Period 4 elements

OverviewPeriod 4 elementA period 4 element is one of the chemical elements in the fourth row (or period) of the periodic table of the elements. The periodic table is laid out in rows to illustrate recurring (periodic) trends in the chemical behaviour of the elements as their atomic number increases: a new row is begun when chemical behaviour begins to repeat, meaning that elements with similar behaviour fall into the same vertical columns. The fourth period contains 18 elements, beginning with potassium and ending with krypton. As a rule, period 4 elements fill their 4s shells first, then their 3d and 4p shells, in that order, however there are exceptions, such as chromium.PropertiesEvery single one of these elements are stable,1 "List of Elements of the Periodic Table Sorted by Abundance in Earth's crust". Science.co.il. . Retrieved 2012-08-14.

and many are extremely common in the earth's crust and/or core. Many of the transition metals in period 4 are incredibly strong, and therefore commonly used in industry, especially iron. Copper is one of three elements that are not silver or gray in color, along with caesium and gold. Two adjacent elements are known to be toxic, with arsenic one of the most well-known poisons, and selenium being toxic to humans in large quantities. Many elements are essential to humans' survival, such as Calcium being what forms bones.2Gray, Theodore (2009). The Elements: A Visual Exploration of Every Known Atom in the Universe. New York: Black Dog & Leventhal Publishers. ISBN978-1-57912-814-2.

Periodic trends

Potassium

Calcium

Scandium

Titanium

Vanadium

Chromium

Manganese

Iron

Cobalt

Nickel

Copper

Zinc

Gallium

Germanium

Arsenic

Selenium

Bromine

Krypton

List of elementsChemical elementChemical seriesElectron configuration

19 KPotassiumAlkali metal [Ar] 4s1

20 CaCalciumAlkaline earth metal [Ar] 4s2

21 ScScandiumTransition metal [Ar] 3d1 4s2

22 TiTitaniumTransition metal [Ar] 3d2 4s2

23 VVanadiumTransition metal [Ar] 3d3 4s2

24 CrChromiumTransition metal [Ar] 3d5 4s1 (*)

25 MnManganeseTransition metal [Ar] 3d5 4s2

26 FeIronTransition metal [Ar] 3d6 4s2

27 CoCobaltTransition metal [Ar] 3d7 4s2

28 NiNickelTransition metal [Ar] 3d9 4s1 (*)

29 CuCopperTransition metal [Ar] 3d10 4s1 (*)

30 ZnZincTransition metal [Ar] 3d10 4s2

31 GaGalliumPost-transition metal [Ar] 3d10 4s2 4p1

32 GeGermaniumMetalloid [Ar] 3d10 4s2 4p2

33 AsArsenicMetalloid [Ar] 3d10 4s2 4p3

34 SeSeleniumNonmetal [Ar] 3d10 4s2 4p4

35 BrBromineHalogen [Ar] 3d10 4s2 4p5

36 KrKryptonNoble gas [Ar] 3d10 4s2 4p6

(*) Exception to the Madelung rule

Some element categories in the periodic table

Metals Metalloids Nonmetals Unknown
chemical
properties

Alkali
metals Alkaline
earth metals Inner transition metalsTransition
metals Post-transition
metals Other
nonmetals Halogens Noble
gases

Lanthanides Actinides

s-block elementsPotassium

Potassium (K) is an alkali metal, placed under sodium and over rubidium, and is the first element of period 4.3 "Elements in the Modern Periodic Table, Periodic Classification of Elements". Tutorvista.com. . Retrieved 2012-08-14.

It is one of the most reactive elements in the periodic table, therefore usually only found in compounds. It tends to oxidize in air very rapidly, thus accounting for its rapid reaction with oxygen when freshly exposed to air. When freshly exposed, it is rather silvery, but it quickly begins to tarnish as it reacts with air. It is soft enough to be cut with a knife and it is the second least dense element.4 "It's Elemental The Element Potassium". Education.jlab.org. . Retrieved 2012-08-14.

Potassium has a relatively low melting point; it will melt just by putting it under a small open flame.5 "Potassium, Chemical Element Overview, Discovery and naming, Physical properties, Chemical properties, Occurrence in nature, Isotopes". Chemistryexplained.com. . Retrieved 2012-08-14.

It also is less dense than water, and can, in turn, float.6 "Potassium (K) Chemical properties, Health and Environmental effects". Lenntech.com. . Retrieved 2012-08-14.

Calcium

Calcium (Ca) is the second element in period 4, between potassium and scandium. An alkali earth metal, calcium is almost never found in nature due to its high reactivity with water.7 "Reactions of the Group 2 elements with water". Chemguide.co.uk. . Retrieved 2012-08-14.

It has one of the most widely known and acknowledged biological roles in all animals and some plants, making up bones and teeth, and used in some applications in cells, such as signals for cellular processeses. It is regarded as the most abundant mineral in the body's mass.8 "Chapter 11. Calcium". Fao.org. . Retrieved 2012-08-14.

d-block elementsScandium

Scandium (Sc) is the third element in period 4, between calcium and titanium, and is the first transition metal in the periodic table. Scandium is quite common in nature, but difficult to find because it is most prevalent in rare earth compounds, which are difficult to isolate elements from. Scandium has very few commercial applications because of the aforementioned facts, and currently its only major application is in aluminium alloys. Titanium

Titanium (Ti) is an element in period 4, between scandium and vanadium. Titanium is well known for being one of the least dense, strong, and corrosion-resistant elements, and as such has many applications, especially in alloys with other elements, such as iron. Due to its aforementioned properties, it is commonly used in airplanes, golf clubs, and other objects that must be strong, but lightweight. Vanadium

Vanadium (V) is an element in period 4, between titanium and chromium. Vanadium is never found in pure form in nature, but is commonly found in compounds. Vanadium is similar to titanium in many ways, such as being very corrosion-resistant, however, unlike titanium, it oxidizes in air even at room temperature. All vanadium compounds have at least some level of toxicity, with some of them being extremely toxic. Chromium

Chromium (Cr) is an element in period 4, between vanadium and manganese. Chromium is, like titanium and vanadium before it, extremely resistant to corrosion, and is indeed one of the main components of stainless steel. Chromium also has many colorful compounds, and as such is very commonly used in pigments, such as chrome green. Manganese

Manganese (Mn) is an element in period 4, between chromium and iron. Manganese is often found free in nature, but is also found in combination with iron. Manganese, like chromium before it, is an important component in stainless steel, preventing the iron from rusting. Manganese is also often used in pigments, again like chromium. Manganese is also poisonous; if enough is inhaled, it can cause irreversible neurological damage. Iron

Iron (Fe) is an element in period 4, between manganese and cobalt. Iron is probably the most well-known element in period 4, being the most common element in the earth and a major component of steel. Iron-56 has the lowest energy density of any isotope of any element, meaning that it is the most massive element that can be produced in supergiant stars. Iron also has some applications in the human body; hemoglobin is partly iron. Cobalt

Cobalt (Co) is an element in period 4, between iron and nickel. Cobalt is commonly used in pigments, as many compounds of cobalt are blue in color. Cobalt is also a core component of many magnetic and high-strength alloys. The only stable isotope, cobalt-59, is an important component of vitamin B-12, while cobalt-60 is a component of nuclear fallout and can be dangerous in large enough quantities due to its radioactivity. Nickel

Nickel (Ni) is an element in period 4, between cobalt and copper. Nickel is rare in the earth's crust, mainly due to the fact that it reacts with oxygen in the air, with most of the nickel on earth coming from nickel iron meteorites. However, nickel is incredibly common in the earths core; along with iron it is one of the two main components. Nickel is an important component of stainless steel, and in many superalloys. Copper

Copper (Cu) is an element in period 4, between nickel and zinc. Copper is one of the few metals that is not white or gray in color, the only others being gold and caesium. Copper has been used by humans for thousands of years to provide a reddish tint to many objects, and is even an essential nutrient to humans, although too much is poisonous. Copper is also commonly used as a wood preservative or fungicides. Zinc

Zinc (Zn) is an element in period 4, between copper and gallium. Zinc is one of the main components of brass, being used since the 10th century BCE. Zinc is also incredibly important to humans; almost 2 billion people in the world suffer from zinc deficiency. However, too much zinc can cause copper deficiency. Zinc is often used in batteries, aptly named carbon-zinc batteries, and is important in many platings, as zinc is very corrosion resistant. p-block elementsGallium

Gallium (Ga) is an element in period 4, between zinc and germanium. Gallium is noteworthy because it has a melting point at about 303 Kelvin, right around room temperature. For example, it will be solid on a typical spring day, but will be liquid on a hot summer day. Gallium is an important component in the alloy galinstan, along with tin. Gallium can also be found in semiconductors. Germanium

Germanium (Ge) is an element in period 4, between gallium and arsenic. Germanium, like silicon above it, is an important semiconductor and is commonly used in diodes and transistors, often in combination with arsenic. Germanium is fairly rare on earth, leading to its comparatively late discovery. Germanium, in compounds, can sometimes irritate the eyes, skin, or lungs. Arsenic

Arsenic (As) is an element in period 4, between germanium and selenium. Arsenic, as mentioned above, is often used in semiconductors in alloys with germanium. Arsenic, in pure form and some alloys, is incredibly poisonous to all multicellular life, and as such is a common component in pesticides to get rid of bugs. Arsenic was also used in some pigments before its toxicity was discovered. Selenium

Selenium (Se) is an element in period 4, between arsenic and bromine. Selenium is the first nonmetal in period 4, with properties similar to sulfer. Selenium is quite rare in pure form in nature, mostly being found in minerals such as pyrite, and even then it is quite rare. Selenium is necessary for humans in trace amounts, but is toxic in larger quantities. Bromine

Bromine (Br) is an element in period 4, between selenium and krypton. Bromine is a halogen, never existing in pure form in nature. Bromine is barely liquid at room temperature, boiling at about 330 Kelvin. Bromine is also quite toxic and corrosive, but bromide ions, which are relatively inert, can be found in halite, or table salt. Bromine is often used as a fire retardant because many compounds can be made to release free bromine atoms. Krypton

Krypton (Kr) is a noble gas, placed under argon and over xenon. Being a noble gas, krypton rarely interacts with itself or other elements; although compounds have been detected, they are all unstable and decay rapidly, and as such, krypton is often used in fluorescent lights. Krypton, like most noble gases, is also used in lighting because of its many spectral lines and the aforementioned reasons. ElementsPotassiumPotassium

Appearance

silvery gray
Potassium pearls under paraffin oil. The large pearl measures 0.5 cm. Below: spectral lines of potassium

General properties

Name, symbol, number potassium, K, 19

Pronunciation/ptsim/ po-TAS-ee-m

Element categoryalkali metal

Group, period, block1,4, s

Standard atomic weight39.0983(1)gmol1

Electron configuration [Ar] 4s1

Electrons per shell 2, 8, 8, 1 (Image)

Physical properties

Phasesolid

Density (near r.t.) 0.862 gcm3

Liquid density at m.p. 0.828 gcm3

Melting point 336.53K,63.38C,146.08F

Boiling point 1032K,759C,1398F

Triple point 336.35K(63C),kPa

Heat of fusion 2.33 kJmol1

Heat of vaporization 76.9 kJmol1

Specific heat capacity (25 C) 29.6 Jmol1K1

Atomic properties

Oxidation states 1
(strongly basic oxide)

Electronegativity 0.82 (Pauling scale)

Ionization energies
(more) 1st: 418.8 kJmol1

2nd: 3052 kJmol1

3rd: 4420 kJmol1

Atomic radius227 pm

Covalent radius20312 pm

Van der Waals radius275 pm

Miscellanea

Crystal structure body-centered cubic

Magnetic ordering paramagnetic

Electrical resistivity (20C) 72 nm

Thermal conductivity (300 K) 102.5Wm1K1

Thermal expansion (25 C) 83.3 mm1K1

Speed of sound (thin rod) (20 C) 2000 m/s

Young's modulus 3.53 GPa

Shear modulus 1.3 GPa

Bulk modulus 3.1 GPa

Mohs hardness 0.4

Brinell hardness 0.363 MPa

CAS registry number 7440-09-7

Most stable isotopes

isoNAhalf-lifeDMDE (MeV)DP

39K 93.26% 39K is stable with 20 neutron

40K 0.012% 1.248(3)109 y 1.311 40Ca

1.505 40Ar

+ 1.505 40Ar

41K 6.73% 41K is stable with 22 neutron

Potassium (/ptsim/ po-TAS-ee-m) is a chemical element with symbolK (from Neo-Latin kalium) and atomic number19. Elemental potassium is a soft silvery-white alkali metal that oxidizes rapidly in air and is very reactive with water, generating sufficient heat to ignite the hydrogen emitted in the reaction.Because potassium and sodium are chemically very similar, it took a long time before their salts were differentiated. The existence of multiple elements in their salts was suspected from 1702,1Marggraf, Andreas Siegmund (1761). Chymische Schriften. p.167. .

and this was proven in 1807 when potassium and sodium were individually isolated from different salts by electrolysis. Potassium in nature occurs only in ionic salts. As such, it is found dissolved in seawater (which is 0.04% potassium by weight2Webb, D. A. (April 1939). "The Sodium and Potassium Content of Sea Water". The Journal of Experimental Biology: 183. .

3Anthoni, J. (2006). "Detailed composition of seawater at 3.5% salinity". seafriends.org.nz. . Retrieved 23 September 2011.

), and is part of many minerals. Most industrial chemical applications of potassium employ the relatively high solubility in water of potassium compounds, such as potassium soaps. Potassium metal has only a few special applications, being replaced in most chemical reactions with sodium metal. Potassium ions are necessary for the function of all living cells. Potassium ion diffusion is a key mechanism in nerve transmission, and potassium depletion in animals, including humans, results in various cardiac dysfunctions. Potassium is found in especially high concentrations within plant cells, and in a mixed diet it is mostly concentrated in fruits. The high concentration of potassium in plants, associated with comparatively low amounts of sodium there, resulted in potassium's being first isolated from potash, the ashes of plants, giving the element its name. For the same reason, heavy crop production rapidly depletes soils of potassium, and agricultural fertilizers consume 95% of global potassium chemical production.4Greenwood 1997, p.73

PropertiesPhysicalThe flame test of potassium

Potassium atoms have 19 electrons, which is one more than the extremely stable configuration of argon. A potassium atom is thus much more likely to lose the "extra" electron than to gain one; however, the alkalide ions, K, are known.5Dye, J. L. (1979). "Compounds of Alkali Metal Anions". Angewandte Chemie International Edition 18 (8): 587598. doi:10.1002/anie.197905871.

Because of the low first ionization energy (418.8 kJ/mol) the potassium atom easily loses an electron and oxidizes into the monopositive cation, K+.6James, A. M.; Lord, M. P. (1992). Macmillan's chemical and physical data. London: Macmillan. ISBN0-333-51167-0.

This process requires so little energy that potassium is readily oxidized by atmospheric oxygen. In contrast, the second ionization energy, is very high (3052 kJ/mol), because removal of two electrons breaks the stable noble gas electronic configuration.7

Potassium therefore does not readily form compounds with the oxidation state of +2 (or higher).8

Potassium is the second least dense metal after lithium. It is a soft solid that has a low melting point and can easily be cut with a knife. Freshly cut potassium is silvery in appearance, but it begins to tarnish toward gray immediately after being exposed to air.9Greenwood 1997, p.76

10Holleman, Arnold F.; Wiberg, Egon; Wiberg, Nils (1985). "Potassium" (in German). Lehrbuch der Anorganischen Chemie (91100 ed.). Walter de Gruyter. ISBN3-11-007511-3.

In a flame test, potassium and its compounds emit a lilac color with a peak emission wavelength of 766.5nm (see movie below).11Greenwood 1997, p.75

12

Chemical Potassium is an extremely active metal, which reacts violently with oxygen and water in air. With oxygen, it converts to potassium peroxide and with water potassium hydroxide. The reaction of potassium with water is dangerous because of its violent exothermic character and the production of hydrogen gas. Hydrogen reacts again with atmospheric oxygen, producing water, which reacts with the remaining potassium.13

This reaction requires only traces of water; because of this, potassium and its liquid alloy with sodium NaK are potent desiccants that can be used to dry solvents prior to distillation.14

15Burkhardt, p. 35

Because of the sensitivity of potassium to water and air, the reactions are possible only in inert atmosphere, such as argon gas using air-free techniques. Potassium does not react with most hydrocarbons, such as mineral oil or kerosene.16

It readily dissolves in liquid ammonia, up to 480 g per 1000 g of ammonia at 0 C. Depending on the concentration, the ammonia solutions are blue to yellow, and their electrical conductivity is similar to that of liquid metals. In a pure solution, potassium slowly reacts with ammonia to form KNH2, but this reaction is accelerated by minute amounts of transition metal salts.17Burkhardt, p. 32

It can reduce the salts to the metal; potassium is often used as the reductant in the preparation of finely divided metals from their salts by the Rieke method.18Rieke, R. D. (1989). "Preparation of Organometallic Compounds from Highly Reactive Metal Powders". Science 246 (4935): 12601264. Bibcode1989Sci...246.1260R. doi:10.1126/science.246.4935.1260. PMID17832221.

For example, the preparation of Rieke magnesium employs potassium as the reductant: MgCl2 + 2 K Mg + 2 KCl

Compounds The only common oxidation state for potassium is +1. Potassium metal is a powerful reducing agent that is easily oxidized to the monopositive cation, K+. Once oxidized, it is very stable and difficult to reduce back to the metal.19

Potassium hydroxide reacts readily with carbon dioxide to produce potassium carbonate, and is used to remove traces of the gas from air. In general, potassium compounds have excellent water solubility, owing to the high hydration energy of the K+ ion. The potassium ion is colorless in water and is very difficult to precipitate; possible precipitation methods include reactions with sodium tetraphenylborate, hexachloroplatinic acid, and sodium cobaltinitrite.20

Potassium oxidizes faster than most metals and forms oxides with oxygen-oxygen bonds, as do all alkali metals except lithium. Three species are formed during the reaction: potassium oxide, potassium peroxide, and potassium superoxide,21Lide, David R. (1998). Handbook of Chemistry and Physics (87 ed.). Boca Raton, Florida, United States: CRC Press. pp.477; 520. ISBN0-8493-0594-2.

which contain three different oxygen-based ions: oxide (O2), peroxide (O), and superoxide (O). The last two species, especially the superoxide, are rare and are formed only in reaction with very electropositive metals; these species contain oxygenoxygen bonds.22

All potassiumoxygen binary compounds are known to react with water violently, forming potassium hydroxide. This compound is a very strong alkali, and 1.21kg of it can dissolve as much as a liter of water.23Lide, D. R., ed. (2005). CRC Handbook of Chemistry and Physics (86th ed.). Boca Raton (FL): CRC Press. p.4-80. ISBN0-8493-0486-5.

24Schultz, p. 94

Structure of solid potassium superoxide (KO2).

In aqueous solution Potassium compounds are typically highly ionic and thus most of them are soluble in water. The main species in water are the aquo complexes [K(H2O)n]+ where n = 6 and 7.25S. F. Lincoln, D. T. Richens, A. G. Sykes "Metal Aqua Ions" Comprehensive Coordination Chemistry II Volume 1, Pages 515-555. doi:10.1016/B0-08-043748-6/01055-0

Some of the few salts that are poorly soluble include potassium tetraphenylborate, potassium hexachloroplatinate, and potassium cobaltinitrite.26

27Hyde, Earl K. (1960). The radiochemistry of thorium. Subcommittee on Radiochemistry, National Academy of SciencesNational Research Council. .

Isotopes There are 24 known isotopes of potassium, three of which occur naturally: 39K (93.3%), 40K (0.0117%), and 41K (6.7%). Naturally occurring 40K has a half-life of 1.250109 years. It decays to stable 40Ar by electron capture or positron emission (11.2%) or to stable 40Ca by beta decay (88.8%).28Georges, Audi (2003). "The NUBASE Evaluation of Nuclear and Decay Properties". Nuclear Physics A (Atomic Mass Data Center) 729: 3128. Bibcode2003NuPhA.729....3A. doi:10.1016/j.nuclphysa.2003.11.001.

The decay of 40K to 40Ar enables a commonly used method for dating rocks. The conventional K-Ar dating method depends on the assumption that the rocks contained no argon at the time of formation and that all the subsequent radiogenic argon (i.e., 40Ar) was quantitatively retained. Minerals are dated by measurement of the concentration of potassium and the amount of radiogenic 40Ar that has accumulated. The minerals that are best suited for dating include biotite, muscovite, metamorphic hornblende, and volcanic feldspar; whole rock samples from volcanic flows and shallow instrusives can also be dated if they are unaltered.29

30Bowen, Robert; Attendorn, H. G. (1988). "Theory and Assumptions in PotassiumArgon Dating". Isotopes in the Earth Sciences. Springer. pp.203208. ISBN978-0-412-53710-3. .

Outside of dating, potassium isotopes have been used as tracers in studies of weathering and for nutrient cycling studies because potassium is a macronutrient required for life.31D. Anac; P. Martin-Prvel (1 August 1999). Improved crop quality by nutrient management. Springer. pp.290. ISBN978-0-7923-5850-3. . Retrieved 20 June 2011.

40K occurs in natural potassium (and thus in some commercial salt substitutes) in sufficient quantity that large bags of those substitutes can be used as a radioactive source for classroom demonstrations. In healthy animals and people, 40K represents the largest source of radioactivity, greater even than 14C. In a human body of 70kg mass, about 4,400 nuclei of 40K decay per second.32 "Radiation and Radioactive Decay. Radioactive Human Body". Harvard Natural Sciences Lecture Demonstrations. . Retrieved 2011-05-18.

The activity of natural potassium is 31 Bq/g.33Winteringham, F. P. W; Effects, F.A.O. Standing Committee on Radiation, Land And Water Development Division, Food and Agriculture Organization of the United Nations (1989). Radioactive fallout in soils, crops and food: a background review. Food & Agriculture Org.. p.32. ISBN978-92-5-102877-3. .

Creation and occurrencePotassium in feldspar

Potassium is formed in the universe by nucleosynthesis from lighter atoms. The stable form of potassium is created in supernovas via the explosive Oxygen-burning process.34Shimansky, V. (September 2003). "Observational constraints on potassium synthesis during the formation of stars of the Galactic disk". Astronomy Reports. . Retrieved 2011-08-07.

Elemental potassium does not occur in nature because it reacts violently with water (see section Precautions below).35

As various compounds, potassium makes up about 2.6% of the weight of the Earth's crust and is the seventh most abundant element, similar in abundance to sodium at approximately 1.8% of the crust.36Greenwood 1997, p.69

In seawater, potassium at 0.39g/L37

(0.039 wt/v%) is far less abundant than sodium at 10.8g/L (1.08 wt/v%).38Micale, Giorgio; Cipollina, Andrea; Rizzuti, Lucio (2009). Seawater Desalination: Conventional and Renewable Energy Processes. Springer. p.3. ISBN978-3-642-01149-8. .

39Prud'homme, Michel; Krukowski, Stanley T. (2006). "Potash". Industrial minerals & rocks: commodities, markets, and uses. Society for Mining, Metallurgy, and Exploration. pp.723740. ISBN978-0-87335-233-8. .

Orthoclase (potassium feldspar) is a common rock-forming mineral. Granite for example contains 5% potassium, which is well above the average in the Earth's crust. Sylvite (KCl), carnallite (KClMgCl26(H2O)), kainite (MgSO4KCl3H2O) and langbeinite (MgSO4K2SO4)) are the minerals found in large evaporite deposits worldwide. The deposits often show layers starting with the least soluble at the bottom and the most soluble on top.40

Deposits of niter (potassium nitrate) are formed by decomposition of organic material in contact with atmosphere, mostly in caves; because of the good water solubility of niter the formation of larger deposits requires special environmental conditions.41Ross, William H. (1914). "The Origin of Nitrate Deposits". Popular Science. Bonnier Corporation. pp.134145. .

History Neither elemental potassium nor potassium salts (as separate entities from other salts) were known in Roman times, and the Latin name of the element is not Classical Latin but rather neo-Latin. The Latin name kalium was taken from the word "alkali", which in turn came from Arabic: al-qalyah "plant ashes." The similar-sounding English term alkali is from this same root (potassium in Modern Standard Arabic is btsym).Humphry Davy

The English name for the element potassium comes from the word "potash",42Davy, Humphry (1808). "On some new phenomena of chemical changes produced by electricity, in particular the decomposition of the fixed alkalies, and the exhibition of the new substances that constitute their bases; and on the general nature of alkaline bodies". Philosophical Transactions of the Royal Society of London 98: 32. doi:10.1098/rstl.1808.0001. .

referring to the method by which potash was obtained leaching the ash of burnt wood or tree leaves and evaporating the solution in a pot. Potash is primarily a mixture of potassium salts because plants have little or no sodium content, and the rest of a plant's major mineral content consists of calcium salts of relatively low solubility in water. While potash has been used since ancient times, it was not understood for most of its history to be a fundamentally different substance from sodium mineral salts. Georg Ernst Stahl obtained experimental evidence that led him to suggest the fundamental difference of sodium and potassium salts in 1702,43

and Henri Louis Duhamel du Monceau was able to prove this difference in 1736.44du Monceau, H. L. D.. "Sur la Base de Sel Marine" (in French). Memoires de l'Academie royale des Sciences: 6568. .

The exact chemical composition of potassium and sodium compounds, and the status as chemical element of potassium and sodium, was not known then, and thus Antoine Lavoisier did include the alkali in his list of chemical elements in 1789.45Weeks, Mary Elvira (1932). "The discovery of the elements. IX. Three alkali metals: Potassium, sodium, and lithium". Journal of Chemical Education 9 (6): 1035. Bibcode1932JChEd...9.1035W. doi:10.1021/ed009p1035.

46Siegfried, R. (1963). "The Discovery of Potassium and Sodium, and the Problem of the Chemical Elements". Isis 54 (2): 247258. doi:10.1086/349704. JSTOR228541.

Potassium metal was first isolated in 1807 in England by Sir Humphry Davy, who derived it from caustic potash (KOH), by the use of electrolysis of the molten salt with the newly discovered voltaic pile. Potassium was the first metal that was isolated by electrolysis.47Enghag, P. (2004). "11. Sodium and Potassium". Encyclopedia of the elements. Wiley-VCH Weinheim. ISBN3-527-30666-8.

Later in the same year, Davy reported extraction of the metal sodium from a mineral derivative (caustic soda, NaOH, or lye) rather than a plant salt, by a similar technique, demonstrating that the elements, and thus the salts, are different.48

49

50Davy, Humphry (1808). "On some new phenomena of chemical changes produced by electricity, in particular the decomposition of the fixed alkalies, and the exhibition of the new substances that constitute their bases; and on the general nature of alkaline bodies". Philosophical Transactions of the Royal Society of London 98: 144. doi:10.1098/rstl.1808.0001. .

51Shaposhnik, V. A. (2007). "History of the discovery of potassium and sodium (on the 200th anniversary of the discovery of potassium and sodium)". Journal of Analytical Chemistry 62 (11): 11001102. doi:10.1134/S1061934807110160.

Although the production of potassium and sodium metal should have shown that both are elements, it took some time before this view was universally accepted.52

For a long time the only significant applications for potash were the production of glass, bleach, and soap.53Browne, C. A. (1926). "Historical notes upon the domestic potash industry in early colonial and later times". Journal of Chemical Education 3 (7): 749756. Bibcode1926JChEd...3..749B. doi:10.1021/ed003p749.

Potassium soaps from animal fats and vegetable oils were especially prized, as they tended to be more water-soluble and of softer texture, and were known as soft soaps.54

The discovery by Justus Liebig in 1840 that potassium is a necessary element for plants and that most types of soil lack potassium55Liebig, Justus von (1840) (in German). Die organische Chemie in ihrer Anwendung auf Agricultur und Physiologie. .

caused a steep rise in demand for potassium salts. Wood-ash from fir trees was initially used as a potassium salt source for fertilizer, but, with the discovery in 1868 of mineral deposits containing potassium chloride near Stafurt, Germany, the production of potassium-containing fertilizers began at an industrial scale.56Cordel, Oskar (1868) (in German). Die Stassfurter Kalisalze in der Landwirthschalt: Eine Besprechung .... L. Schnock. . Retrieved 29 May 2011.

57Birnbaum, Karl (1869) (in German). Die Kalidngung in ihren Vortheilen und Gefahren. .

58Organization, United Nations Industrial Development; Center, Int'l Fertilizer Development (1998-03-31). Fertilizer Manual. ISBN978-0-7923-5032-3. .

59Organization, United Nations Industrial Development; Center, Int'l Fertilizer Development (1998-03-31). Fertilizer Manual. ISBN978-0-7923-5032-3. .

Other potash deposits were discovered, and by the 1960s Canada became the dominant producer.60Miller, H. (1980). "Potash from Wood Ashes: Frontier Technology in Canada and the United States". Technology and Culture 21 (2): 187208. doi:10.2307/3103338. JSTOR3103338.

61Rittenhouse, P. A. (1979). "Potash and politics". Economic Geology 74 (2): 353357. doi:10.2113/gsecongeo.74.2.353.

Commercial productionSylvite from New Mexico

Potassium salts such as carnallite, langbeinite, polyhalite, and sylvite form extensive deposits in ancient lake and seabeds,62

making extraction of potassium salts in these environments commercially viable. The principal source of potassium potash is mined in Canada, Russia, Belarus, Germany, Israel, United States, Jordan, and other places around the world.63Garrett, Donald E. (1995-12-31). Potash: deposits, processing, properties and uses. Springer. ISBN978-0-412-99071-7. .

64Ober, Joyce A.. "Mineral Commodity Summaries 2008:Potash". United States Geological Survey. . Retrieved 2008-11-20.

65Ober, Joyce A.. "Mineral Yearbook 2006:Potash". United States Geological Survey. . Retrieved 2008-11-20.

The first mined deposits were located near Stafurt, Germany, but the deposits span from Great Britain over Germany into Poland. They are located in the Zechstein and were deposited in the Middle to Late Permian. The largest deposits ever found lie 1000 meters (3000feet) below the surface of the Canadian province of Saskatchewan. The deposits are located in the Elk Point Group produced in the Middle Devonian. Saskatchewan, where several large mines have operated since the 1960s, pioneered the use of freezing of wet sands (the Blairmore formation) in order to drive mine shafts through them. The main potash mining company in Saskatchewan is the Potash Corporation of Saskatchewan.66Wishart, David J (2004). Encyclopedia of the Great Plains. U of Nebraska Press. p.433. ISBN978-0-8032-4787-1. .

The water of the Dead Sea is used by Israel and Jordan as a source for potash, while the concentration in normal oceans is too low for commercial production.67

68

Mining and beneficiation waste heaps from potash mining in Germany, consisting mostly of sodium chloride.

Several methods are applied to separate the potassium salts from the present sodium and magnesium compounds. The most-used method is to precipitate some compounds relying on the solubility difference of the salts at different temperatures. Electrostatic separation of the ground salt mixture is also used in some mines. The resulting sodium and magnesium waste is either stored underground or piled up in slag heaps. Most of the mined potassium minerals end up as potassium chloride after processing. The mineral industry refers to potassium chloride either as potash, muriate of potash, or simply MOP.69

Pure potassium metal can be isolated by electrolysis of its hydroxide in a process that has changed little since Davy. Although the electrolysis process was developed and used in industrial scale in the 1920s the thermal method by reacting sodium with potassium chloride in a chemical equilibrium reaction became the dominant method in the 1950s. The production of sodium potassium alloys is possible by changing the reaction time and the amount of sodium used in the reaction. The Griesheimer process employing the reaction of potassium fluoride with calcium carbide was also used to produce potassium.70

71. doi:10.1002/0471238961.161520010308092.

Na + KCl NaCl + K (Thermal method)

2 KF + CaC2 2K + CaF2 + 2 C (Griesheimer process)

Reagent-grade potassium metal cost about $10.00/pound ($22/kg) in 2010 when purchased in tonne quantities. Lower purity metal is considerably cheaper. The market is volatile due to the difficulty of the long-term storage of the metal. It must be stored under a dry inert gas atmosphere or anhydrous mineral oil to prevent the formation of a surface layer of potassium superoxide. This superoxide is a pressure-sensitive explosive that will detonate when scratched. The resulting explosion will usually start a fire that is difficult to extinguish.72Burkhardt, p. 34

73Delahunt, J; Lindeman, T (2007). "Review of the safety of potassium and potassium oxides, including deactivation by introduction into water". Journal of Chemical Health and Safety 14 (2): 2132. doi:10.1016/j.jchas.2006.09.010.

Biological roleBiochemical functionThe action of the sodium-potassium pump is an example of primary active transport. The two carrier proteins on the left are using ATP to move sodium out of the cell against the concentration gradient. The proteins on the right are using secondary active transport to move potassium into the cell.

Potassium is the eighth or ninth most common element by mass (0.2%) in the human body, so that a 60kg adult contains a total of about 120g of potassium.74Abdelwahab, M.; Youssef, S.; Aly, A.; Elfiki, S.; Elenany, N.; Abbas, M. (1992). "A simple calibration of a whole-body counter for the measurement of total body potassium in humans". International Journal of Radiation Applications and Instrumentation. Part A. Applied Radiation and Isotopes 43 (10): 12851289. doi:10.1016/0883-2889(92)90208-V.

The body has about as much potassium as sulfur and chlorine, and only the major minerals calcium and phosphorus are more abundant.75Chang, Raymond (1 July 2007). Chemistry. McGraw-Hill Higher Education. p.52. ISBN978-0-07-110595-8. . Retrieved 29 May 2011.

Potassium cations are important in neuron (brain and nerve) function, and in influencing osmotic balance between cells and the interstitial fluid, with their distribution mediated in all animals (but not in all plants) by the so-called Na+/K+-ATPase pump.76Campbell, Neil (1987). Biology. Menlo Park, California: Benjamin/Cummings Pub. Co.. p.795. ISBN0-8053-1840-2.

This ion pump uses ATP to pump three sodium ions out of the cell and two potassium ions into the cell, thus creating an electrochemical gradient over the cell membrane. In addition, the highly selective potassium ion channels (which are tetramers) are crucial for the hyperpolarization, in for example neurons, after an action potential is fired. The most recently resolved potassium ion channel is KirBac3.1, which gives a total of five potassium ion channels (KcsA, KirBac1.1, KirBac3.1, KvAP, and MthK) with a determined structure.77Hellgren, Mikko; Sandberg, Lars; Edholm, Olle (2006). "A comparison between two prokaryotic potassium channels (KirBac1.1 and KcsA) in a molecular dynamics (MD) simulation study". Biophysical Chemistry 120 (1): 19. doi:10.1016/j.bpc.2005.10.002. PMID16253415.

All five are from prokaryotic species. Potassium can be detected by taste because it triggers three of the five types of taste sensations, according to concentration. Dilute solutions of potassium ions taste sweet, allowing moderate concentrations in milk and juices, while higher concentrations become increasingly bitter/alkaline, and finally also salty to the taste. The combined bitterness and saltiness of high-potassium solutions makes high-dose potassium supplementation by liquid drinks a palatability challenge.78Institute of Medicine (U.S.). Committee on Optimization of Nutrient Composition of Military Rations for Short-Term, High-Stress Situations; Institute of Medicine (U.S.). Committee on Military Nutrition Research (2006). Nutrient composition of rations for short-term, high-intensity combat operations. National Academies Press. pp.287. ISBN978-0-309-09641-6. . Retrieved 29 May 2011.

79Shallenberger, R. S. (1993). Taste chemistry. Springer. pp.120. ISBN978-0-7514-0150-9. . Retrieved 29 May 2011.

Membrane polarization Potassium is also important in preventing muscle contraction and the sending of all nerve impulses in animals through action potentials. By nature of their electrostatic and chemical properties, K+ ions are larger than Na+ ions, and ion channels and pumps in cell membranes can distinguish between the two types of ions, actively pumping or passively allowing one of the two ions to pass, while blocking the other.80Lockless, S. W.; Zhou, M.; MacKinnon, R. (2007). "Structural and thermodynamic properties of selective ion binding in a K+ channel". PLoS Biol 5 (5): e121. doi:10.1371/journal.pbio.0050121. PMC1858713. PMID17472437.

A shortage of potassium in body fluids may cause a potentially fatal condition known as hypokalemia, typically resulting from vomiting, diarrhea, and/or increased diuresis.81Slonim, Anthony D.; Pollack, Murray M. (2006). "Potassium". Pediatric critical care medicine. Lippincott Williams & Wilkins. p.812. ISBN978-0-7817-9469-5. .

Deficiency symptoms include muscle weakness, paralytic ileus, ECG abnormalities, decreased reflex response and in severe cases respiratory paralysis, alkalosis and cardiac arrhythmia.82Visveswaran, Kasi (2009). "hypokalemia". Essentials of Nephrology, 2/e. BI Publications. p.257. ISBN978-81-7225-323-3. .

Filtration and excretion Potassium is an essential macromineral in human nutrition; it is the major cation (positive ion) inside animal cells, and it is thus important in maintaining fluid and electrolyte balance in the body. Sodium makes up most of the cations of blood plasma at a reference range of about 145 mmol/L (3.345g)(1mmol/L = 1mEq/L), and potassium makes up most of the cell fluid cations at about 150 mmol/L (4.8g). Plasma is filtered through the glomerulus of the kidneys in enormous amounts, about 180liters per day.83Potts, W. T. W.; Parry, G. (1964). Osmotic and ionic regulation in animals. Pergamon Press.

Thus 602g of sodium and 33g of potassium are filtered each day. All but the 110g of sodium and the 14g of potassium likely to be in the diet must be reabsorbed. Sodium must be reabsorbed in such a way as to keep the blood volume exactly right and the osmotic pressure correct; potassium must be reabsorbed in such a way as to keep serum concentration as close as possible to 4.8 mmol/L (about 0.190g/L).84Lans, H. S.; Stein, I. F.; Meyer, KA (1952). "The relation of serum potassium to erythrocyte potassium in normal subjects and patients with potassium deficiency". American Journal of Medical Science 223 (1): 6574. doi:10.1097/00000441-195201000-00011. PMID14902792.

Sodium pumps in the kidneys must always operate to conserve sodium. Potassium must sometimes be conserved also, but, as the amount of potassium in the blood plasma is very small and the pool of potassium in the cells is about thirty times as large, the situation is not so critical for potassium. Since potassium is moved passively85Bennett, C. M.; Brenner, B. M.; Berliner, R. W. (1968). "Micropuncture study of nephron function in the rhesus monkey". Journal of Clinical Investigation 47 (1): 203216. doi:10.1172/JCI105710. PMC297160. PMID16695942.

86Solomon, A. K. (1962). "Pumps in the living cell". Scientific American 207 (2): 1008. doi:10.1038/scientificamerican0862-100. PMID13914986.

in counter flow to sodium in response to an apparent (but not actual) Donnan equilibrium,87Kernan, Roderick P. (1980). Cell potassium (Transport in the life sciences). New York: Wiley. pp.40, 48. ISBN0-471-04806-2.

the urine can never sink below the concentration of potassium in serum except sometimes by actively excreting water at the end of the processing. Potassium is secreted twice and reabsorbed three times before the urine reaches the collecting tubules.88Wright, F. S. (1977). "Sites and mechanisms of potassium transport along the renal tubule". Kidney International 11 (6): 415432. doi:10.1038/ki.1977.60. PMID875263.

At that point, it usually has about the same potassium concentration as plasma. At the end of the processing, potassium is secreted one more time if the serum levels are too high. If potassium were removed from the diet, there would remain a minimum obligatory kidney excretion of about 200mg per day when the serum declines to 3.03.5 mmol/L in about one week,89Squires, R. D.; Huth, E. J. (1959). "Experimental potassium depletion in normal human subjects. I. Relation of ionic intakes to the renal conservation of potassium". Journal of Clinical Investigation 38 (7): 11341148. doi:10.1172/JCI103890. PMC293261. PMID13664789.

and can never be cut off completely, resulting in hypokalemia and even death.90Nicholas H. Fiebach; Lee Randol Barker; John Russell Burton; Philip D. Zieve (2007). Principles of ambulatory medicine. Lippincott Williams & Wilkins. pp.748750. ISBN978-0-7817-6227-4. . Retrieved 20 June 2011.

The potassium moves passively through pores in the cell membrane. When ions move through pumps there is a gate in the pumps on either side of the cell membrane and only one gate can be open at once. As a result, approximately 100 ions are forced through per second. Pores have only one gate, and there only one kind of ion can stream through, at 10 million to 100 million ions per second.91Gadsby, D. C. (2004). "Ion transport: spot the difference". Nature 427 (6977): 795797. Bibcode2004Natur.427..795G. doi:10.1038/427795a. PMID14985745.; for a diagram of the potassium pores are viewed, see Miller, C (2001). "See potassium run". Nature 414 (6859): 2324. doi:10.1038/35102126. PMID11689922.

The pores require calcium in order to open92Jiang, Y.; Lee, A.; Chen, J.; Cadene, M.; Chait, B. T.; MacKinnon, R. (2002). "Crystal structure and mechanism of a calcium-gated potassium channel". Nature 417 (6888): 51522. doi:10.1038/417515a. PMID12037559.

although it is thought that the calcium works in reverse by blocking at least one of the pores.93Shi, N.; Ye, S.; Alam, A.; Chen, L.; Jiang, Y (2006). "Atomic structure of a Na+- and K+-conducting channel". Nature 440 (7083): 570574. Bibcode2006Natur.440..570S. doi:10.1038/nature04508. PMID16467789.; includes a detailed picture of atoms in the pump.

Carbonyl groups inside the pore on the amino acids mimic the water hydration that takes place in water solution94Zhou, Y.; Morais-Cabral, J. H.; Kaufman, A.; MacKinnon, R. (2001). "Chemistry of ion coordination and hydration revealed by a K+ channel-Fab complex at 2.0 A resolution". Nature 414 (6859): 4348. doi:10.1038/35102009. PMID11689936.

by the nature of the electrostatic charges on four carbonyl groups inside the pore.95Noskov, S. Y.; Bernche, S.; Roux, B. (2004). "Control of ion selectivity in potassium channels by electrostatic and dynamic properties of carbonyl ligands". Nature 431 (7010): 830834. Bibcode2004Natur.431..830N. doi:10.1038/nature02943. PMID15483608.

In dietAdequate intake A potassium intake sufficient to support life can in general be guaranteed by eating a variety of foods. Clear cases of potassium deficiency (as defined by symptoms, signs and a below-normal blood level of the element) are rare in healthy individuals. Foods rich in potassium include parsley, dried apricots, dried milk, chocolate, various nuts (especially almonds and pistachios), potatoes, bamboo shoots, bananas, avocados, soybeans, and bran, although it is also present in sufficient quantities in most fruits, vegetables, meat and fish.96 "Potassium Food Charts". Asia Pacific Journal of Clinical Nutrition. . Retrieved 2011-05-18.

Optimal intake Epidemiological studies and studies in animals subject to hypertension indicate that diets high in potassium can reduce the risk of hypertension and possibly stroke (by a mechanism independent of blood pressure), and a potassium deficiency combined with an inadequate thiamine intake has produced heart disease in rats.97Folis, R. H. (1942). "Myocardial Necrosis in Rats on a Potassium Low Diet Prevented by Thiamine Deficiency". Bull. Johns-Hopkins Hospital 71: 235.

There is some debate regarding the optimal amount of dietary potassium. For example, the 2004 guidelines of the Institute of Medicine specify a DRI of 4,000mg of potassium (100mEq), though most Americans consume only half that amount per day, which would make them formally deficient as regards this particular recommendation.98Grim, C. E.; Luft, F. C.; Miller, J. Z.; Meneely, G.R.; Battarbee, H. D.; Hames, C. G.; Dahl, L. K. (1980). "Racial differences in blood pressure in Evans County, Georgia: relationship to sodium and potassium intake and plasma renin activity". Journal of Chronicle Diseases 33 (2): 8794. doi:10.1016/0021-9681(80)90032-6. PMID6986391.

Likewise, in the European Union, in particular in Germany and Italy, insufficient potassium intake is somewhat common.99Karger, S. (2004). "Energy and nutrient intake in the European Union" (PDF). Annals of Nutrition and Metabolism 48 (2 (suppl)): 116. doi:10.1159/000083041. .

Italian researchers reported in a 2011 meta-analysis that a 1.64g higher daily intake of potassium was associated with a 21% lower risk of stroke.100D'Elia, L.; Barba, G.; Cappuccio, F.; Strazzullo, P. (2011). "Potassium Intake, Stroke, and Cardiovascular Disease: A Meta-Analysis of Prospective Studies". The Journal of the American College of Cardiology 57 (10): 12101219. doi:10.1016/j.jacc.2010.09.070.

Medical supplementation and disease Supplements of potassium in medicine are most widely used in conjunction with loop diuretics and thiazides, classes of diuretics that rid the body of sodium and water, but have the side-effect of also causing potassium loss in urine. A variety of medical and non-medical supplements are available. Potassium salts such as potassium chloride may be dissolved in water, but the salty/bitter taste of high concentrations of potassium ion make palatable high concentration liquid supplements difficult to formulate.101

Typical medical supplemental doses range from 10 mmol (400mg, about equal to a cup of milk or 6USfloz (180ml). of orange juice) to 20 mmol (800mg) per dose. Potassium salts are also available in tablets or capsules, which for therapeutic purposes are formulated to allow potassium to leach slowly out of a matrix, as very high concentrations of potassium ion (which might occur next to a solid tablet of potassium chloride) can kill tissue, and cause injury to the gastric or intestinal mucosa. For this reason, non-prescription supplement potassium pills are limited by law in the US to only 99mg of potassium. Individuals suffering from kidney diseases may suffer adverse health effects from consuming large quantities of dietary potassium. End stage renal failure patients undergoing therapy by renal dialysis must observe strict dietary limits on potassium intake, as the kidneys control potassium excretion, and buildup of blood concentrations of potassium (hyperkalemia) may trigger fatal cardiac arrhythmia. ApplicationsFertilizerPotassium and magnesium sulfate fertilizer

Potassium ions are an essential component of plant nutrition and are found in most soil types.102

They are used as a fertilizer in agriculture, horticulture, and hydroponic culture in the form of chloride (KCl), sulfate (K2SO4), or nitrate (KNO3). Agricultural fertilizers consume 95% of global potassium chemical production, and about 90% of this potassium is supplied as KCl.103

The potassium content of most plants range from 0.5% to 2% of the harvested weight of crops, conventionally expressed as amount of K2O. Modern high-yield agriculture depends upon fertilizers to replace the potassium lost at harvest. Most agricultural fertilizers contain potassium chloride, while potassium sulfate is used for chloride-sensitive crops or crops needing higher sulfur content. The sulfate is produced mostly by decomposition of the complex minerals kainite (MgSO4KCl3H2O) and langbeinite (MgSO4K2SO4). Only a very few fertilizers contain potassium nitrate.104Roy, Amit H. (2007). Kent and Riegel's handbook of industrial chemistry and biotechnology. Springer. pp.11351157. ISBN978-0-387-27843-8. .

In 2005, about 93% of world potassium production was consumed by the fertilizer industry.105

Food The potassium cation is a nutrient necessary for human life and health. Potassium chloride is used as a substitute for table salt by those seeking to reduce sodium intake so as to control hypertension. The USDA lists tomato paste, orange juice, beet greens, white beans, potatoes, bananas and many other good dietary sources of potassium, ranked in descending order according to potassium content.106 "Potassium Content of Selected Foods per Common Measure, sorted by nutrient content". USDA National Nutrient Database for Standard Reference, Release 20. .

Potassium sodium tartrate (KNaC4H4O6, Rochelle salt) is the main constituent of baking powder; it is also used in the silvering of mirrors. Potassium bromate (KBrO3) is a strong oxidizer (E924), used to improve dough strength and rise height. Potassium bisulfite (KHSO3) is used as a food preservative, for example in wine and beer-making (but not in meats). It is also used to bleach textiles and straw, and in the tanning of leathers.107Figoni, Paula I (2010). "Bleaching and Maturing Agents". How Baking Works: Exploring the Fundamentals of Baking Science. John Wiley and Sons. p.86. ISBN978-0-470-39267-6. .

108Chichester, C. O. (1986-07). "Uses and Exposure to Sulfites in Food". Advances in food research. Academic Press. pp.46. ISBN978-0-12-016430-1. .

Industrial Major potassium chemicals are potassium hydroxide, potassium carbonate, potassium sulfate, and potassium chloride. Megatons of these compounds are produced annually.109H. Schultz, G. Bauer, Erich Schachl, F. Hagedorn, P. Schmittinger "Potassium Compounds" in Ullmann's Encyclopedia of Industrial Chemistry, 2005, Wiley-VCH, Weinheim. doi:10.1002/14356007.a22_039

Potassium hydroxide KOH is a strong base, which is used in industry to neutralize strong and weak acids, to control pH and to manufacture potassium salts. It is also used to saponify fats and oils, in industrial cleaners, and in hydrolysis reactions, for example of esters.110Toedt, John; Koza, Darrell; Cleef-Toedt, Kathleen Van (2005). "Personal Cleansing Products: Bar Soap". Chemical composition of everyday products. Greenwood Publishing Group. ISBN978-0-313-32579-3. .

111Schultz, p. 95

Potassium nitrate (KNO3) or saltpeter is obtained from natural sources such as guano and evaporites or manufactured via the Haber process; it is the oxidant in gunpowder (black powder) and an important agricultural fertilizer. Potassium cyanide (KCN) is used industrially to dissolve copper and precious metals, in particular silver and gold, by forming complexes. Its applications include gold mining, electroplating, and electroforming of these metals; it is also used in organic synthesis to make nitriles. Potassium carbonate (K2CO3 or potash) is used in the manufacture of glass, soap, color TV tubes, fluorescent lamps, textile dyes and pigments.112Schultz, p. 99

Potassium permanganate (KMnO4) is an oxidizing, bleaching and purification substance and is used for production of saccharin. Potassium chlorate (KClO3) is added to matches and explosives. Potassium bromide (KBr) was formerly used as a sedative and in photography.113

Potassium chromate (K2CrO4) is used in inks, dyes, stains (bright yellowish-red color); in explosives and fireworks; in the tanning of leather, in fly paper and safety matches,114Siegel, Richard S. (1940). "Ignition of the safety match". Journal of Chemical Education 17 (11): 515. Bibcode1940JChEd..17..515S. doi:10.1021/ed017p515.

but all these uses are due to the properties of chromate ion containment rather than potassium ions. Niche uses Potassium compounds are so pervasive that thousands of small uses are in place. The superoxide KO2 is an orange solid that acts as a portable source of oxygen and a carbon dioxide absorber. It is widely used in respiration systems in mines, submarines and spacecraft as it takes less volume than the gaseous oxygen.115Greenwood 1997, p.74

116Marx, Robert F. (1990). The history of underwater exploration. Courier Dover Publications. ISBN978-0-486-26487-5. .

4KO2 + 2CO2 2K2CO3 + 3O2

Potassium cobaltinitrite K3[Co(NO2)6] is used as artist's pigment under the name of Aureolin or Cobalt yellow.117Gettens, Rutherford John; Stout, George Leslie (1966). Painting materials: A short encyclopaedia. Courier Dover Publications. pp.109110. ISBN978-0-486-21597-6. .

Laboratory uses An alloy of sodium and potassium, NaK is a liquid used as a heat-transfer medium and a desiccant for producing dry and air-free solvents. It can also be used in reactive distillation.118Jackson, C. B.; Werner, R. C. (1957-01-01). 18. "Handling and uses of the alkali metals". Advances in Chemistry 19: 169173. doi:10.1021/ba-1957-0019.ch018. ISBN978-0-8412-0020-3. .

The ternary alloy of 12% Na, 47% K and 41% Cs has the lowest melting point of 78 C of any metallic compound.119

Metallic potassium is used in several types of magnetometers.120Kearey, Philip; Brooks, M; Hill, Ian (2002). "Optical Pumped Magnetometer". An introduction to geophysical exploration. Wiley-Blackwell. pp.164. ISBN978-0-632-04929-5. .

PrecautionsA reaction of potassium metal with water. Hydrogen is liberated that burns with a pink or lilac flame, the flame color owing to burning potassium vapor. Strongly alkaline potassium hydroxide is formed in solution.

Potassium reacts very violently with water producing potassium hydroxide (KOH) and hydrogen gas. 2 K (s) + 2 H2O (l) 2 KOH (aq) + H2 (g)

This reaction is exothermic and releases enough heat to ignite the resulting hydrogen. It in turn may explode in the presence of oxygen. Potassium hydroxide is a strong alkali that causes skin burns. Finely divided potassium will ignite in air at room temperature. The bulk metal will ignite in air if heated. Because its density is 0.89 g/cm3, burning potassium floats in water that exposes it to atmospheric oxygen. Many common fire extinguishing agents, including water, either are ineffective or make a potassium fire worse. Nitrogen, argon, Sodium chloride (table salt), sodium carbonate (soda ash), and silicon dioxide (sand) are effective if they are dry. Some Class D dry powder extinguishers designed for metal fires are also effective. These agents deprive the fire of oxygen and cool the potassium metal.121Solomon, Robert E. (2002). Fire and Life Safety Inspection Manual. Jones & Bartlett Learning. p.459. ISBN978-0-87765-472-8. .

Potassium reacts violently with halogens and will detonate in the presence of bromine. It also reacts explosively with sulfuric acid. During combustion potassium forms peroxides and superoxides. These peroxides may react violently with organic compounds such as oils. Both peroxides and superoxides may react explosively with metallic potassium.122 "DOE Handbook-Alkali Metals Sodium, Potassium, NaK, and Lithium". Hss.doe.gov. Archived from the original on 2010-09-28. . Retrieved 2010-10-16.

Because potassium reacts with water vapor present in the air, it is usually stored under anhydrous mineral oil or kerosene. Unlike lithium and sodium, however, potassium should not be stored under oil for longer than 6 months, unless in an inert (oxygen free) atmosphere, or under vacuum. After prolonged storage in air dangerous shock-sensitive peroxides can form on the metal and under the lid of the container, and can detonate upon opening.123Wray, Thomas K.. "Danger: peroxidazable chemicals". Environmental Health & Public Safety (North Carolina State University). .

Because of the highly reactive nature of potassium metal, it must be handled with great care, with full skin and eye protection and preferably an explosion-resistant barrier between the user and the metal. Ingestion of large amounts of potassium compounds can lead to hyperkalemia strongly influencing the cardiovascular system.124Schonwald, Seth (2004). "Potassium Chloride and Potassium Permanganate". Medical toxicology. Lippincott Williams & Wilkins. pp.903905. ISBN978-0-7817-2845-4. .

125Markovchick, Vincent J.; Pons, Peter T. (2003). Emergency medicine secrets. Elsevier Health Sciences. p.223. ISBN978-1-56053-503-4. .

Potassium chloride is used in the United States for death penalty via lethal injection.126

BibliographyBurkhardt, Elizabeth R. et al. (2006). "Potassium and Potassium Alloys". Ullmann's Encyclopedia of Industrial Chemistry. A22. pp.3138. doi:10.1002/14356007.a22_031.pub2. ISBN3-527-30673-0.

Greenwood, Norman N; Earnshaw, Alan (1997). Chemistry of the Elements (2 ed.). Oxford: Butterworth-Heinemann. ISBN0-08-037941-9.

Holleman, Arnold F.; Wiberg, Egon; Wiberg, Nils (1985). "Potassium" (in German). Lehrbuch der Anorganischen Chemie (91100 ed.). Walter de Gruyter. ISBN3-11-007511-3.

Schultz, H. et al. (2006). "Potassium compounds". Ullmann's Encyclopedia of Industrial Chemistry. A22. pp.39103. doi:10.1002/14356007.a22_031.pub2. ISBN3-527-30673-0.

CalciumCalcium

Appearance

dull gray, silver
Spectral lines of Calcium

General properties

Name, symbol, number calcium, Ca, 20

Pronunciation/klsim/ KAL-see-m

Element categoryalkaline earth metal

Group, period, block2,4, s

Standard atomic weight40.078(4)gmol1

Electron configuration [Ar] 4s2

Electrons per shell 2, 8, 8, 2 (Image)

Physical properties

Phasesolid

Density (near r.t.) 1.55 gcm3

Liquid density at m.p. 1.378 gcm3

Melting point 1115K,842C,1548F

Boiling point 1757K,1484C,2703F

Heat of fusion 8.54 kJmol1

Heat of vaporization 154.7 kJmol1

Specific heat capacity (25 C) 25.929 Jmol1K1

Vapor pressure

P/Pa 1 10 100 1 k 10 k 100 k

at T/K 864 956 1071 1227 1443 1755

Atomic properties

Oxidation states +2, +11Krieck, Sven; Grls, Helmar; Westerhausen, Matthias (2010). "Mechanistic Elucidation of the Formation of the Inverse Ca(I) Sandwich Complex [(thf)3Ca(-C6H3-1,3,5-Ph3)Ca(thf)3] and Stability of Aryl-Substituted Phenylcalcium Complexes". Journal of the American Chemical Society 132 (35): 100818110534020. doi:10.1021/ja105534w. PMID20718434.

(strongly basic oxide)

Electronegativity 1.00 (Pauling scale)

Ionization energies
(more) 1st: 589.8 kJmol1

2nd: 1145.4 kJmol1

3rd: 4912.4 kJmol1

Atomic radius197 pm

Covalent radius17610 pm

Van der Waals radius231 pm

Miscellanea

Crystal structure face-centered cubic

Magnetic orderingdiamagnetic

Electrical resistivity (20C) 33.6 nm

Thermal conductivity (300 K) 201Wm1K1

Thermal expansion (25 C) 22.3 mm1K1

Speed of sound (thin rod) (20 C) 3810 m/s

Young's modulus 20 GPa

Shear modulus 7.4 GPa

Bulk modulus 17 GPa

Poisson ratio 0.31

Mohs hardness 1.75

Brinell hardness 167 MPa

CAS registry number 7440-70-2

Most stable isotopes

isoNAhalf-lifeDMDE (MeV)DP

40Ca 96.941% 40Ca is stable with 20 neutron

41Ca trace1.03105 y - 41K

42Ca 0.647% 42Ca is stable with 22 neutron

43Ca 0.135% 43Ca is stable with 23 neutron

44Ca 2.086% 44Ca is stable with 24 neutron

45Ca syn162.7 d 0.258 45Sc

46Ca 0.004% >2.81015 y ? 46Ti

47Ca syn4.536 d 0.694, 1.99 47Sc

1.297 -

48Ca 0.187% >41019 y ? 48Ti

Calcium (/klsim/ KAL-see-m) is the chemical element with symbolCa and atomic number20. Calcium is a soft gray alkaline earth metal, and is the fifth-most-abundant element by mass in the Earth's crust. Calcium is also the fifth-most-abundant dissolved ion in seawater by both molarity and mass, after sodium, chloride, magnesium, and sulfate.2Dickson, A. G. and Goyet, C. (1994). "5". Handbook of method for the analysis of the various parameters of the carbon dioxide system in sea water, version 2. ORNL/CDIAC-74. .

Calcium is essential for living organisms, in particular in cell physiology, where movement of the calcium ion Ca2+ into and out of the cytoplasm functions as a signal for many cellular processes. As a major material used in mineralization of bone, teeth and shells, calcium is the most abundant metal by mass in many animals. Notable characteristicsFlame test. Brick-red color originates from calcium.

In chemical terms, calcium is reactive and soft for a metal (though harder than lead, it can be cut with a knife with difficulty). It is a silvery metallic element that must be extracted by electrolysis from a fused salt like calcium chloride.3Pauling, Linus (1970). General Chemistry. Dover Publications. p.627. ISBN0-7167-0149-9.

Once produced, it rapidly forms a gray-white oxide and nitride coating when exposed to air. In bulk form (typically as chips or "turnings"), the metal is somewhat difficult to ignite, more so even than magnesium chips; but, when lit, the metal burns in air with a brilliant high-intensity orange-red light. Calcium metal reacts with water, evolving hydrogen gas at a rate rapid enough to be noticeable, but not fast enough at room temperature to generate much heat. In powdered form, however, the reaction with water is extremely rapid, as the increased surface area of the powder accelerates the reaction with the water. Part of the slowness of the calcium-water reaction results from the metal being partly protected by insoluble white calcium hydroxide. In water solutions of acids, where this salt is soluble, calcium reacts vigorously. Calcium, with a density of 1.55 g/cm3, is the lightest of the alkaline earth metals; magnesium (specific gravity 1.74) and beryllium (1.84) are more dense, although lighter in atomic mass. From strontium onward, the alkali earth metals become more dense with increasing atomic mass. It has two allotropes.4doi:10.1007/BF02873196

Calcium has a higher electrical resistivity than copper or aluminium, yet weight-for-weight, due to its much lower density, it is a rather better conductor than either. However, its use in terrestrial applications is usually limited by its high reactivity with air. Calcium salts are colorless from any contribution of the calcium, and ionic solutions of calcium (Ca2+) are colorless as well. As with magnesium salts and other alkaline earth metal salts, calcium salts are often quite soluble in water. Notable exceptions include the hydroxide, the sulfate (unusual for sulfate salts), the carbonate and the phosphates. With the exception of the sulfate, even the insoluble ones listed are in general more soluble than its transition metal counterparts. When in solution, the calcium ion to the human taste varies remarkably, being reported as mildly salty, sour, "mineral like" or even "soothing." It is apparent that many animals can taste, or develop a taste, for calcium, and use this sense to detect the mineral in salt licks or other sources.5Tordoff, M. G. (2001). "Calcium: Taste, Intake, and Appetite". Physiological Reviews 81 (4): 156797. PMID11581497. .

In human nutrition, soluble calcium salts may be added to tart juices without much effect to the average palate. Calcium is the fifth-most-abundant element by mass in the human body, where it is a common cellular ionic messenger with many functions, and serves also as a structural element in bone. It is the relatively high-atomic-number calcium in the skeleton that causes bone to be radio-opaque. Of the human body's solid components after drying and burning of organics (as for example, after cremation), about a third of the total "mineral" mass remaining, is the approximately one kilogram of calcium that composes the average skeleton (the remainder being mostly phosphorus and oxygen). H and K lines Visible spectra of many stars, including the Sun, exhibit strong absorption lines of singly ionized calcium. Prominent among these are the H-line at 3968.5 and the K line at 3933.7 of singly ionized calcium, or CaII. For the Sun and stars with low temperatures, the prominence of the H and K lines can be an indication of strong magnetic activity in the chromosphere. Measurement of periodic variations of these active regions can also be used to deduce the rotation periods of these stars.6Staff (1995). "H-K Project". Mount Wilson Observatory. . Retrieved 2006-08-10.

Compounds Calcium, combined with phosphate to form hydroxylapatite, is the mineral portion of human and animal bones and teeth. The mineral portion of some corals can also be transformed into hydroxylapatite. Calcium hydroxide (slaked lime) is used in many chemical refinery processes and is made by heating limestone at high temperature (above 825C) and then carefully adding water to it. When lime is mixed with sand, it hardens into a mortar and is turned into plaster by carbon dioxide uptake. Mixed with other compounds, lime forms an important part of Portland cement. Calcium carbonate (CaCO3) is one of the common compounds of calcium. It is heated to form quicklime (CaO), which is then added to water (H2O). This forms another material known as slaked lime (Ca(OH)2), which is an inexpensive base material used throughout the chemical industry. Chalk, marble, and limestone are all forms of calcium carbonate. When water percolates through limestone or other soluble carbonate rocks, it partially dissolves the rock and causes cave formation and characteristic stalactites and stalagmites and also forms hard water. Other important calcium compounds are calcium nitrate, calcium sulfide, calcium chloride, calcium carbide, calcium cyanamide and calcium hypochlorite. A few calcium compounds in the oxidation state +1 have also been investigated recently.7

Nucleosynthesis Calcium-40 is created in extremely large and hot (over 2.5 109 K) stars, as part of the silicon-burning process in which alpha particles are added to silicon atoms. The process fuses an atom of argon and an atom of helium: 36Ar + 4He = 40Ca Isotopes Calcium has four stable isotopes (40Ca, 42Ca, 43Ca and 44Ca), plus two more isotopes (46Ca and 48Ca) that have such long half-lives that for all practical purposes they also can be considered stable. The 20% range in relative mass among naturally occurring calcium isotopes is greater than for any element except hydrogen and helium. Calcium also has a cosmogenic isotope, radioactive 41Ca, which has 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 metre 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. Ninety-seven percent of naturally occurring calcium is in the form of 40Ca. 40Ca 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 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 the most common calcium isotope (containing 10 helium nuclei) has been synthesized. Isotope fractionation As with the isotopes of other elements, a variety of processes fractionate, or alter the relative abundance of, calcium isotopes.8Russell, WA; Papanastassiou, DA; Tombrello, TA (1978). "Ca isotope fractionation on the earth and other solar system materials". Geochim Cosmochim Acta 42 (8): 107590. Bibcode1978GeCoA..42.1075R. doi:10.1016/0016-7037(78)90105-9.

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. Isotopically light calcium is preferentially incorporated into minerals, leaving the solution from which the mineral precipitated enriched in isotopically heavy calcium. At room temperature the magnitude of this fractionation is roughly 0.25 (0.025%) per atomic mass unit (AMU). Mass-dependent differences in calcium isotope composition conventionally are expressed 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.9Skulan, J; DePaolo, DJ (1999). "Calcium isotope fractionation between soft and mineralized tissues as a monitor of calcium use in vertebrates". Proc Natl Acad Sci USA 96 (24): 1370913. Bibcode1999PNAS...9613709S. doi:10.1073/pnas.96.24.13709. PMC24129. PMID10570137. .

Calcium isotope fractionation during mineral formation has led to several applications of calcium isotopes. In particular, the 1997 observation by Skulan and DePaolo10Skulan, J; DePaolo, DJ; Owens, TL (June 1997). "Biological control of calcium isotopic abundances in the global calcium cycle". Geochimica et Cosmochimica Acta 61 (12): 250510. Bibcode1997GeCoA..61.2505S. doi:10.1016/S0016-7037(97)00047-1.

that calcium minerals are isotopically lighter than the solutions from which the minerals precipitate is the basis of analogous applications in medicine and in paleooceanography. In animals with skeletons mineralized 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, soft tissue 44Ca/40Ca rises. Soft tissue 44Ca/40Ca falls when bone resorption exceeds bone formation. Because of this relationship, calcium isotopic measurements of urine or blood may be useful in the early detection of metabolic bone diseases like osteoporosis.11Skulan, J; Bullen, T; Anbar, AD; Puzas, JE; Shackelford, L; Leblanc, A; Smith, SM (2007). "Natural calcium isotopic composition of urine as a marker of bone mineral balance". Clinical Chemistry 653 (6): 11551158. doi:10.1373/clinchem.2006.080143. PMID17463176. .

A similar system exists in the ocean, where seawater 44Ca/40Ca tends to rise when the rate of removal of Ca2+ from seawater by mineral precipitation exceeds the input of new calcium into the ocean, and fall when calcium input exceeds mineral precipitation. It follows that rising 44Ca/40Ca corresponds to falling seawater Ca2+ concentration, and falling 44Ca/40Ca corresponds to rising seawater Ca2+ concentration. In 1997 Skulan and DePaolo presented the first evidence of change in seawater 44Ca/40Ca over geologic time, along with a theoretical explanation of these changes. More recent papers have confirmed this observation, demonstrating that seawater Ca2+ concentration is not constant, and that the ocean probably never is in steady state with respect to its calcium input and output.12Fantle, M; DePaolo, D (2007). "Ca isotopes in carbonate sediment and pore fluid from ODP Site 807A: The Ca2+(aq)calcite equilibrium fractionation factor and calcite recrystallization rates in Pleistocene sediments". Geochim Cosmochim Acta 71 (10): 25242546. Bibcode2007GeCoA..71.2524F. doi:10.1016/j.gca.2007.03.006.

13Griffith, Elizabeth M.; Paytan, Adina; Caldeira, Ken; Bullen, Thomas; Thomas, Ellen (2008). "A Dynamic marine calcium cycle during the past 28 million years". Science 322 (12): 16711674. Bibcode2008Sci...322.1671G. doi:10.1126/science.1163614. PMID19074345.

This has important climatological implications, as the marine calcium cycle is closely tied to the carbon cycle (see below). Geochemical cycling Calcium provides an important link between tectonics, climate and the carbon cycle. In the simplest terms, uplift of mountains exposes Ca-bearing rocks to chemical weathering and releases Ca2+ into surface water. This Ca2+ eventually is transported to the ocean where it reacts with dissolved CO2 to form limestone. Some of this limestone settles to the sea floor where it is incorporated into new rocks. Dissolved CO2, along with carbonate and bicarbonate ions, are referred to as dissolved inorganic carbon (DIC). Travertine terraces Pamukkale, Turkey

The actual reaction is more complicated and involves the bicarbonate ion (HCO3-) that forms when CO2 reacts with water at seawater pH: Ca2+ + 2HCO CaCO3 (limestone) + CO2 + H2O

Note that at ocean pH most of the CO2 produced in this reaction is immediately converted back into HCO. The reaction results in a net transport of one molecule of CO2 from the ocean/atmosphere into the lithosphere.14Zeebe (2006). "Marine carbonate chemistry". National Council for Science and the Environment. . Retrieved 2010-03-13.

The result is that each Ca2+ ion released by chemical weathering ultimately removes one CO2 molecule from the surficial system (atmosphere, ocean, soils 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 CO2 from the ocean and atmosphere, exerting a strong long-term effect on climate.15Berner, Robert (2003). "The long-term carbon cycle, fossil fuels and atmospheric composition". Nature 426 (6964): 323326. doi:10.1038/nature02131. PMID14628061.

Analogous cycles involving magnesium, and to a much smaller extent strontium and barium, have the same effect. As the weathering of limestone (CaCO3) liberates equimolar amounts of Ca2+ and CO2, it has no net effect on the CO2 content of the atmosphere and ocean. The weathering of silicate rocks like granite, on the other hand, is a net CO2 sink because it produces abundant Ca2+ but very little CO2. History Lime as building material was used since prehistoric times going as far back as 7000 to 14000BC.16Miller, M. Michael. "Commodity report:Lime". United States Geological Survey. . Retrieved 2012-03-06.

The first dated lime kiln dates back to 2500BC and was found in Khafajah mesopotamia.17Williams, Richard (2004). Lime Kilns and Lime Burning. p.4. ISBN978-0-7478-0596-0. .

18Oates, J. A. H (2008-07-01). Lime and Limestone: Chemistry and Technology, Production and Uses. ISBN978-3-527-61201-7. .

Calcium (from Latin calx, genitive calcis, meaning "lime")19 calx. Charlton T. Lewis and Charles Short. A Latin Dictionary on Perseus Project.

was known as early as the first century when the Ancient Romans prepared lime as calcium oxide. Literature dating back to 975 AD notes that plaster of paris (calcium sulfate), is useful for setting broken bones. It was not isolated until 1808 in England when Sir Humphry Davy electrolyzed a mixture of lime and mercuric oxide.20Davy H (1808). "Electro-chemical researches on the decomposition of the earths; with observations on the metals obtained from the alkaline earths, and on the amalgam procured from ammonia". Philosophical Transactions of the Royal Society of London 98: 333370. Bibcode1808RSPT...98..333D. doi:10.1098/rstl.1808.0023. .

Davy was trying to isolate calcium; when he heard that Swedish chemist Jns Jakob Berzelius and Pontin prepared calcium amalgam by electrolyzing lime in mercury, he tried it himself. He worked with electrolysis throughout his life and also discovered/isolated sodium, potassium, magnesium, boron and barium. Calcium metal was not available in large scale until the beginning of the 20th century.Occurrence Calcium is not naturally found in its elemental state. Calcium occurs most commonly in sedimentary rocks in the minerals calcite, dolomite and gypsum. It also occurs in igneous and metamorphic rocks chiefly in the silicate minerals: plagioclases, amphiboles, pyroxenes and garnets. Applications Calcium is used21Lide, D. R., ed. (2005). CRC Handbook of Chemistry and Physics (86th ed.). Boca Raton (FL): CRC Press. ISBN0-8493-0486-5.

as a reducing agent in the extraction of other metals, such as uranium, zirconium, and thorium.

as a deoxidizer, desulfurizer, or decarbonizer for various ferrous and nonferrous alloys.

as an alloying agent used in the production of aluminium, beryllium, copper, lead, and magnesium alloys.

in the making of cements and mortars to be used in construction.

in the making of cheese, where calcium ions influence the activity of rennin in bringing about the coagulation of milk.

Calcium compoundsCalcium carbonate (CaCO3) is used in manufacturing cement and mortar, lime, limestone (usually used in the steel industry) and aids in production in the glass industry. It also has chemical and optical uses as mineral specimens in toothpastes, for example.

Calcium hydroxide solution (Ca(OH)2) (also known as limewater) is used to detect the presence of carbon dioxide by being bubbled through a solution. It turns cloudy where CO2 is present.

Calcium arsenate (Ca3(AsO4)2) is used in insecticides.

Calcium carbide (CaC2) is used to make acetylene gas (for use in acetylene torches for welding) and in the manufacturing of plastics.

Calcium chloride (CaCl2) is used in ice removal and dust control on dirt roads, in conditioner for concrete, as an additive in canned tomatoes, and to provide body for automobile tires.

Calcium cyclamate (Ca(C6H11NHSO3)2) was used as a sweetening agent but is no longer permitted for use because of suspected cancer-causing properties.

Calcium gluconate (Ca(C6H11O7)2) is used as a food additive and in vitamin pills.

Calcium hypochlorite (Ca(OCl)2) is used as a swimming pool disinfectant, as a bleaching agent, as an ingredient in deodorant, and in algaecide and fungicide.

Calcium permanganate (Ca(MnO4)2) is used in liquid rocket propellant, textile production, as a water sterilizing agent and in dental procedures.

Calcium phosphate (Ca3(PO4)2) is used as a supplement for animal feed, fertilizer, in commercial production for dough and yeast products, in the manufacture of glass, and in dental products.

Calcium phosphide (Ca3P2) is used in fireworks, rodenticide, torpedoes and flares.

Calcium stearate (Ca(C18H35O2)2) is used in the manufacture of wax crayons, cements, certain kinds of plastics and cosmetics, as a food additive, in the production of water resistant materials and in the production of paints.

Calcium sulfate (CaSO42H2O) is used as common blackboard chalk, as well as, in its hemihydrate form better known as Plaster of Paris.

Calcium tungstate (CaWO4) is used in luminous paints, fluorescent lights and in X-ray studies.

Hydroxylapatite (Ca5(PO4)3(OH), but is usually written Ca10(PO4)6(OH)2) makes up seventy percent of bone. Also carbonated-calcium deficient hydroxylapatite is the main mineral of which dental enamel and dentin are comprised.

Nutrition Recommended adequate intake by the IOM for calcium:22 "Dietary Supplement Fact Sheet: Calcium". . Retrieved 8 March 2011.

23 "Dietary Reference Intakes for Calcium and Vitamin D". November 2010. .

Age Calcium (mg/day)

06 months 200

712 months 260

13 years 700

48 years 1000

918 years 1300

1950 years 1000

5170 years (male) 1000

5170 years (female) 1200

71+ years 1200

Calcium is an important component of a healthy diet and a mineral necessary for life. The National Osteoporosis Foundation says, "Calcium plays an important role in building stronger, denser bones early in life and keeping bones strong and healthy later in life." Approximately 99 percent of the body's calcium is stored in the bones and teeth.24 "Dietary Supplement Fact Sheet: Calcium". Office of Dietary Supplements, NIH. . Retrieved 31 March 2011.

The rest of the calcium in the body has other important uses, such as some exocytosis, especially neurotransmitter release, and muscle contraction. In the electrical conduction system of the heart, calcium replaces sodium as the mineral that depolarizes the cell, proliferating the action potential. In cardiac muscle, sodium influx commences an action potential, but during potassium efflux, the cardiac myocyte experiences calcium influx, prolonging the action potential and creating a plateau phase of dynamic equilibrium. Long-term calcium deficiency can lead to rickets and poor blood clotting and in case of a menopausal woman, it can lead to osteoporosis, in which the bone deteriorates and there is an increased risk of fractures. While a lifelong deficit can affect bone and tooth formation, over-retention can cause hypercalcemia (elevated levels of calcium in the blood), impaired kidney function and decreased absorption of other minerals.25Standing Committee on the Scientific Evaluation of Dietary Reference Intakes, Food and Nutrition Board, Institute of Medicine (1997). Dietary Reference Intakes for Calcium, Phosphorus, Magnesium, Vitamin D and fluoride. Washington DC: The National Academies Press. ISBN0-309-06403-1. .

26Committee to Review Dietary Reference Intakes for Vitamin D and Calcium; Institute of Medicine (2011). A. Catharine Ross, Christine L. Taylor, Ann L. Yaktine, Heather B. Del Valle. ed. Dietary Reference Intakes for Calcium and Vitamin D. ISBN978-0-309-16394-1. .

Several sources suggest a correlation between high calcium intake (2000mg per day, or twice the U.S. recommended daily allowance, equivalent to six or more glasses of milk per day) and prostate cancer.27Giovannucci E, Rimm EB, Wolk A, et al. (February 1998). "Calcium and fructose intake in relation to risk of prostate cancer". Cancer Research 58 (3): 4427. PMID9458087. .

High calcium intakes or high calcium absorption were previously thought to contribute to the development of kidney stones. However, a high calcium intake has been associated with a lower risk for kidney stones in more recent research.28Curhan, GC; Willett, WC; Rimm, EB; Stampfer, MJ (1993). "A prospective study of dietary calcium and other nutrients and the risk of symptomatic kidney stones" (PDF). The New England Journal of Medicine 328 (12): 8338. doi:10.1056/NEJM199303253281203. PMID8441427. .

29Bihl G, Meyers A. (2001). "Recurrent renal stone disease-advances in pathogenesis and clinical management". Lancet 358 (9282): 651656. doi:10.1016/S0140-6736(01)05782-8. PMID11530173.

30Hall WD, Pettinger M, Oberman A (2001). "Risk factors for kidney stones in older women in the Southern United States". Am J Med Sci 322 (1): 1218. doi:10.1097/00000441-200107000-00003. PMID11465241.

Vitamin D is needed to absorb calcium. Dairy products, such as milk and cheese, are a well-known source of calcium. Some individuals are allergic to dairy products and even more people, in particular those of non Indo-European descent, are lactose-intolerant, leaving them unable to consume non-fermented dairy products in quantities larger than about half a liter per serving. Others, such as vegans, avoid dairy products for ethical and health reasons. Many good sources of calcium exist, including seaweeds such as kelp, wakame and hijiki; nuts and seeds like almonds, hazelnuts, sesame, pistachio; blackstrap molasses; beans; figs; quinoa; okra; rutabaga; broccoli; dandelion leaves; kale; and fortified products such as orange juice and soy milk. Numerous vegetables, notably spinach, chard and rhubarb have a high calcium content, but they may also contain varying amounts of oxalic acid that binds calcium and reduces its absorption. The same problem may to a degree affect the absorption of calcium from amaranth, collard greens, chicory greens. This process may also be related to the generation of calcium oxalate. An overlooked source of calcium is eggshell, which can be ground into a powder and mixed into food or a glass of water.31Schaafsma, Anne and Beelen, Gerard M (1999). "Eggshell powder, a comparable or better source of calcium than purified calcium carbonate: piglet studies". Journal of the Science of Food and Agriculture 79 (12): 15961600. doi:10.1002/(SICI)1097-0010(199909)79:123.0.CO;2-A.

32Schaafsma A, van Doormaal JJ, Muskiet FA, Hofstede GJ, Pakan I, van der Veer E (2002). "Positive effects of a chicken eggshell powder-enriched vitamin-mineral supplement on femoral neck bone mineral density in healthy late post-menopausal Dutch women". Br. J. Nutr. 87 (3): 26775. doi:10.1079/BJNBJN2001515. PMID12064336.

33Rovensk J, Stanckov M, Masaryk P, Svk K, Istok R (2003). "Eggshell calcium in the prevention and treatment of osteoporosis". Int J Clin Pharmacol Res 23 (23): 8392. PMID15018022.

The calcium content of most foods can be found in the USDA National Nutrient Database.34 "USDA National Nutrient Database". .

Dietary calcium supplements500 milligram calcium supplements made from calcium carbonate

Calcium supplements are used to prevent and to treat calcium deficiencies. Most experts recommend that supplements be taken with food and that no more than 600mg should be taken at a time because the percent of calcium absorbed decreases as the amount of calcium in the supplement increases.35

It is recommended to spread doses throughout the day. Recommended daily calcium intake for adults ranges from 1000 to 1500mg. It is recommended to take supplements with food to aid in absorption. Vitamin D is added to some calcium supplements. Proper vitamin D status is important because vitamin D is converted to a hormone in the body, which then induces the synthesis of intestinal proteins responsible for calcium absorpti