| Related sites for http://en.wikipedia.org/wiki/Potassium |
| Tellurium Data tables and historic information. | | Visual_Elements__Tellurium General and physical information, source, uses, key isotopes, and ionisation energies. | | WebElements__Tellurium Extensive information on history, uses, occurrence, compounds, and properties. | | Wikipedia__Tellurium Properties of the element, including its history, applications, and characteristics. | | Radiochemistry_of_Tellurium Full text of the monograph by G. W. Leddicotte (Oak Ridge National Laboratory, Oak Ridge, Tennessee). [PDF] (July, 1961) | | ChemicalElements_com__Cobalt Basic information, atomic structure, and isotopes. | | Cobalt Data tables and historic information. | | EnvironmentalChemistry_com__Cobalt Atomic structure, chemical and physical properties, and table of nuclides. | | It\'s_Elemental__Cobalt Basic physical and historical information. | | LANL__Cobalt History, sources, properties, and uses. | | Lenntech__Cobalt Physical data, chemical properties, health and environmental effects. | | USGS_Minerals_Information__Cobalt Statistics and information on the worldwide supply, demand, and flow of the element (PDF format). | | Visual_Elements__Cobalt Image, general and physical information, source, uses, biological role, key isotopes, and ionisation energies. | | WebElements__Cobalt Extensive information on history, uses, occurrence, compounds, and properties of the element. | | Wikipedia__Cobalt Properties of the element, including its history, applications, and characteristics. | | Radiochemistry_of_Cobalt Full text of the monograph by L.C. Bate (Oak Ridge National Laboratory, Oak Ridge, Tennessee). [PDF] (September, 1961) | | ChemicalElements_com__Nickel Basic information, atomic structure, and isotopes. | | EnvironmentalChemistry_com__Nickel Atomic structure, chemical and physical properties, and table of nuclides. | | It\'s_Elemental__Nickel Basic physical and historical information. | | LANL__Nickel Sources, properties, and uses. | | Lenntech__Nickel Physical data, chemical properties, health and environmental effects. | | Nickel Data tables and historic information. | | USGS_Minerals_Information__Nickel Statistics and information on the worldwide supply, demand, and flow of the element (PDF format). | | Visual_Elements__Nickel Image, general and physical information, source, uses, key isotopes, and ionisation energies. | | WebElements__Nickel Extensive information on history, uses, occurrence, compounds, and properties of the element. | | Wikipedia__Nickel Properties of the element, including its history, applications, and characteristics. | | Radiochemistry_of_Nickel Full text of the monograph by L. J. Kirby (Hanford Atomic Products Operation, Richland, Washington). [PDF] (November, 1961) | | ChemicalElements_com__Platinum Basic information, atomic structure, and isotopes. | | EnvironmentalChemistry_com__Platinum Atomic structure, chemical and physical properties, and table of nuclides. | | It\'s_Elemental__Platinum Basic physical and historical information. | | LANL__Platinum Sources, properties, and uses. | | Lenntech__Platinum Physical data, chemical properties, health and environmental effects. | | Platinum Data tables and historic information. | | Visual_Elements__Platinum Image, general and physical information, source, uses, key isotopes, and ionisation energies. | | WebElements__Platinum Extensive information on history, uses, occurrence, compounds, and properties. | | Wikipedia__Platinum Properties of the element, including its history, applications, and characteristics. | | Radiochemistry_of_Platinum Full text of the monograph by G. W. Leddicotte (Oak Ridge National Laboratory, Oak Ridge, Tennessee). [PDF] (October, 1961) | | Cesium Data tables and historic information. | | ChemicalElements_com__Cesium Basic information, atomic structure, and isotopes table. | | EnvironmentalChemistry_com__Cesium Atomic structure, chemical and physical properties, and table of nuclides. |
|
Potassium - Wikipedia, the free encyclopedia /**/ /**/ if (wgNotice != '') document.writeln(wgNotice); Potassium From Wikipedia, the free encyclopedia Jump to: navigation, search 19argon ← potassium → calciumNa↑K↓Rb Periodic table - Extended periodic tableGeneralName, symbol, numberpotassium, K, 19Element categoryalkali metalsGroup, period, block1, 4, sAppearancesilvery white Standard atomic weight39.0983(1) g·mol−1Electron configuration[Ar] 4s1Electrons per shell2, 8, 8, 1Physical propertiesPhasesolidDensity (near r.t.)0.89 g·cm−3Liquid density at m.p.0.828 g·cm−3Melting point336.53 K(63.38 °C, 146.08 °F)Boiling point1032 K(759 °C, 1398 °F)Triple point336.35 K (63°C), kPaAtomic propertiesCrystal structurecubic body centeredOxidation states1(strongly basic oxide)Electronegativity0.82 (Pauling scale)Ionization energies(more)1st: 418.8 kJ·mol−12nd: 3052 kJ·mol−13rd: 4420 kJ·mol−1Atomic radius220 pmAtomic radius (calc.)243 pmCovalent radius196 pmVan der Waals radius275 pmMiscellaneousMagnetic orderingparamagneticThermal conductivity(300 K) 102.5 W·m−1·K−1Thermal expansion(25 °C) 83.3 µm·m−1·K−1Speed of sound (thin rod)(20 °C) 2000 m/sYoung's modulus3.53 GPaShear modulus1.3 GPaBulk modulus3.1 GPaMohs hardness0.4Brinell hardness0.363 MPaCAS registry number7440-09-7Selected isotopesMain article: Isotopes of potassiumisoNAhalf-lifeDMDE (MeV)DP39K93.26%39K is stable with 20 neutrons40K0.012%1.248(3)×109 yβ-1.31140Caε1.50540Arβ+1.50540Ar41K6.73%41K is stable with 22 neutronsReferencesThis box: view • talk • editPotassium (pronounced /pəˈtæsiəm/) is a chemical element. It has the symbol K (Latin: kalium, from Arabic: القَلْيَه al qalya “plant ashes”, cf. Alkali from the same root), atomic number 19, and atomic mass 39.0983. The name "potassium" comes from the word "potash", as potassium was first isolated from potash. Potassium is a soft silvery-white metallic alkali metal that occurs naturally bound to other elements in seawater and many minerals. It oxidizes rapidly in air and is very reactive with water, generating sufficient heat to ignite the evolved hydrogen. In many respects, potassium and sodium are chemically similar, although they have very different functions in organisms in general, and in animal cells in particular. It is found in bananas, potatoes and especially avocados as part of a regular diet.Contents1 Occurence2 Production3 Isotopes4 Properties4.1 Physical4.2 Chemical5 Potassium cations in the body5.1 Biochemical function5.2 Membrane polarization5.3 Filtration and excretion5.4 Potassium in the diet6 Applications6.1 Biological applications6.2 Food applications6.3 Industrial applications7 History8 Precautions9 References10 See also11 External links//[edit] Occurence  Potassium in feldsparPotassium is a metal that is never found free, as it reacts violently with the abundant water in nature.[1] As various compounds, potassium makes up about 1.5% of the weight of the Earth's crust and is the seventh most abundant element.[1] As it is very electropositive, potassium metal is difficult to obtain from its minerals.See also: Category:Potassium minerals[edit] ProductionPure potassium metal can be isolated by electrolysis of its hydroxide in a process that has changed little since Davy.[1] Thermal methods also are employed in potassium production, using potassium chloride Humphry Davy extracted this metal in 1807 along with sodium.[2]Potassium salts such as carnallite, langbeinite, polyhalite, and sylvite form extensive deposits in ancient lake and seabeds, making extraction of potassium salts in these environments commercially viable. The principal source of potassium, potash, is mined in Saskatchewan, California, Germany, New Mexico, Utah, and in other places around the world. It is also found abundantly in the Dead Sea. Three thousand feet below the surface of Saskatchewan are large deposits of potash which are important sources of this element and its salts, with several large mines in operation since the 1960s. Saskatchewan pioneered the use of freezing of wet sands (the Blairmore formation) in order to drive mine shafts through them. The main mining company is the Potash Corporation of Saskatchewan. The oceans are another source of potassium, but the quantity present in a given volume of seawater is relatively low compared with sodium.[3][4][edit] IsotopesMain article: isotopes of potassiumThere are 24 known isotopes of potassium. Three isotopes occur naturally: 39K (93.3%), 40K (0.0117%) and 41K (6.7%). Naturally occurring 40K decays to stable 40Ar (11.2%) by electron capture and by positron emission, and decays to stable 40Ca (88.8%) by beta decay; 40K has a half-life of 1.250×109 years. 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, plutonic/high grade metamorphic hornblende, and volcanic feldspar; whole rock samples from volcanic flows and shallow instrusives can also be dated if they are unaltered.Outside of dating, potassium isotopes have been used extensively as tracers in studies of weathering. They have also been used for nutrient cycling studies because potassium is a macronutrient required for life.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 70 kg mass, about 4,400 nuclei of 40K decay per second.[5]The activity of natural potassium is 31 Bq/g.[edit] Properties[edit] Physical The flame-test colour for potassiumPotassium is the second least dense metal; only lithium is less dense. It is a soft, low-melting solid that can easily be cut with a knife. Freshly cut potassium is silvery in appearance, but in air it begins to tarnish toward grey immediately.[1]In a flame test, potassium and its compounds emit a pale violet color, which may be masked by the strong yellow emission of sodium if it is also present. Cobalt glass can be used to filter out the yellow sodium color.[6] Potassium concentration in solution is commonly determined by flame photometry, atomic absorption spectrophotometry, inductively coupled plasma, or ion selective electrodes.[edit] ChemicalPotassium must be protected from air for storage to prevent disintegration of the metal from oxide and hydroxide corrosion. Often samples are maintained under a reducing medium such as kerosene.Like the other alkali metals, potassium reacts violently with water, producing hydrogen. The reaction is notably more violent than that of lithium or sodium with water, and is sufficiently exothermic that the evolved hydrogen gas ignites.2K(s) + 2H2O(l) → H2(g) + 2KOH(aq)Because potassium reacts quickly with even traces of water, and its reaction products are nonvolatile, it is sometimes used alone, or as NaK (an alloy with sodium which is liquid at room temperature) to dry solvents prior to distillation. In this role, it serves as a potent desiccant.Potassium hydroxide reacts strongly with carbon dioxide to produce potassium carbonate, and is used to remove traces of CO2 from air.Potassium compounds generally have excellent water solubility, due to the high hydration energy of the K+ ion. The potassium ion is colorless in water.Methods of separating potassium by precipitation, sometimes used for gravimetric analysis, include the use of sodium tetraphenylborate, hexachloroplatinic acid, and sodium cobaltinitrite[edit] Potassium cations in the body[edit] Biochemical functionMain article: Action potentialPotassium cations are important in neuron (brain and nerve) function, and in influencing osmotic balance between cells and the interstitial fluid.[7].Potassium may be detected by taste because it triggers three of the five types of taste sensations, according to concentration. Dilute solutions of potassium ion 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 content solutions makes high-dose potassium supplementation by liquid drinks a palatability challenge.[citation needed][edit] Membrane polarizationPotassium is also important in allowing 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. [8]A shortage of potassium in body fluids may cause a potentially fatal condition known as hypokalemia, typically resulting from diarrhea, increased diuresis and vomiting. Deficiency symptoms include muscle weakness, paralytic ileus, ECG abnormalities, decreased reflex response and in severe cases respiratory paralysis, alkalosis and cardiac arrhythmia.[edit] Filtration and excretionPotassium is an essential mineral micronutrient 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 milliequivalents per liter (3345 milligrams) and potassium makes up most of the cell fluid cations at about 150 milliequivalents per liter (4800 milligrams). Plasma is filtered through the glomerulus of the kidneys in enormous amounts, about 180 liters per day.[9] Thus 602,000 milligrams of sodium and 33,000 milligrams of potassium are filtered each day. All but the 1000-10,000 milligrams of sodium and the 1000-4000 milligrams 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 milliequivalents (about 190 milligrams) per liter.[10] Sodium pumps must always operate to conserve sodium. Potassium must sometimes be conserved also, but since 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 passively[11][12] in counter flow to sodium in response to an apparent (but not actual) Donnan equilibrium,[13] 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.[14] At that point, it usually has about the same potassium concentration as plasma. If potassium were removed from the diet, there would remain a minimum obligatory kidney excretion of about 200 mg per day when the serum declines to 3.0-3.5 milliequivalents per liter in about one week,[15] and can never be cut off completely. Because it cannot be cut off completely, death will result when the whole body potassium declines to the vicinity of one-half full capacity. At the end of the processing, potassium is secreted one more time if the serum levels are too high. Reference ranges for blood tests, showing blood content of potassium in blue in right part of the spectrum.The potassium moves passively through pores in the cell wall. When ions move through pumps there is a gate in the pumps on either side of the cell wall and only one gate can be open at once. As a result 100 ions are forced through per second. Pores have only one gate and there one kind of ion only can stream through at 10 million to 100 million ions per second.[16] The pores require calcium in order to open[17] although it is thought that the calcium works in reverse by blocking at least one of the pores.[18] Carbonyl groups inside the pore on the amino acids mimics the water hydration that takes place in water solution[19] by the nature of the electrostatic charges on four carbonyl groups inside the pore.[20][edit] Potassium in the dietAdequate intake can generally be guaranteed by eating a variety of foods containing potassium and deficiency is rare in healthy individuals eating a balanced diet. Foods with high sources of potassium include orange juice, potatoes, bananas, avocados, tomatoes, broccoli, soybeans, brown rice, garlic and apricots, although it is also common in most fruits, vegetables and meats [21]. Diets high in potassium can reduce the risk of hypertension and a potassium deficiency combined with an inadequate thiamine intake has produced heart disease in rats.[22] The 2004 guidelines of the Institute of Medicine specify a DRI of 4,000mg of potassium, though most Americans consume only half that amount per day.[23] Similarly, in the European Union, particularly in Germany and Italy, insufficient potassium intake is somewhat common.[24]Supplements of potassium in medicine are most widely used in conjunction with loop diuretics and thiazides, classes of diuretics which rid the body of sodium and water, but have the side effect of also causing potassium loss in urine. A variety of medical supplements are available. If potassium supplements are used, such as sodium free baking powder and sodium free table salt, inadequate thiamine can cause beriberi.[25][26][citation needed]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, since the kidneys control potassium excretion, and buildup of blood concentrations of potassium may trigger fatal cardiac arrhythmia. Acute hyperkalemia can be reduced through eating baking soda,[27] or glucose,[28][29] hyperventilation[30] and perspiration.[31][edit] ApplicationsAbout 93% of the world potassium production was consumed by the fertilizer industry.[4][edit] Biological applications Potassium and Magnesium sulfate fertilizerPotassium ion is an essential component of plant nutrition and is found in most soil types. Its primary use in agriculture, horticulture and hydroponic culture is as a fertilizer as the chloride (KCl), sulfate (K2SO4) or nitrate (KNO3).In animal cells, potassium ions are vital to keeping cells alive (see Na-K pump).[edit] Food applicationsPotassium ion 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, bananas, and many other good dietary sources of potassium, ranked according to potassium content per measure shown.[32]Potassium sodium tartrate, or Rochelle salt (KNaC4H4O6) is the main constituent of baking powder. Potassium bromate (KBrO3) is a strong oxidiser, used as a flour improver (E924) to improve dough strength and rise height.The sulfite compound, 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.Non-dietary uses of potassium chloride include its use to stop the heart, e.g. in cardiac surgery and in a solution used in executions by lethal injection.[edit] Industrial applicationsPotassium vapor is used in several types of magnetometers. An alloy of sodium and potassium, NaK (usually pronounced "nack"[citation needed]), that is liquid at room temperature, is used as a heat-transfer medium. It can also be used as a desiccant for producing dry and air-free solvents.Potassium metal reacts vigorously with all of the halogens to form the corresponding potassium halides, which are white, water-soluble salts with cubic crystal morphology. Potassium bromide (KBr), potassium iodide (KI) and potassium chloride (KCl) are used in photographic emulsion to make the corresponding photosensitive silver halides.Potassium hydroxide KOH is a strong base, used in industry to neutralize strong and weak acids and thereby finding uses in pH control and in the manufacture of potassium salts. Potassium hydroxide is also used to saponify fats and oils and in hydrolysis reactions, for example of esters and in industrial cleaners.Potassium nitrate KNO3 or saltpeter is obtained from natural sources such as guano and evaporites or manufactured by the Haber process and is the oxidant in gunpowder (black powder) and an important agricultural fertilizer. Potassium cyanide KCN is used industrially to dissolve copper and precious metals particularly silver and gold by forming complexes; applications include gold mining, electroplating and electroforming of these metals. It is also used in organic synthesis to make nitriles. Potassium carbonate K2CO3, also known as potash, is used in the manufacture of glass and soap and as a mild desiccant.Potassium chromate (K2CrO4) is used in dyes and stains (bright yellowish-red colour), in explosives and fireworks, in safety matches, in the tanning of leather and in fly paper. Potassium fluorosilicate (K2SiF6) is used in specialized glasses, ceramics, and enamels. Potassium sodium tartrate, or Rochelle salt (KNaC4H4O6) is used in the silvering of mirrors.The superoxide KO2 is an orange-colored solid used as a portable source of oxygen and as a carbon dioxide absorber. It is useful in portable respiration systems. It is widely used in submarines and spacecraft as it takes less volume than O2(g).4KO2 + 2CO2 --- 2K2CO3 + O2Potassium chlorate KClO3 is a strong oxidant, used in percussion caps and safety matches and in agriculture as a weedkiller. Glass may be treated with molten potassium nitrate KNO3 to make toughened glass, which is much stronger than regular glass.[edit] History Please help improve this section by expanding it. Further information might be found on the talk page. (June 2008)Potassium was discovered in 1807 in England by Sir Humphry Davy, who derived it from caustic potash (KOH). Before the 18th century, no distinction was made between potassium and sodium. Potassium was the first metal that was isolated by electrolysis.[33]Potassium was not known in Roman times, and its names are not Classical Latin but rather neo-Latin.The name kalium was taken from the word "alkali", which came from Arabic al qalīy = "the calcined ashes".The name potassium was made from the word "potash", which is English, and originally meant an alkali extracted in a pot from the ash of burnt wood or tree leaves.[edit] PrecautionsPlease help improve this section by expanding itwith:additional citations. Further information might be found on the talk page. (February 2008) Peroxides (Yellow) and Ozonides (Red) on surface of potassium metal.Potassium reacts very violently with water producing hydrogen gas which then usually catches fire. Potassium is usually kept under a mineral oil such as kerosene to stop the metal reacting with water vapour present in the air. Unlike lithium and sodium, however, potassium should not be stored under oil indefinitely. If stored longer than 6 months to a year, dangerous shock-sensitive peroxides can form on the metal and under the lid of the container, which can detonate upon opening. It is recommended that potassium, rubidium or caesium not be stored for longer than three months unless stored in an inert (oxygen free) atmosphere, or under vacuum.[34]As potassium reacts with water to produce highly flammable hydrogen gas, a potassium fire is only exacerbated by the addition of water, and only a few dry chemicals are effective for putting out such a fire (see the precaution section in sodium).Potassium also produces potassium hydroxide (KOH) in the reaction with water. Potassium hydroxide is a strong alkali and so is a caustic hazard, causing burns.Due to the highly reactive nature of potassium, it should be handled with great care, with full skin and eye protection being used and preferably a explosive resistant barrier between the user and the source of the potassium.[edit] References^ a b c d Mark Winter. "Potassium: Key Information". Webelements.^ Davy, Humphry (1808). "On some new Phenomena of Chemical Changes produced by Electricity, particularly the Decomposition of the fixed Alkalies, and the Exhibition of the new Substances, which constitute their Bases". Philosophical Transactions of the Royal Society of London 1: 1–45, http://books.google.com/books?id=Kg9GAAAAMAAJ. ^ Ober, Joyce A.. "Mineral Commodity Summaries 2008:Potash". United States Geological Survey. Retrieved on 2008-11-20.^ a b Ober, Joyce A.. "Mineral Yearbook 2006:Potash". United States Geological Survey. Retrieved on 2008-11-20.^ "background radiation - potassium-40 - γ radiation".^ Anne Marie Helmenstine. "Qualitative Analysis - Flame Tests". About.com.^ Campbell, Neil (1987). Biology. pp.795. ISBN 0-8053-1840-2. ^ Lockless SW, Zhou M, MacKinnon R.. "Structural and thermodynamic properties of selective ion binding in a K+ channel". Laboratory of Molecular Neurobiology and Biophysic, Rockefeller University. Retrieved on 2008-03-08.^ Potts, W.T.W.; Parry, G. (1964). Osmotic and ionic regulation in animals, Pergamon Press. ^ Lans HS, Stein IF, Meyer KA (1952). "The relation of serum potassium to erythrocyte potassium in normal subjects and patients with potassium deficiency". Am. J. Med. Sci. 223 (1): 65–74. doi:10.1097/00000441-195201000-00011. PMID 14902792. ^ Bennett CM, Brenner BM, Berliner RW (1968). "Micropuncture study of nephron function in the rhesus monkey". J Clin Invest 47 (1): 203–216. PMID 16695942. ^ Solomon AK (1962). "Pumps in the living cell". Sci. Am. 207: 100–8. PMID 13914986. ^ Kernan, Roderick P. (1980). Cell potassium (Transport in the life sciences). New York: Wiley. ISBN 0471048062. ; p. 40 & 48.^ Wright FS (1977). "Sites and mechanisms of potassium transport along the renal tubule". Kidney Int. 11 (6): 415–32. doi:10.1038/ki.1977.60. PMID 875263. ^ Squires RD, Huth EJ (1959). "Experimental potassium depletion in normal human subjects. I. Relation of ionic intakes to the renal conservation of potassium". J. Clin. Invest. 38 (7): 1134–48. doi:10.1172/JCI103890. PMID 13664789. ^ Gadsby DC (2004). "Ion transport: spot the difference". Nature 427 (6977): 795–7. doi:10.1038/427795a. PMID 14985745. ; for a diagram of the potassium pores are viewed, see Miller C (2001). "See potassium run". Nature 414 (6859): 23–4. doi:10.1038/35102126. PMID 11689922. ^ Jiang Y, Lee A, Chen J, Cadene M, Chait BT, MacKinnon R (2002). "Crystal structure and mechanism of a calcium-gated potassium channel". Nature 417 (6888): 515–22. doi:10.1038/417515a. PMID 12037559. ^ Shi N, Ye S, Alam A, Chen L, Jiang Y (2006). "Atomic structure of a Na+- and K+-conducting channel". Nature 440 (7083): 570–4. doi:10.1038/nature04508. PMID 16467789. ; includes a detailed picture of atoms in the pump.^ Zhou Y, Morais-Cabral JH, 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): 43–8. doi:10.1038/35102009. PMID 11689936. ^ Noskov SY, Bernèche S, Roux B (2004). "Control of ion selectivity in potassium channels by electrostatic and dynamic properties of carbonyl ligands". Nature 431 (7010): 830–4. doi:10.1038/nature02943. PMID 15483608. ^ "Potassium Content of Food and Drink". Retrieved on 2008-09-18.^ Folis, R.H. (1942). "Myocardial Necrosis in Rats on a Potassium Low Diet Prevented by Thiamine Deficiency". Bull. Johns-Hopkins Hospital 71: 235. ^ Grim CE, Luft FC, Miller JZ, et al (1980). "Racial differences in blood pressure in Evans County, Georgia: relationship to sodium and potassium intake and plasma renin activity". J Chronic Dis 33 (2): 87–94. doi:10.1016/0021-9681(80)90032-6. PMID 6986391. ^ Karger, S. (2004). "Energy and nutrient intake in the European Union" (pdf). Ann Nutr Metab 48 (2 (suppl)): 1–16, http://content.karger.com/ProdukteDB/produkte.asp?Aktion=ShowPDF&ProduktNr=223977&Ausgabe=230671&ArtikelNr=83312&filename=83312.pdf. ^ Mineno, T (1969). "Effect of some vitamins and other substances on K metabolism in the myocardia of vitamin deficient rats - Experimental investigation". J. Nagoya Med. Assoc. 92;: 80–95. ^ Gould, SE (ed) (1968). Pathology of the Heart and Blood Vessels, Charles C. Thomas. pp.851. p. 508.^ Berliner RW, Kennedy TJ, Orloff J (1951). "Relationship between acidification of the urine and potassium metabolism; effect of carbonic anhydrase inhibition on potassium excretion". Am. J. Med. 11 (3): 274–82. doi:10.1016/0002-9343(51)90165-9. PMID 14877833. ^ Knochel JP (1984). "Diuretic-induced hypokalemia". Am. J. Med. 77 (5A): 18–27. doi:10.1016/S0002-9343(84)80004-2. PMID 6496556. ^ Kolb H, Burkart V (1999). "Nicotinamide in type 1 diabetes. Mechanism of action revisited". Diabetes Care 22 Suppl 2: B16–20. PMID 10097894. ^ Kilburn KH (1966). "Movements of potassium during acute respiratory acidosis and recovery". J Appl Physiol 21 (2): 679–84. PMID 5934480. ^ Consolazio CF, Matoush LO, Nelson RA, Harding RS, Canham JE (1963). "Excretion of sodium, potassium, magnesium and iron in human sweat and the relation of each to balance and requirements". J. Nutr. 79: 407–15. PMID 14022653. ^ )Potassium / K (mg.) Content of Selected Foods per Common Measure, sorted by nutrient content | USDA National Nutrient Database for Standard Reference, Release 20 http://www.nal.usda.gov/fnic/foodcomp/Data/SR20/nutrlist/sr20w306.pdf^ Enghag, P. (2004). Encyclopedia of the elements, Wiley-VCH Weinheim. ^ Thomas K. Wray. "DANGER: PEROXIDIZABLE CHEMICALS". Environmental Health & Public Safety (North Carolina State University).[edit] See alsoPotassium compoundsPotassium in biology[edit] External links Wikimedia Commons has media related to: Potassium Look up potassium inWiktionary, the free dictionary.WebElements.com – Potassiumv • d • eAlkali metals LithiumLiAtomic Number: 3Atomic Weight: 6.941Melting Point: 453.69Boiling Point: 1615Electronegativity: 0.98SodiumNaAtomic Number: 11Atomic Weight: 22.990Melting Point: 370.87Boiling Point: 1156Electronegativity: 0.96PotassiumKAtomic Number: 19Atomic Weight: 39.098Melting Point: 336.58Boiling Point: 1032Electronegativity: 0.82RubidiumRbAtomic Number: 37Atomic Weight: 85.468Melting Point: 312.46Boiling Point: 961Electronegativity: 0.82CaesiumCsAtomic Number: 55Atomic Weight: 132.905Melting Point: 301.59Boiling Point: 944Electronegativity: 0.79FranciumFrAtomic Number: 87Atomic Weight: (223)Melting Point: ?295Boiling Point: ?950Electronegativity: 0.7v • d • ePeriodic tableH HeLiBe BCNOFNeNaMg AlSiPSClArKCa ScTiVCrMnFeCoNiCuZnGaGeAsSeBrKrRbSr YZrNbMoTcRuRhPdAgCdInSnSbTeIXeCsBaLaCePrNdPmSmEuGdTbDyHoErTmYbLuHfTaWReOsIrPtAuHgTlPbBiPoAtRnFrRaAcThPaUNpPuAmCmBkCfEsFmMdNoLrRfDbSgBhHsMtDsRgUubUutUuqUupUuhUusUuoUueUbn Alkali metalsAlkaline earth metalsLanthanoidsActinoidsTransition metalsOther metalsMetalloidsOther nonmetalsHalogensNoble gasesRetrieved from "http://en.wikipedia.org/wiki/Potassium" Categories: Alkali metals | Chemical elements | Desiccants | Dietary minerals | PotassiumHidden categories: All articles with unsourced statements | Articles with unsourced statements since December 2007 | Articles with unsourced statements since January 2008 | Articles with unsourced statements since September 2008 | Articles to be expanded since June 2008 | All articles to be expanded | Articles to be expanded since February 2008 Views Article Discussion Edit this page History Personal tools Log in / create account if (window.isMSIE55) fixalpha(); Navigation Main page Contents Featured content Current events Random article Search Interaction About Wikipedia Community portal Recent changes Contact Wikipedia Donate to Wikipedia Help Toolbox What links here Related changesUpload fileSpecial pages Printable version Permanent linkCite this page Languages Afrikaans العربية Asturianu বাংলা Bân-lâm-gú Беларуская Bosanski Български Català Česky Corsu Cymraeg Dansk Deutsch ދިވެހިބަސް Eesti Ελληνικά Español Esperanto Euskara فارسی Français Furlan Gaeilge Gaelg Galego 한국어 Hawai`i Հայերեն हिन्दी Hrvatski Ido Bahasa Indonesia Íslenska Italiano עברית Basa Jawa ಕನ್ನಡ Kiswahili Kreyòl ayisyen Kurdî / كوردی Latina Latviešu Lëtzebuergesch Lietuvių Lojban Magyar Македонски മലയാളം Māori Bahasa Melayu Nederlands 日本語 Norsk (bokmål) Norsk (nynorsk) Novial Occitan O'zbek ਪੰਜਾਬੀ Plattdüütsch Polski Português Română Runa Simi Русский Shqip Sicilianu Simple English Slovenčina Slovenščina Српски / Srpski Srpskohrvatski / Српскохрватски Suomi Svenska ไทย Tiếng Việt Тоҷикӣ Türkçe Українська اردو Walon West-Vlams 粵語 中文   This page was last modified on 2 December 2008, at 01:07. All text is available under the terms of the GNU Free Documentation License. (See Copyrights for details.) Wikipedia® is a registered trademark of the Wikimedia Foundation, Inc., a U.S. registered 501(c)(3) tax-deductible nonprofit charity. Privacy policy About Wikipedia Disclaimers if (window.runOnloadHook) runOnloadHook(); |
|
