CALCIUM

Description

Calcium is the chemical element with the symbol Ca and atomic number 20. It has an atomic weight of 40.078 g/mol. Calcium is the fifth most abundant dissolved ion in seawater by both molarity and mass, after Na⁺, Cl^-, Mg²⁺ and SO4^2-. Calcium metal is quite reactive. It readily forms a white coating of calcium nitride (Ca3N2) in air at room temperature. It reacts with water and the metal burns in air with an orangered flame, forming largely the nitride, but some oxide. In the visible portion of the spectrum of many stars, including the Sun, strong absorption lines of singly ionized calcium ions are evident. Prominent among these are the H line at 3968.5 A ° and the K line at 3933.7 A°of singly ionized calcium, or Ca II. For the Sun, and stars with low temperatures, the prominence of the H and K lines is an indication of strong magnetic activity in the chromosphere. Measurement of periodic variations of these active regions can also be used to deduce the rotational periods of these stars.

Physical properties

Before the nineteenth century, calcium and some other alkali earth metals were consideredminerals rather than metals because they readily formed hydroxides and, thus, were classedas alkaline substances (bases). Calcium and other alkali earth metals were mined as ores.Therefore, early chemists considered them earth elements.Calcium metal is moderately hard and has a lustrous, silvery color when freshly cut, but itsoon oxidizes to a dull gray. It is a good conductor of heat and electricity, but has few uses inthe electronics industries. Calcium is an important element for the nutrition of living organisms, particularly vertebrates. Vitamin D from foods, sunlight, and milk aids in the depositionof calcium in bones and teeth. The elemental metal is very reactive in water as it releases hydrogen from the water. It is somewhat less reactive in air. Its melting point is 839°C, its boilingpoint is 1484°C, its density is 1.54 g/cm³, and it has a cubic crystal structure.

Isotopes

There are 20 isotopes of calcium, ranging from Ca-35 to Ca-54. Of the sixstable isotopes, Ca-40 makes up 96.941% of the calcium found in the Earth’s crust, andCa-42 = 0.647%, Ca-43 = 0.135%, Ca-44 = 2.086%, Ca-46 = 0.004%, and Ca-48 =0.187% found on Earth. All the other isotopes of calcium are radioactive and are artificially produced with half-lives ranging from a few microseconds to 1×10⁵ years.Radioactive Ca-45 emits beta particles (high-speed electrons) and has a half-life ofabout 163 days. It is used to determine the calcium levels in bones and in soils.

Isotopes

Calcium has four stable isotopes (⁴⁰Ca and ⁴²Ca through ⁴⁴Ca), plus two more isotopes (⁴⁶Ca and ⁴⁸Ca) that have such a long half-lives that for all practical purposes they can be considered stable. It also has a cosmogenic isotope, radioactive ⁴⁴Ca, with a half-life of 103,000 years. Unlike cosmogenic isotopes that are produced in the atmosphere, ⁴¹Ca is produced by neutron activation of ⁴⁰Ca. Most of its production is in the upper meter or so of the soil column, where the cosmogenic neutron flux is still sufficiently strong. 41Ca has received much attention in stellar studies because it decays to 41K, a critical indicator of solar system anomalies. A major part (97%) of naturally occurring calcium is in the form of 40Ca. This isotope is one of the daughter products of 40K decay, along with 40Ar. While K/Ar dating has been used extensively in the geological sciences, the prevalence of ⁴⁰Ca in nature has impeded its use in dating. Techniques using mass spectrometry and a “double-spike” isotope dilution have been used for K/Ca age dating.
The most abundant isotope, ⁴⁰Ca, has a nucleus of 20 protons and 20 neutrons. This is the heaviest stable isotope of any element known that has equal numbers of protons and neutrons. In supernova explosions, calcium is formed from the reaction of carbon with various numbers of alpha particles (helium nuclei), until this most common calcium isotope (containing 10 helium nuclei) has been formed.

Origin of Name

Its modern name was derived from the word calcis or calx, which is Latin for “lime.

Occurrence

Calcium is the fifth most abundant element found in the Earth’s crust. It is not found asa free element, but as calcium compounds (mostly salts and oxides), which are found on alllandmasses of the world as limestone, marble, and chalk. Calcium, particularly as the compound calcium chloride (CaCl₂), is found in the oceans to the extent of 0.15%.Calcium is produced by two methods. One method is the electrolysis of calcium chloride(Ca++ + 2Cl- → CaCl₂) as the electrolyte at a temperature of ?800°C, during which processmetallic calcium cations (Ca++) are deposited at the cathode as elemental calcium metal.Calcium can also be produced through a thermal process under very low pressure (vacuum)in which lime is reduced by using aluminum.In addition to limestone, calcium is also found in other rocks, coral, shells, eggshells, bones,teeth, and stalactites and stalagmites.

History

Though lime was prepared by the Romans in the first century under the name calx, Calcium was not discovered until 1808. After learning that Berzelius and Pontin prepared calcium amalgam by electrolyzing lime in mercury, Davy was able to isolate the impure metal. Calcium is a metallic element, fifth in abundance in the Earth’s crust, of which it forms more than 3%. It is an essential constituent of leaves, bones, teeth, and shells. Never found in nature uncombined, it occurs abundantly as limestone (CaCO3), gypsum (CaSO4·2H2O), and fluorite (CaF2); apatite is the fluorophosphate or chlorophosphate of calcium. Calcium has a silvery color, is rather hard, and is prepared by electrolysis of the fused chloride to which calcium fluoride is added to lower the melting point. Chemically it is one of the alkaline earth elements; it readily forms a white coating of oxide in air, reacts with water, burns with a yellowred flame, largely forming the oxide. The metal is used as a reducing agent in preparing other metals such as thorium, The Elements 4-7 uranium, zirconium, etc., and is used as a deoxidizer, desulfurizer, and inclusion modifier for various ferrous and nonferrous alloys. It is also used as an alloying agent for aluminum, beryllium, copper, lead, and magnesium alloys, and serves as a “getter” for residual gases in vacuum tubes. Its natural and prepared compounds are widely used. Quicklime (CaO), made by heating limestone and changed into slaked lime by the careful addition of water, is the great cheap base of the chemical industry with countless uses. Mixed with sand it hardens as mortar and plaster by taking up carbon dioxide from the air. Calcium from limestone is an important element in Portland cement. The solubility of the carbonate in water containing carbon dioxide causes the formation of caves with stalactites and stalagmites and is responsible for hardness in water. Other important compounds are the carbide (CaC2), chloride (CaCl2), cyanamide (CaCN2), hypochlorite (Ca(OCl)2), nitrate (Ca(NO3)2), and sulfide (CaS). Calcium sulfide is phosphorescent after being exposed to light. Natural calcium contains six isotopes. Sixteen other radioactive isotopes are known. Metallic calcium (99.5%) costs about $200/kg.

Characteristics

Finely powdered calcium metal is flammable in air because it liberates hydrogen from themoisture. It can be extremely reactive in water but can be dissolved in acids. Calcium is harderthan sodium metal, but softer than aluminum. In its elemental form it can be machined (cuton a lathe), extruded (pushed through a die), and drawn (stretched into rods or wires).Calcium is present in a number of products used in our everyday lives. It is found inclassroom chalk, teeth, and bones. About 2% of the human body consists of various forms ofcalcium compounds. Calcium is an essential inorganic element (usually in compound form)for plant and animal life.

History

Calcium (Latin word calcis meaning “lime”) was known as early as the first century when the Ancient Romans prepared lime as calcium oxide from limestone and used “slaked lime” as a “whitewash” on various homes and buildings. Literature dating back to 975 AD notes that “Plaster-of-Paris” (calcium sulfate), is useful for setting broken bones. The metal was not isolated until 1808 when Sir Humphrey Davy of England electrolyzed a mixture of lime and mercuric oxide, using then the new “Voltaic Cell” as an energy source. Davy was trying to isolate calcium. When he heard that Swedish chemist J?ns Berzelius and his colleague, Pontin, had prepared calcium amalgam by electrolyzing lime in mercury, he set up a similar system and was successful. He worked with electrolysis throughout his life and also discovered/isolated boron, sodium, potassium, magnesium, calcium, strontium and barium. Calcium metal was not available on a large scale until the beginning of the twentieth century.

The Uses of CALCIUM

Calcium is an alkaline earth element that contributes toward bone and teeth formation, muscle contraction, and blood clotting. it occurs in milk, vegetables, and egg yolk.

The Uses of CALCIUM

Calcium oxide was used in ancient times to make mortar for building with stone. Boththe metal and calcium compounds have many industrial as well as biological uses. Metalliccalcium is used as an alloy agent for copper and aluminum. It is also used to purify lead andis a reducing agent for beryllium.It is used to remove carbon and sulfur impurities during the processing of iron, producinga higher-grade iron for use in the manufacture of steel. It is also used as a reducing agent inthe preparation of several other important metals.Calcium is an important ingredient in the diets of all plants and animals. It is found in thesoft tissues and fluids of animals (e.g., blood) as well as in bones and teeth. Calcium makes upabout 2% of human body weight.Calcium is the main ingredient of Portland cement and is used to reduce the acid contentof soils.

The Uses of CALCIUM

Calcium, plasma standard solution is used as a standard solution in analytical chemistry and atomic absorption spectroscopy. It is also used as a single-element standard solution for plasma emission spectrometry.

Definition

A moderately soft, low-melting reactive metal; the third element in group 2 (formerly IIA) of the periodic table. The electronic configuration is that of argon with an additional pair of 4s electrons.
Calcium is widely distributed in the Earth’s crust and is the third most abundant element. Large deposits occur as chalk or marble, CaCO₃; gypsum, CaSO₄. 2H₂O; anhydrite, CaSO₄; fluorspar, CaF; and apatite, CaF₂.Ca₃(PO₄)₃. However sufficiently large quantities of calcium chloride are available as waste from the Solvay process to satisfy industrial requirements for the metal, which is produced by electrolysis of the fused salt. Large quantities of lime, Ca(OH)₂, and quicklime, CaO, are produced by decomposition of the carbonate for use in both building and agriculture. Several calcium minerals are mined as a source of other substances. Thus, limestone is a cheap source of carbon dioxide, gypsum and anhydrite are used in the manufacture of sulfuric acid, phosphate rock for phosphoric acid, and fluorspar for a range of fluorochemicals.
Calcium has a low ionization potential and a relatively large atomic radius. It is therefore a very electropositive element. The metal is very reactive and the compounds contain the divalent ion Ca²⁺. Calcium forms the oxide (CaO), a white ionic solid, on burning in air, but for practical purposes the oxide is best prepared by heating the carbonate, which decomposes at about 800°C. Both the oxide and the metal itself react with water to give the basic hydroxide (Ca(OH)₂). On heating with nitrogen, sulfur, or the halogens, calcium reacts to form the nitride (Ca₃N₂), sulfide (CaS), or the halides (CaX₂). Calcium also reacts directly with hydrogen to give the hydride CaH₂ and borides, arsenides, carbides, and silicides can be prepared in a similar way. Both the carbonate and sulfate are insoluble. Calcium salts impart a characteristic brick-red color to flames which is an aid to qualitative analysis. At ordinary temperatures calcium has the face-centered cubic structure with a transition at 450°C to the close-packed hexagonal structure.

Definition

calcium: Symbol Ca. A soft greymetallic element belonging togroup 2 (formerly IIA) of the periodictable; a.n. 20; r.a.m. 40.08; r.d.1.54; m.p. 839°C; b.p. 1484°C. Calciumcompounds are common in theearth’s crust; e.g. limestone and marble(CaCO₃), gypsum (CaSO₄.2H₂O),and fluorite (CaF₂). The element is extractedby electrolysis of fused calciumchloride and is used as a getterin vacuum systems and a deoxidizerin producing nonferrous alloys. It isalso used as a reducing agent in theextraction of such metals as thorium,zirconium, and uranium.
Calcium is an essential element forliving organisms, being required fornormal growth and development. Inanimals it is an important constituentof bones and teeth and ispresent in the blood, being requiredfor muscle contraction and othermetabolic processes. In plants it is aconstituent (in the form of calciumpectate) of the middle lamella.

Indications

Calcium is the principal extracellular electrolyte regulated by PTH, calcitonin, and D3. Extracellular calcium is a critical component of signal transduction across the plasma membrane, which regulates a wide spectrum of physiological events including muscle contraction, secretion of neurotransmitters and hormones, and the action of growth factors, cytokines, and protein hormones. Intracellular calcium is an important cofactor in many enzymatic reactions.

Hazard

The metallic form of calcium, particularly the powdered form, combines with water oroxidizing agents to release hydrogen that may explode, as do the other alkali metals. There aremany useful calcium compounds; some are excellent reducing agents, some are explosive, andothers are essential for life.
The radioactive isotope calcium-45 is deposited in bones and teeth as well as other plantand animal tissues. Because our bodies cannot distinguish between Ca-45 and the stable Ca-40, the radioactive isotope Ca-45 is used as a tracer to study diseased bone and tissue. At thesame time, a massive overexposure to Ca-45 can displace the stable form of Ca-40 in animalsand can cause radiation sickness or even death.
A few calcium compounds, when in powder or vapor form, are toxic when ingested orinhaled.

Agricultural Uses

Calcium (Ca), an element essential for the growth of living organisms, is a soft grey metallic element belonging to Group 2 (formerly Ⅱ A) of the Periodic Table. It is an important constituent of bones and teeth in animals and is present in blood. It is required for muscle contraction and other metabolic processes. In plants, it is a constituent of the middle lamella and is essential for plant cell elongation, cell division and various anabolic and catabolic processes.
Calcium is an essential secondary plant nutrient. It is absorbed by plant roots as calcium ions and supplied to the root surface by mass flow and root interception. In cells, calcium levels are micromolar, whereas externally they are millimolar, facilitating a major role in metabolic control. Calcium levels directly control many enzyme reactions and enhance the uptake of nitrate, thus becoming interrelated with nitrogen metabolism.
Calcium also facilitates the uptake of potassium ions preferentially over sodium ions. Therefore, an optimum ratio of potassium to calcium is important for a favorable water balance in plants.
Calcium is generally an immobile element in plants.
Calcium concentration usually averages to about 33 ppm in the soil solution. A level of 15 ppm is adequate for a high corn yield. A higher concentration or excess of calcium in soil can cause deficiency of potassium and some micronutrients, particularly zinc. Crops like alfalfa, cabbage, Potato and Sugar beet9 need higher amounts of calcium.
Calcium content varies widely in soil. Plagioclase mineral anorthite is the most important primary source of calcium. The availability of calcium from soil to plants is determined by many factors, such as the total calcium content, soil pH, cation exchange capacity (CEC), calcium percentage saturation, type of soil colloid, ratio of calcium to other cations in solution, etc. Highly acidic soils impede the uptake of calcium ions. As the calcium percentage saturation decreases relative to the total CEC, the amount of calcium ions absorbed by the plant decreases.
Many crops respond to calcium applications when the exchangeable calcium ion saturation falls below 25%. The 2: 1 clays (meaning, a lattice with 2 silica sheets for 1 alumina sheet) require a higher calcium ion saturation than 1 : 1 clays (a lattice with 1 silica sheet for 1 alumina sheet). An increased aluminum ion concentration in the soil solution reduces calcium uptake by crops such as corn, cotton, soybean and wheat. The calcium supply for most crops is considered adequate when the soil calcium to cation ratio is between 0.10 and 0.15. Ammonium, potassium, magnesium, manganese and aluminum ions suppress the calcium ion uptake by plants, whereas nitrate ions enhance it.
Calcium deficiency leads to a decrease in the development of terminal shoot buds and apical root tips, thereby inhibiting plant growth. Calcium deficiency also causes distortion of new leaves; in maize, for instance, the deficiency prevents the emergence and unfolding of new leaves and renders the tips of the existing leaves almost colorless. A sticky gelatinous material causes the leaves to adhere to one another.
Calcium deficiency is rare in fruits (excepting apples which show pitting of flesh and skin) and vegetables; however, it is seen as a disorder in the storage tissues.
The quantity of calcium lost by leaching is 75 to 200 kg/ha/yr. This loss is greater than that of sodium ions as the quantity of calcium present in a solution is higher and is in an exchangeable form.
Calcium is also lost or neutralized from the soil by other routes. These are (a) rapid neutralization by acidforming fertilizers (ammonium), (b) slow neutralization by acid formed by carbon dioxide in water due to decomposition of organic matter, (c) erosion, leaving more acidic subsoil to be limed, (d) slow removal of harvested or grazed crops, and (e) slow neutralization by acidic rain.
Calcium is added to plants as a component of other nutrients, particularly phosphorus, as in superphosphates. Primary sources of liming materials are calcium carbonate and magnesium carbonate. Gypsum is used for supplementing calcium as a nutrient without correcting soil acidity. Synthetic chelates, such as ethylenediaminetetraacetic acid (EDTA) complexes of calcium containing 3 to 5% calcium, are used as fertilizers both for soil and foliar applications.
The uptake of calcium is progressively reduced by the presence of ammonium, magnesium, potassium and sodium ions, in that order. The foliar calcium to magnesium ratio of 2: 1 and potassium to calcium ratio of 4:l are considered optimum for plant growth. Phosphorus favors calcium uptake under acidic conditions. A pH increase after excessive liming results indeficiencies of iron, manganese, boron or zinc, leading to subsequent chlorosis. Calcium and boron exhibit a synergistic effect in reducing plant disorders.

Pharmaceutical Applications

Calcium is mostly found in limestone and its related forms, such as chalk, and marble and lime (CaO). Calcium is the most abundant inorganic element in the human body and is an essential key for many physiological processes. Calcium ions are a critical factor in several life-defining biochemical processes as well as in the endocrine, neural and renal aspects of blood pressure homeostasis. Calcium has the symbol Ca and atomic number 20 and is a soft grey alkaline earth metal. Calcium has four stable isotopes (⁴⁰Ca and ⁴²Ca–⁴⁴Ca) and the metal reacts with water with the formation of calcium hydroxide and hydrogen.

Industrial uses

Calcium (symbol Ca) is a metallic elementbelonging to the group of alkaline earths. It isone of the most abundant materials, occurringin combination in limestones and calcareousclays. The metal is obtained 98.6% pure byelectrolysis of the fused anhydrous chloride. Byfurther subliming, it is obtained 99.5% pure.Calcium metal is yellowish white in color. Itoxidizes easily and, when heated in air, burnswith a brilliant white light. It has a density of1.55 g/cm3, a melting point of 838°C, and aboiling point of 1440°C. Its strong affinity for O2 and sulfur is utilized as a cleanser for nonferrous alloys. As a deoxidizer and desulfurizerit is employed in the form of lumps or sticks ofcalcium metal or in ferroalloys and Ca–Cu.
Many compounds of calcium are employedindustrially, in fertilizers, foodstuffs, and medicine.It is an essential element in the formationof bones, teeth, shells, and plants. Oyster shellsform an important commercial source of calciumfor animal feeds. They are crushed, andthe fine flour is marketed for stock feeds andthe coarse for poultry feeds. The shell is calciumcarbonate.