation exceeds the rate of bone resorption, the ratio of 44 Ca/ 40 Ca in the soft tissue increases and vice versa. Because of this relationship, measurement of urinary or blood calcium isotopes may be useful in early detection of metabolic bone disease such as osteoporosis. Similar systems exist in seawater, where 44 Ca/ 40 Ca tends to rise when the removal rate of Ca 2 by mineral precipitation exceeds the new calcium input into oceans. In 1997 the Sculan and DePaolo presented the first evidence of alteration in seawater 44 Ca/ 40 Ca during geological time, along with the theoretical explanations of this change. Newer papers have confirmed these observations, suggesting that the concentration of Ca 2 sea water is not constant, and that the oceans are never in "steady state" with respect to calcium input and output. This has important climatological implications, since the marine calcium cycle is closely related to the carbon cycle.
Many calcium compounds are used in foods, such as medicines, and in medicines, among others. For example, calcium and phosphorus are added in the diet through the addition of calcium lactate, calcium diphosphate, and tricalcium phosphate. The latter is also used as a polishing agent in toothpaste and antacids. Calcium lactobionate is a white powder used as an agent suspending drugs. In baking, calcium monophosphate is used as a yeast agent. Calcium sulfite is used as a bleach in paper making and as a disinfectant, calcium silicate is used as a reinforcing agent in rubber, and calcium acetate is the limiting component of resin and is used to make metallic soaps and synthetic resins.
Biological and pathological roles
Calcium is an important element that is needed in large quantities. The Ca 2 ion acts as an electrolyte and is vital to the health of the muscular, circulatory, and digestive systems; very necessary to build bones; and supports the synthesis and function of blood cells. For example, it regulates muscle contraction, nerve conduction, and blood clotting. As a result, the levels of intra- and extracellular calcium are strictly regulated by the body. Calcium can play this role because Ca 2 ions form a stable coordination complex with many organic compounds, especially proteins; it also forms compounds with various solubilities, enabling the formation of skeletons.
Calcium ions can be complexed by proteins by binding of carboxyl groups of glutamic acid or aspartic acid residues; through interaction with phosphorylated serine, tyrosine, or threonine; or by chelated by amino acid carboxylic residues. Trypsin, a digestive enzyme, uses the first method; osteocalcin, bone matrix protein, using a third. Some other bone matrix proteins such as osteopontin and bone sialoprotein use the first and the second. Direct activation of enzymes by binding calcium is common; some other enzymes are activated by noncovalent associations with a direct calcium-binding enzyme. Calcium also binds to the cell membrane phospholipid layer, binding to proteins associated with the cell surface. As an example of the various solubility of calcium compounds, monocalcium phosphate is very soluble in water, 85% of extracellular calcium is as dicalcium phosphate with 2.0 mM solubility and bone hydroxyapatite in organic matrix is ​​tricalcium phosphate at 100 ÂμM.
About three quarters of the dietary calcium comes from dairy products and whole grains, the rest accounted for by vegetables, protein-rich foods, fruits, sugars, fats, and oils. Calcium supplements are controversial, because calcium bioavailability is highly dependent on salt solubility involved: calcium citrate, malate, and lactate are highly bioavailable while oxalates are much less. The intestine absorbs about one-third of the calcium eaten as free ions, and plasma calcium levels are then regulated by the kidneys. Parathyroid hormone and vitamin D promote bone formation by allowing and increasing the deposition of calcium ions there, allowing rapid bone turnover without affecting bone mass or mineral content. When plasma calcium levels decrease, cell surface receptors are activated and parathyroid hormone secretion occurs; followed by stimulating the entry of calcium into the plasma pool by taking it from the kidneys, intestines, and targeted bone cells, by the action of bone-forming parathyroid hormone to be antagonized by calcitonin, whose secretion increases with increasing plasma calcium levels.
Excessive calcium intake can cause hypercalcaemia, but because of the inefficient absorption of calcium by the intestine, the more likely cause is excessive vitamin D intake or excessive secretion of parathyroid hormone. This can also happen because of bone damage that occurs when the tumor metastasizes to the bone. This results in the deposition of calcium salts to the heart, blood vessels, and kidneys. Symptoms include anorexia, nausea, vomiting, memory loss, confusion, muscle weakness, increased urination, dehydration, and metabolic bone disease. Chronic hypercalcaemia may cause soft tissue calcification, which can lead to serious consequences: for example, calcification of blood vessel walls can lead to loss of elasticity and laminar blood flow disruption, and from there to plaque rupture and thrombosis. Similarly, inadequate calcium or vitamin D intake causes hypocalcemia, often due to inadequate secretion of parathyroid hormones or receptors in cells. Symptoms include neuromuscular stimuli, potentially leading to tetany and defects in cardiac conduction.
Because calcium is heavily involved in bone making, many bone diseases can be traced to problems with organic or hydroxyapatite matrices in molecular structures or organizations. For example, osteoporosis is the reduction of bone mineral content per unit volume, and can be treated with calcium, vitamin D, and bisphosphate supplementation. Calcium supplements may be useful for serum lipids in women who have passed menopause as well as older men; calcium supplementation in postmenopausal women also seems to correlate inversely with cardiovascular disease. Insufficient quantities of calcium, vitamin D, or phosphate can cause bone softening, known as osteomalacia.
Security
Because calcium reacts exothermally with water and acids, the contact of calcium metal with moisture produces severe corrosive irritation. When ingested, calcium metal has the same effect on the mouth, throat, and stomach, and can be fatal. However, long-term exposure is not known to have different side effects.
Because of concerns of long-term side effects such as calcification of arteries and kidney stones, the US Institute of Medicine (IOM) and the European Food Safety Authority (EFSA) both set Tolerable Upper Intake Levels (ULs) for a combination of diet and extra calcium. From IOM, people aged 9-18 years should not exceed 3,000 mg/day; for ages 19-50 not exceeding 2,500 mg/day; for ages 51 and older, not exceeding 2,000 mg/day. EFSA establishes UL at 2,500 mg/day for adults but decides information for children and adolescents is not enough to determine UL.
See also
References
Bibliography
- Greenwood, Norman N.; Earnshaw, Alan (1997). Chemical Elements (2nd ed.). Butterworth-Heinemann. ISBN: 0-08-037941-9.
- Hluchan, Stephen E.; Pomerantz, Kenneth (2005), "Calcium and Calcium Alloys", Ullmann's Encyclopedia of Industrial Chemistry , Weinheim: Wiley-VCH, doi: 10.1002/14356007.a04_515.pub2
External links
- WebElements.com - Calcium
- Calcium in Periodic Video Table (University of Nottingham)
- Nutrition fact sheet from National Institutes of Health
Source of the article : Wikipedia
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