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silvery gray
General properties
Name, symbol, number potassium, K, 19
Element category alkali metal
Group, period, block 14, s
Standard atomic weight 39.0983(1)g·mol−1
Electron configuration [Ar] 4s1
Electrons per shell 2, 8, 8, 1 (Image)
Physical properties
Phase solid
Density (near r.t.) 0.89 g·cm−3
Liquid density at m.p. 0.828 g·cm−3
Melting point 336.53 K, 63.38 °C, 146.08 °F
Boiling point 1032 K, 759 °C, 1398 °F
Triple point 336.35 K (63°C),  kPa
Heat of fusion 2.33 kJ·mol−1
Heat of vaporization 76.9 kJ·mol−1
Specific heat capacity (25 °C) 29.6 J·mol−1·K−1
Atomic properties
Oxidation states 1
(strongly basic oxide)
Electronegativity 0.82 (Pauling scale)
Ionization energies
1st: 418.8 kJ·mol−1
2nd: 3052 kJ·mol−1
3rd: 4420 kJ·mol−1
Atomic radius 227 pm
Covalent radius 203±12 pm
Van der Waals radius 275 pm
Crystal structure body-centered cubic
Magnetic ordering paramagnetic
Electrical resistivity (20 °C) 72 nΩ·m
Thermal conductivity (300 K) 102.5 W·m−1·K−1
Thermal expansion (25 °C) 83.3 µm·m−1·K−1
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
Main article: Isotopes of potassium
iso NA half-life DM DE (MeV) DP
39K 93.26% 39K is stable with 20 neutrons
40K 0.012% 1.248(3)×109 y β 1.311 40Ca
ε 1.505 40Ar
β+ 1.505 40Ar
41K 6.73% 41K is stable with 22 neutrons

Potassium (pronounced /pɵˈtæsiəm/ po-TAS-ee-əm) is the chemical element with the symbol K (Latin: kalium, from Arabic: القَلْيَهal-qalyah "plant ashes", cf. Alkali from the same root), atomic number 19, and atomic mass 39.0983. Potassium was first isolated from potash. Elemental potassium is a soft silvery-white metallic alkali metal that oxidizes rapidly in air and is very reactive with water, generating sufficient heat to ignite the evolved hydrogen.

Potassium in nature occurs only as ionic salt. As such, it is found dissolved in seawater, and as part of many minerals. Potassium ion is necessary for the function of all living cells, and is thus present in all plant and animal tissues. It is found in especially high concentrations in plant cells, and in a mixed diet, it is most highly concentrated in fruits.

Potassium and sodium are chemically similar, since both are alkali metals. However, their functions in organisms are quite different, especially in animal cells.



Potassium in feldspar

Elemental potassium does not occur in nature because it reacts violently with water. As various compounds, potassium makes up about 1.5% of the weight of the Earth's crust and is the seventh most abundant element. As it is very electropositive and highly reactive potassium metal is difficult to obtain from its minerals.[1]

History of the free element

Elemental 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. The structure of potash was not then known, but is now understood to be mostly potassium carbonate. By heating, the carbonate could be freed of carbon dioxide, leaving "caustic potash", so called because it caused chemical burns in contact with human tissue.

Potassium metal was discovered 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. Before the 18th century, no distinction was made between potassium and sodium. Potassium was the first metal that was isolated by electrolysis.[2] Davy extracted sodium by a similar technique, demonstrating the elements to be different.[3]

Commercial Production

Pure potassium metal may 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

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 much lower than that of sodium.[4][5]

Potassium metal in reagent-grade sells for about $10.00/pound ($22/kg) in 2010 when purchased in tonnage quantities. Lower purity metal sells for considerably less. The market in this metal is volatile due to the difficulty in its long term storage. 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 which will detonate when scratched. The resulting explosion will usually start a fire which is difficult to extinguish.[6]

Kilogram quantities of potassium cost far more, in the range of $700/kg. This is partially due to the cost of hazardous material shipping requirements. [7]


There 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% of decays) by electron capture or positron emission, or decays to stable 40Ca (88.8% of decays) 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.[8] The activity of natural potassium is 31 Bq/g.



The flame-test color for potassium

Potassium 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.[9] Potassium concentration in solution is commonly determined by flame photometry, atomic absorption spectrophotometry, inductively coupled plasma, or ion selective electrodes.


Potassium must be protected from air for storage to prevent disintegration of the metal from oxide and hydroxide corrosion. Often samples are maintained under a hydrocarbon medium which does not react with alkali metals, such as mineral oil or 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.

2 K(s) + 2 H2O(l) → H2(g) + 2 KOH(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

Potassium cations in the body

Biochemical function

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.[10] This ion pump uses ATP to pump 3 sodium ions out of the cell and 2 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 hyperpolarisation, 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, MthK) with a determined structure.[11] All five are from prokaryotic species.

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.[12]

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.[13]

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. Deficiency symptoms include muscle weakness, paralytic ileus, ECG abnormalities, decreased reflex response and in severe cases respiratory paralysis, alkalosis and cardiac arrhythmia.

Filtration and excretion

Potassium 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 (3.345 grams) and potassium makes up most of the cell fluid cations at about 150 milliequivalents per liter (4.8 grams). Plasma is filtered through the glomerulus of the kidneys in enormous amounts, about 180 liters per day.[14] Thus 602 grams of sodium and 33 grams of potassium are filtered each day. All but the 1–10 grams of sodium and the 1–4 grams 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 0.190 grams) per liter.[15] 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 passively[16][17] in counter flow to sodium in response to an apparent (but not actual) Donnan equilibrium,[18] 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.[19] 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,[20] 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 (3.6 to 5.2 mmol/L) 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 only one kind of ion can stream through, at 10 million to 100 million ions per second.[21] The pores require calcium in order to open[22] although it is thought that the calcium works in reverse by blocking at least one of the pores.[23] Carbonyl groups inside the pore on the amino acids mimics the water hydration that takes place in water solution[24] by the nature of the electrostatic charges on four carbonyl groups inside the pore.[25]

Potassium in the diet and by supplement

Adequate intake

A potassium intake sufficient to support life can generally be guaranteed by eating a variety of foods, especially plant 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 eating a balanced diet. Foods rich in potassium include orange juice, potatoes, bananas, avocados, cantaloupes, tomatoes, broccoli, soybeans, brown rice, garlic and apricots, although it is also abundant in most fruits, vegetables and meats.[26]

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.[27] With these findings, the question of what is the intake of potassium consistent with optimal health, is debated. For example, the 2004 guidelines of the Institute of Medicine specify a DRI of 4,000 mg of potassium (100 mEq), though most Americans consume only half that amount per day, which would make them formally deficient as regards this particular recommendation.[28] Similarly, in the European Union, particularly in Germany and Italy, insufficient potassium intake is somewhat common.[29]

Medical supplementation and disease

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 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.[12] Typical medical supplemental doses range from 10 milliequivalents (400 mg, about equal to a cup of milk or 6 oz. of orange juice) to 20 milliequivalents (800 mg) 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 U.S. to only 99 mg 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.


About 93% of the world potassium production was consumed by the fertilizer industry.[5]

Biological applications

Potassium and magnesium sulfate fertilizer

Potassium ions are an essential component of plant nutrition and are 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 cell function. They participate in the Na-K pump.

In the form of potassium chloride, it is used to stop the heart, e.g. in cardiac surgery and execution by lethal injection.

Food applications

Potassium 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, potatoes, bananas and many other good dietary sources of potassium, ranked according to potassium content per measure shown.[30]

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.

Industrial applications

Potassium vapor is used in several types of magnetometers. An alloy of sodium and potassium, NaK (usually pronounced "nack"), 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 inks, 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).

4 KO2 + 2 CO2 → 2 K2CO3 + 3 O2

Potassium 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.

Potassium cobaltinitrite K3[Co(NO2)6] is used as artist's pigment under the name of Aureolin or Cobalt yellow.


Potassium reacts very violently with water producing potassium hydroxide (KOH) and hydrogen gas.

2K (s) + 2H2O (l) → 2KOH (aq) + H2 (g)

This reaction is exothermic and temperature produced is sufficient to ignite the resulting hydrogen. It in turn may explode in the presence of oxygen. Potassium hydroxide is a strong alkali which causes skin burns.

Finely divided potassium will ignite in air at room temperature. The bulk metal will ignite in air if heated. Water makes a potassium fire worse. Because its density is 0.89, burning potassium floats which exposes it to more atmospheric oxygen. The water also produces potentially explosive hydrogen gas. Many common fire extinguishing agents are either ineffective or make a potassium fire worse. 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 powders deprive the fire of oxygen and cool the potassium metal. Nitrogen or argon are also effective.

Potassium reacts violently in the presence of halogens and will detonate in the presence of bromine. It also reacts explosively with sulphuric acid. During combustion potassium forms peroxides and superoxides . These peroxides may react violently with organics present such as oils. Both peroxides and superoxides may react explosively with metallic potassium. [31]

Since 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 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 not be stored for longer than three months unless stored in an inert (oxygen free) atmosphere, or under vacuum.[32]

Due to the highly reactive nature of potassium metal, it must be handled with great care, with full skin and eye protection being used and preferably an explosive resistant barrier between the user and the potassium.

See also


  1. ^ a b c Mark Winter. "Potassium: Key Information". Webelements. 
  2. ^ Enghag, P. (2004). "11. Sodium and Potassium". Encyclopedia of the elements. Wiley-VCH Weinheim. ISBN 3527306668. 
  3. ^ 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 98: 1–45. doi:10.1098/rstl.1808.0001. 
  4. ^ Ober, Joyce A.. "Mineral Commodity Summaries 2008:Potash". United States Geological Survey. Retrieved 2008-11-20. 
  5. ^ a b Ober, Joyce A.. "Mineral Yearbook 2006:Potash". United States Geological Survey. Retrieved 2008-11-20. 
  6. ^ Potassium Metal 98.50% Purity
  7. ^ 004 – Potassium Metal
  8. ^ "background radiation – potassium-40 – γ radiation". 
  9. ^ Anne Marie Helmenstine. "Qualitative Analysis – Flame Tests". 
  10. ^ Campbell, Neil (1987). Biology. Menlo Park, Calif.: Benjamin/Cummings Pub. Co.. p. 795. ISBN 0-8053-1840-2. 
  11. ^ Mikko Hellgren, Lars Sandberg, Olle Edholm (2006). "A comparison between two prokaryotic potassium channels (KirBac1.1 and KcsA) in a molecular dynamics (MD) simulation study". Biophys. Chem. 120 (1): 1–9. doi:10.1016/j.bpc.2005.10.002. PMID 16253415. 
  12. ^ a b "Potassium Without the Taste". Retrieved Feb 14, 2009. 
  13. ^ Lockless SW, Zhou M, MacKinnon R.. "Structural and thermodynamic properties of selective ion binding in a K+ channel". Laboratory of Molecular Neurobiology and Biophysics, Rockefeller University. Retrieved 2008-03-08. 
  14. ^ Potts, W.T.W.; Parry, G. (1964). Osmotic and ionic regulation in animals. Pergamon Press. 
  15. ^ 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. 
  16. ^ Bennett CM, Brenner BM, Berliner RW (1968). "Micropuncture study of nephron function in the rhesus monkey". J. Clin. Invest. 47 (1): 203–216. doi:10.1172/JCI105710 (inactive 2010-03-18). PMID 16695942. 
  17. ^ Solomon AK (1962). "Pumps in the living cell". Scientific American 207: 100–8. doi:10.1038/scientificamerican0862-100. PMID 13914986. 
  18. ^ Kernan, Roderick P. (1980). Cell potassium (Transport in the life sciences). New York: Wiley. pp. 40, 48. ISBN 0471048062. 
  19. ^ 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. 
  20. ^ 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. 
  21. ^ 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. 
  22. ^ 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. 
  23. ^ 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.
  24. ^ 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. 
  25. ^ 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. 
  26. ^ "Potassium Content of Food and Drink". Retrieved 2008-09-18. 
  27. ^ Folis, R.H. (1942). "Myocardial Necrosis in Rats on a Potassium Low Diet Prevented by Thiamine Deficiency". Bull. Johns-Hopkins Hospital 71: 235. 
  28. ^ 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. 
  29. ^ Karger, S. (2004). "Energy and nutrient intake in the European Union" (pdf). Ann Nutr Metab 48 (2 (suppl)): 1–16. doi:10.1159/000083041. 
  30. ^ "Potassium Content of Selected Foods per Common Measure, sorted by nutrient content". USDA National Nutrient Database for Standard Reference, Release 20. 
  31. ^ DOE HANDBOOK-Alkali Metals Sodium, Potassium, NaK, and Lithium
  32. ^ Thomas K. Wray. "Danger: peroxidazable chemicals". Environmental Health & Public Safety (North Carolina State University). 

External links

1911 encyclopedia

Up to date as of January 14, 2010

From LoveToKnow 1911

POTASSIUM [[[symbol]] K (from kalium), atomic weight 39.114 0=16)], a metallic chemical element, belonging to the group termed the metals of the alkalis. Although never found free in nature, in combination the metal is abundantly and widely distributed. In the oceans alone there are estimated to be 1141 X 10 12 tons of sulphate, K 2 SO 4, but this inexhaustible store is not much drawn upon; and the "salt gardens" on the coast of France lost their industrial importance as potash-producers since the deposits at Stassfurt in Germany have come to be worked. These deposits, in addition to common salt, include the following minerals: sylvine, KC1; carnallite, KC1 MgC12.6H20 (transparent, deliquescent crystals, often red with diffused oxide of iron); kainite, K 2 SO 4 MgSO 4 MgC1 2 6H 2 O (hard crystalline masses, permanent in the air); kieserite MgS04 H20 (only very slowly dissolved by water); besides polyhalite, MgSO 4 K 2 SO 4.2CaSO 4.2H20anhydrite, CaSO 4; salt, NaC1, and some minor components. These potassium minerals are not confined to Stassfurt; larger quantities of sylvine and kainite are met with in the salt mines of Kalusz in the eastern Carpathian Mountains. The Stassfurt minerals owe their industrial importance to their solubility in water and consequent ready amenability to chemical operations. In point of absolute mass they are insignificant compared with the abundance and variety of potassiferous silicates, which occur everywhere in the earth's crust; orthoclase (potash felspar) and potash mica may be quoted as prominent examples. Such potassiferous silicates are found in almost all rocks, both as normal and as accessory components; and their disintegration furnishes the soluble potassium salts which are found in all fertile soils. These salts are sucked up by the roots of plants, and by taking part in the process of nutrition are partly converted into oxalate, tartrate, and other organic salts, which, when the plants are burned, are converted into the carbonate, K 2 CO 3. It is a remarkable fact that, although in a given soil the soda-content may predominate largely over the potash salts, the plants growing in the soil take up the latter: in the ashes of most land plants the potash (calculated as K20) forms upwards of 90% of the total alkali. The proposition holds, in its general sense, for sea plants likewise. In ocean water the ratio of soda (Na 2 O) to potash (K 2 0) is loo: 3.23 (Dittmar); in kelp it is, on the average, ioo: 5.26 (Richardson). Ashes particularly rich in potash are those of burning nettles, wormwood (Artemisia absinthium), tansy (Tanacetum vulgare), fumitory (Fumaria officinalis), and tobacco. In fact, the ashes of herbs generally are richer in potash than those of the trunks and branches of trees; yet, for obvious reasons, the latter are of greater industrial importance as sources of potassium carbonate. According to Liebig, potassium is the essential alkali of the animal body; and it may be noted that sheep excrete most of the potassium which they take from the land as sweat, one-third of the weight of raw merino consisting of potassium compounds.

To Sir Humphry Davy belongs the merit of isolating this element from potash, which itself had previously been considered an element. On placing a piece of potash on a platinum plate, connected to the negative of a powerful electric battery, and bringing a platinum wire, connected to the positive of the battery, to the surface of the potassium a vivid action was observed: gas was evolved at the upper surface of the fused globule of potash, whilst at the lower surface, adjacent to the platinum plate, minute metallic globules were formed, some of which immediately inflamed, whilst others merely tarnished. In 1808 Gay-Lussac and Thenard (Ann. chim. 6 5, p. 3 2 5) obtained the metal by passing melted potash down a clay tube containing iron turnings or wire heated to whiteness, and Caradau (ibid. 66, p. 97) effected the same decomposition with charcoal at a white heat. This last process was much improved by Brunner, Wdhler, and especially by F. M. L. Donny and J. D. B. Mareska (Ann. chim. phys., 1852, (3), 35, p. 1 47). Brunner's process consisted in forming an intimate mixture of potassium carbonate and carbon by igniting crude tartar in covered iron crucibles, cooling the mass, and then distilling it at a white heat from iron bottles, the vaporized metal being condensed beneath the surface of paraffin or naphtha contained in a copper vessel. It was found, however, that if the cooling be not sufficiently rapid explosions occurred owing to the combination of the metal with carbon monoxide (produced in the oxidation of the charcoal) to form the potassium salt of hexaoxybenzene. In Mareska and Donny's process the condensation is effected in a shallow iron box, which has a large exposed surface, capable of being cooled by damped cloths. When the distillation is finished the iron box, after cooling, is unclamped and the product turned out beneath the surface of paraffin. It is purified by redistilling and condensing directly under paraffin. Electrolytic processes have also been devised. Linnemann (Journ. Prak. Chem., 1858, 73, p. 413) obtained the metal on a small scale by electrolysing potassium cyanide between carbon electrodes; A. Matthiessen (Journ. Chem. Soc., 1856, p. 30) electrolysed an equimolecular mixture of potassium and calcium chlorides (which melts at a lower temperature than potassium chloride) also between carbon electrodes; whilst Castner's process, in which caustic potash is electrolysed, is employed commercially. The metal, however, is not in great demand, for it is generally found that sodium, which is cheaper, and, weight for weight, more reactive, will fulfil any purpose for which potassium may be desired.

Pure potassium is a silvery white metal tinged with blue; but on exposure to air it at once forms a film of oxide, and on prolonged exposure deliquesces into a solution of hydrate and carbonate. Perfectly dry oxygen, however, has no action upon it. At temperatures below o° C. it is pretty hard and brittle; at the ordinary temperature it is so soft that it can be kneaded between the fingers and cut with a blunt knife. Its specific gravity is o 865; hence it is the lightest metal known except lithium. It fuses at 62.5°C. (Bunsen) and boils at 667°, emitting an intensely green vapour. It may be obtained crystallized in quadratic octahedra of a greenish-blue colour, by melting in a sealed tube containing an inert gas, and inverting the tube when the metal has partially solidified. When heated in air it fuses and then takes fire, burning into a mixture of oxides. Most remarkable, and characteristic for the group it represents, is its action on water. A pellet of potassium when thrown on water at once bursts out into a violet flame and the burning metal fizzes about on the surface, its extremely high temperature precluding absolute contact with the liquid, exce p t at the very end, when the last remnant, through loss of temperature, is wetted by the water and bursts with explosive violence. The reaction may be written 2K+ 211 2 0= 2K0H+H2, and the flame is due to the combustion of the hydrogen, the violet colour being occasioned by the potassium vapour. The metal also reacts with alcohol to form potassium ethylate, while hydrogen escapes, this time without inflammation: K+C 2 H 5. OH=C 2 H 5. OK+H. When the oxide-free metal is heated gently in dry ammonia it is gradually transformed into a blue liquid, which on cooling freezes into a yellowish-brown or flesh-coloured solid, potassamide, KNH 2. When heated to redness the amide is decomposed into ammonia and potassium nitride, NK 3, which is an almost black solid. Both it and the amide decompose water readily with formation of ammonia and caustic potash. Potassium at temperatures from 200 to 400°C. occludes hydrogen gas, the highest degree of saturation corresponding approximately to the formula K 2 H. In a vacuum or in sufficiently dilute hydrogen the compound from 200° upwards loses hydrogen, until the tension of the free gas has arrived at the maximum value characteristic of that temperature (Troost and Hautefeuille).

Compounds. Oxides and Hydroxide. - Potassium forms two well-defined oxides, K 2 0 and K204, whilst several others, of less certain existence, have been described. The monoxide, K 2 0, may be obtained by strongly heating the product or burning the metal in slightly moist air; by heating the hydroxide with the metal: 2KHO+2K= 2K 2 0+H 2; or by passing pure and almost dry air over the molten metal (Kiihnemann, Chem. Centr., 186 3, p. 49 1). It forms a grey brittle mass, having a conchoidal fracture; it is very deliquescent, combining very energetically with water to form caustic potash. According to Holt and Sims (Journ. Chem. Soc., 18 94, p. 43 8), the substance as obtained above always contains free potassium.

Potassium hydroxide or caustic potash, KOH, formerly considered to be an oxide but shown subsequently to be a hydroxide of potassium, may be obtained by dissolving the metal or monoxide in water, but is manufactured by double decomposition from potassium carbonate and slaked lime: K 2 CO 3 -E-Ca(OH) 2 =2KOH+CaC03. A solution of one part of the carbonate in 12 parts of water is heated to boiling in a cast-iron vessel (industrially by means of steampipes) and the milk of lime added in instalments until a sample of the filtered mixture no longer effervesces with an excess of acid. The mixture is then allowed to settle in the iron vessel, access of air being prevented as much as practicable, and the clear liquor is syphoned off. The remaining mud of calcium carbonate and hydrate is washed, by decantation, with small instalments of hot water to recover at least part of the alkali diffused throughout it, but this process must not be continued too long or else some of the lime passes into solution. The liquors after a concentration in iron vessels are now evaporated in a silver dish, until the heavy vapour of the hydrate is seen to go off. The residual oily liquid is then poured out into a polished iron tray, or into an iron mould to produce the customary form of "sticks," and allowed to cool. The solid must be at once bottled, because it attracts the moisture and carbonic acid of the air with great avidity and deliquesces. According to Dittmar (Journ. Soc. Chem. Ind., May 1884), nickel basins are far better adapted than iron basins for the preliminary concentration of potash ley. The latter begin to oxidize before the ley has come up to the traditional strength of specific gravity 1.333 when cold, while nickel is not attacked so long as the percentage of real KHO is short of 60. For the fusion of the dry hydrate nickel vessels cannot be used; in fact, even silver is perceptibly attacked as soon as all the excess of water is away; absolutely pure KHO can be produced only in gold vessels. Glass and (to a less extent) porcelain are attacked by caustic potash ley, slowly in the cold, more readily on boiling.

Solid caustic potash forms an opaque, white, stone-like mass of dense granular fracture; specific gravity 2'1. It fuses considerably below and is perceptibly volatile at a red heat. At a white heat the vapour breaks down into potassium, hydrogen and oxygen. It is extremely soluble in even cold water, and in any proportion of water on boiling. On crystallizing a solution, the hydrate KOH 2H 2 0 is deposited; 2KOH 9H 2 0 and 2KOH 5H 2 0 have also been obtained. The solution is intensely "alkaline" to testpapers. It readily dissolves the epidermis of the skin and many other kinds of animal tissue - hence the former application of the "sticks" in surgery. A dilute potash readily emulsionizes fats, and on boiling saponifies them with formation of a soap and glycerin. All commercial caustic potash is contaminated with excess of water (over and above that in the KHO) and with potassium carbonate and chloride; sulphate, as a rule, is absent. A preparation sufficing for most purposes is obtained by digesting the commercial article in absolute alcohol, decanting and evaporating the solution to dryness and fusing in silver vessels.

The peroxide, K204, discovered by Gay-Lussac and Thenard, is obtained by heating the metal in an excess of slightly moist air or oxygen. Vernon Harcourt (Journ. Chem. Soc., 1862, p. 267) recommends melting the metal in a flask filled with nitrogen and gradually displacing this gas by oxygen; the first formed grey film on the metal changes to a deep blue, and then the gas is rapidly absorbed, the film becoming white and afterwards yellow. It is a dark yellow powder, which fuses at a high temperature, the liquid on cooling depositing shining tabular crystals; at a white heat it loses oxygen and yields the monoxide. Exposed to moist air it loses oxygen, possibly giving the dioxide, K 2 0 2; water reacts with it, evolving much heat and giving caustic potash, hydrogen peroxide and oxygen; whilst carbon monoxide gives potassium carbonate and oxygen at temperatures below loo°. A violent reaction ensues with phosphorus and sulphur, and many metals are oxidized by it, some with incandescence.

Halogen Compounds. - Potassium fluoride, KF, is a very deliquescent salt, crystallizing in cubes and having a sharp saline taste, which is formed by neutralizing potassium carbonate or hydroxide with hydrofluoric acid and concentrating in platinum vessels. It forms the acid fluoride KHF 2 when dissolved in aqueous hydrofluoric acid, a salt which at a red heat gives the normal fluoride and hydrofluoric acid. Other salts of composition KF 2HF and KF 3HF, have been described by Moissan (Comet. rend., 1888, 106, p. 547).

Potassium chloride, KC1, also known as muriate of potash, closely resembles ordinary salt. It is produced in immense quantities at Stassfurt from the so-called "Abraumsalze." For the purpose of the manufacturer of this salt these are assorted into a raw material containing approximately, in Ioo parts, 55-65 of carnallite (representing 16 parts of potassium chloride), 20-25 of common salt, 15-20 of kieserite; 2-4 of tachhydrite (CaC12 2MgC12 12H20), and minor components. This mixture is now wrought mainly in two ways. (1) The salt is dissolved in water with the help of steam, and the solution is cooled down to from 60° to 70°, when a quantity of impure common salt crystallizes out, which is removed. The decanted ley deposits on standing a 70% potassium chloride, which is purified by washing with cold water. Common salt principally goes into solution, and the percentage may thus be brought up to from 80 to 95. The mother-liquor from the 70% chloride is evaporated, the common salt which separates out in the heat removed as it appears, and the sufficiently concentrated liquor allowed to crystallize, when almost pure carnallite separates out, which is easily decomposed into its components '(see' infra). (2) Ziervogel and Tuchen's method. - The crude salt is ground up and then heated in a concentrated solution of magnesium chloride with agitation. The carnallite principally dissolves and crystallizes out relatively pure on cooling. The mother-liquor is used for a subsequent extraction of fresh raw salt. The carnallite produced is dissolved in hot water and the solution allowed to cool, when it deposits a coarse granular potassium chloride containing up to 99% of the pure substance. The undissolved residue produced in either process consists chiefly of kieserite and common salt. It is worked up either for Epsom salt and common salt, or for sodium sulphate and magnesium chloride. The potassiferous by-products are utilized for the manufacture of manures.

Chemically pure chloride of potassium is most conveniently prepared from the pure perchlorate by heating it in a platinum basin at the lowest temperature and then fusing the residue in a wellcovered platinum crucible. The fused product solidifies on cooling into a colourless glass.

When hydrochloric acid gas is passed into the solution the salt is completely precipitated as a fine powder. If the original solution contained the chlorides of magnesium or calcium or sulphate of potassium all impurities remain in the mother-liquor (the sulphur as KHS04), and can be removed by washing the precipitate with strong hydrochloric acid. The salt crystallizes in cubes of specific gravity 1.995; it melts at about 800° and volatilizes at a bright red heat. When melted in a current of hydrogen or electrolysed in the same condition, a dark blue mass is obtained of uncertain composition. It is extensively employed for the preparation of other potassium salts, but the largest quantity (especially of the impure product) is used in the production of artificial manures.

Potassium bromide, KBr, may be obtained by dissolving bromine in potash, whereupon bromide and bromate are first formed, evaporating and igniting the product in order to decompose the bromate: 6KHO 3Br 2 =5KBr -}- KBrO 3 -}- 3H 2 0; 2KBrO 3 = 2KBr + 302 (cf. Chlorates); but it is manufactured by acting with bromine water on iron filings and decomposing the iron bromide thus formed with potassium carbonate. In appearance it closely resembles the chloride, forming colourless cubes which readily dissolve in water and melt at 722°. It combines with bromine to form an unstable tribromide, KBr 3 (see F. P. Worley, Journ. Chem. Soc., 1905, 87, p. 1107).

Potassium iodide, KI, is obtained by dissolving iodine in potash, the deoxidation of the iodate being facilitated by the addition of charcoal before ignition, proceeding as with the bromide. The commercial salt usually has an alkaline reaction; it may be purified by dissolving in the minimum amount of water, and neutralizing with dilute sulphuric acid; alcohol is now added to precipitate the potassium sulphate, the solution filtered and crystallized. It forms colourless cubes which are readily soluble in water, melt at 685°, and yield a vapour of normal density. It is sparingly soluble in absolute alcohol. Both the iodide and bromide are used in photography. Iodine dissolves in an aqueous solution of the salt to form a dark brown liquid, which on evaporation over sulphuric acid gives black acicular crystals of the tri-iodide, K1 3. This salt is very deliquescent; it melts at 45°, and at 100° decomposes into iodine and potassium iodide. For the oxyhalogen salts see Chlorate, Chlorine, Bromine and Iodine.

Potassium carbonate, K 2 CO 3, popularly known as "potashes," was originally obtained in countries where wood was cheap by lixiviating wood ashes in wooden tubs, evaporating the solution to dryness in iron pots and calcining the residue; in more recent practice the calcination is carried out in reverberatory furnaces. This product, known as "crude potashes," contains, in addition to carbonate, varying amounts of sulphate and chloride and also insoluble matter. Crude potash is used for the manufacture of glass, and, after being causticized, for the making of soft soap. For many other purposes it must be refined, which is done by treating the crude product with the minimum of cold water required to dissolve the carbonate, removing the undissolved part (which consists chiefly of sulphate), and evaporating the clear liquor to dryness in an iron pan. The purified carbonate (which still contains most of the chloride of the raw material and other impurities) is known as "pearl ashes." Large quantities of carbonate used to be manufactured from the aqueous residue left in the distillation of beet-root spirit, i.e. indirectly from beet-root molasses. The liquors are evaporated to dryness and the residue is ignited to obtain a very impure carbonate, which is purified by methods founded on the different solubilities of the several components. Most of the carbonate which now occurs in commerce is made from the chloride of the Stassfurt beds by an adaptation of the "Leblanc process" for the conversion of common salt into soda ash (see Alkali Manufacture).

Chemically pure carbonate of potash is best prepared by igniting pure bicarbonate (see below) in iron or (better) in silver or platinum vessels, or else by calcining pure cream of tartar. The latter operation furnishes an intimate mixture of the carbonate with charcoal, from which the carbonate is extracted by lixiviation with water and filtration. The filtrate is evaporated to dryness (in iron or platinum vessels) and the residue fully dehydrated by gentle ignition. The salt is thus obtained as a white porous mass, fusible at a red heat (838° C., Carnelley) into a colourless liquid, which solidifies into a white opaque mass. The dry salt is very hygroscopic; it deliquesces into an oily solution ("oleum tartari") in ordinary air. The most saturated solution contains 205 parts of the salt to 100 of water and boils at 135°. On crystallizing a solution monoclinic crystals of 2K2C03.3H20 are deposited, which at 100° lose water and give a white powder of K 2 CO 3 H 2 0; this is completely dehydrated at 130°. The carbonate, being insoluble in strong alcohol (and many other liquid organic compounds), is much used for dehydration of the corresponding aqueous preparations. The pure carbonate is constantly used in the laboratory as a basic substance generally, for the disintegration of silicates, and as a precipitant. The industrial preparation serves for the making of flint glass, of potash soap (soft soap) and of caustic potash.

Potassium bicarbonate, Khco 3, is obtained when carbonic acid is passed through a cold solution of the ordinary carbonate as long as it is absorbed. Any silicate present is also converted into bicarbonate with elimination of silica, which must be filtered off. The filtrate is evaporated at a temperature not exceeding 60° or at most 70° C.; after sufficient concentration it deposits on cooling anhydrous crystals of the salt, while the potassium chloride, which may be present as an impurity, remains mostly in the motherliquor; the rest is easily removed by repeated recrystallization. If an absolutely pure preparation is wanted it is best to follow Water and start with the "black flux" produced by the ignition of pure bitartrate. The flux is moistened with water and exposed to a current of carbonic acid, which, on account of the condensing action of the charcoal, is absorbed with great avidity. The bicarbonate forms large monoclinic prisms, permanent in the air. When the dry salt is heated to 190° it decomposes into normal carbonate, carbon dioxide and water.

Potassium sulphide, K 2 S, was obtained by Berzelius in pale red crystals by passing hydrogen over potassium sulphate, and by Berthier as a flesh-coloured mass by heating the sulphate with carbon. It appears, however, that these products contain higher sulphides. On saturating a solution of caustic potash with sulphuretted hydrogen and adding a second equivalent of alkali, a solution is obtained which on evaporation in a vacuum deposits crystals of K 2 S.5H 2 O. The solution is strongly caustic. It turns yellow on exposure to air, absorbing oxygen and carbon dioxide and forming thiosulphate and potassium carbonate and liberating sulphuretted hydrogen, which decomposes into water and sulphur, the latter combining with the monosulphide to form higher salts. The solution also decomposes on boiling. The hydrosulphide, KHS, was obtained by Gay-Lussac on heating the metal in sulphuretted hydrogen, and by Berzelius on acting with sulphuretted hydrogen on potassium carbonate at a dull red heat. It forms a yellowishwhite deliquescent mass, which melts on heating, and at a sufficiently high temperature it yields a dark red liquid. It is readily soluble in water, and on evaporation in a vacuum over caustic lime it deposits colourless, rhombohedral crystals of 2KHS.H 2 0. The solution is more easily prepared by saturating potash solution with sulphuretted hydrogen. The solution has a bitter taste, and on exposure to the air turns yellow, but on long exposure it recovers its original colourless appearance owing to the formation of thiosulphate. Liver of sulphur or hepar sulphuris, a medicine known to the alchemists, is a mixture of various polysulphides with the sulphate and thiosulphate, in variable proportions, obtained by gently heating the carbonate with sulphur in covered vessels. It forms a liver-coloured mass. In the pharmacopoeia it is designated potassa sulphurata. Potassium sulphite, K 2 S0 3, is prepared by saturating a potash solution with sulphur dioxide, adding a second equivalent of potash, and crystallizing in a vacuum, when the salt separates as small deliquescent, hexagonal crystals. The salt K2S03 H20 may be obtained by crystallizing the metabisulphite, K 2 S 2 0 5 (from sulphur dioxide and a hot saturated solution of the carbonate, or from sulphur dioxide and a mixture of milk of lime and potassium sulphate) with an equivalent amount of potash. The salt K2S03 2H20 is obtained as oblique rhombic octahedra by crystallizing the solution over sulphuric acid. On the isomeric potassium sodium sulphites see Sulphur.

Potassium sulphate, K2S04, a salt known early in the 14th century, and studied by Glauber, Boyle and Tachenius, was styled in the 17th century arcanum or sal duplicatum, being regarded as a combination of an acid salt with an alkaline salt. It was obtained as a by-product in many chemical reactions, and subsequently used to be extracted from kainite, one of the Stassfurt minerals, but the process is now given up because the salt can be produced cheaply enough from the chloride by decomposing it with sulphuric acid and calcining the residue. To purify the crude product it is dissolved in hot water and the solution filtered and allowed to cool, when the bulk of the dissolved salt crystallizes out with characteristic promptitule. The very beautiful (anhydrous) crystals have the habit of a double six-sided pyramid, but really belong to the rhombic system. They are transparent, very hard and absolutely permanent in the air. They have a bitter, salty taste. The salt is soluble in water, but insoluble in caustic potash of sp. gr. 1.35, and in absolute alcohol. It fuses at 1078°. The crude salt is used occasionally in the manufacture of glass. The acid sulphate or bisulphate, Khso 4, is readily produced by fusing thirteen parts of the powdered normal salt with eight parts of sulphuric acid. It forms rhombic pyramids, which melt at 197°. It dissolves in three parts of water of o° C. The solution behaves pretty much as if its two congeners, K 2 SO 4 and H 2 SO 4, were present side by side of each other uncombined. An excess of alcohol, in fact, precipitates normal sulphate (with little bisulphate) and free acid remains in solution. Similar is the behaviour of the fused dry salt at a dull red heat; it acts on silicates, titanates, &c., as if it were sulphuric acid raised beyond its natural boiling point. Hence its frequent application in analysis as a disintegrating agent. For the salts of other sulphur acids, see Sulphur.

Potassamide, NH 2 K, discovered by Gay-Lussac and Thenard in 1871, is obtained as an olive green or brown mass by gently heating the metal in ammonia gas, or as a white, waxy, crystalline mass when the metal is heated in a silver boat. It decomposes in moist air, or with water, giving caustic potash and ammonia, in the latter case with considerable evolution of heat. On strong heating Tithesley (Journ. Chem. Soc., 18 94, p. 511) found that it decomposed into its elements. For the nitrite, see Nitrogen, for the nitrate see Saltpetre and for the cyanide see Prussic Acid; for other salts see the articles wherein the corresponding acid receives treatment.

Analysis, &c. - All volatile potassium compounds impart a violet coloration to the Bunsen flame, which is masked, however, if sodium be present. The emission spectrum shows two lines, Ka, a double line towards the infra-red, and Ka in the violet. The chief insoluble salts are the perchlorate, acid-tartrate and platinochloride. The atomic weight was determined by Stas and more recently by T. W. Richards and A. Stahler, who obtained the value 39.114 from analyses of the chloride, and by Richards and E. Meuller, who obtained the values 39.1135 and 39.1143 from analyses of the bromide '(see' Abs. C. S., 1907, ii. 615).

Medicine. Pharmacology. - Numerous salts and preparations of potassium are used in medicine; viz. Potassii Carbonis (salt of tartar), dose 5 to 20 grs., from which are made (a) Potassii Bicarbonas, dose 5 to 30 grs.; (b) Potassa Caustica, a powerful caustic not used internally. From caustic potash are made (I) Potassii Permanganas, dose 1 to 3 grs., used in preparing Liquor Potassii Permanganatis, a I A solution, dose 2 to 4 drs. (2) Potassii Iodidum, dose 5 to 20 grs.; from this are made the Linamentum Potassii Iodidi cum sapone, strength 1 in 10, and the Unguentum Potassii Iodidi, strength 1 in 10. (3) Potassii Bromidum, dose 5 to 30 grs. (4) Liquor Potassae, strength 27 grs. of caustic potash to the oz. Potassii Citras, dose To to 40 grs. Potassii Acetas, dose Do to 60 grs. Potassii Chloras, dose 5 to 15 grs., from which is made a lozenge. Trochiscus Potassii Chloratis, each containing 3 grs. Potassii Tartras Acidus (cream of tartar), dose 20 to 60 grs., which has a subpreparation Potassii Tartras, dose 30 to 60 grs. Potassii Nitras (saltpetre), dose 5 to 20 grs. Potassii Sulphas, dose To to 4 0 grs. Potassii Bicaromas, dose a to gr.


Poisoning by caustic potash may take place or poisoning by pearl ash containing caustic potash. A caustic taste in the mouth is quickly followed by burning abdominal pain, vomiting and diarrhoea, with a feeble pulse and a cold clammy skin; the post-mortem appearances are those of acute gastrointestinal irritation. The treatment is washing out the stomach or giving emetics followed by vinegar or lemon juice and later oil and white of egg.


Externally: Caustic potash is a most powerful irritant and caustic; it is used with lime in making Vienna paste, which is occasionally used to destroy morbid growths. Liquor potassae is also used in certain skin diseases. The permanganate of potash is an irritant if used pure. Its principal action is as an antiseptic and disinfectant. If wet it oxidizes the products of decomposition. It is used in the dressing of_ foul ulcers. The 1 solution is an antidote for snake-bite.

Internally: Dilate solutions of potash, like other alkalis, are used to neutralize the poisonous effects of strong acids. In the stomach potassium salts neutralize the gastric acid, and hence small doses are useful in hyperchloridia. Potassium salts are strongly diuretic, acting directly on the renal epithelium. 'They are quickly excreted in the urine, rendering it alkaline and thus more able to hold uric acid in solution. They also hinder the formation of uric acid calculi. The acetate and the citrate are valuable mild diuretics in Bright's disease and in feverish conditions, and by increasing the amount of urine diminish the pathological fluids in pleuritic effusion, ascites, &c. In tubal nephritis they aid the excretion of fatty casts. The tartrate and acid tartrate are also diuretic in their action and, as well as the sulphate, are valuable hydragogue saline purgatives. Potassium nitrate is chiefly used to make nitre paper, which on burning emits fumes useful in the treatment of the asthmatic paroxysm. Lozenges of potassium chlorate are used in stomatitis, tonsilitis and pharyngitis, it can also be used in a gargle, to grs. to I fl. oz. of water. Its therapeutic action is said to be due to nascent oxygen given off, so it is local in its action. In large doses it is a dangerous poison, converting the oxyhaemoglobin of the blood into methaemoglobin. Internally the permanganate is a valuable antidote in opium poisoning. The action of potassium bromide and potassium iodide has been treated under bromine and iodine (q.v.). All potassium salts if taken in large doses are cardiac depressants, they also depress the nervous system, especially the brain and spinal cord. Like all alkalis if given in quantities they increase metabolism.

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Simple English

[[File:|thumb|Potassium metal]] Potassium is a chemical element in the periodic table. It has the symbol K. This symbol is taken from the Latin word kalium. Potassium's atomic number is 19. It has 19 protons and electrons. Potassium is not found as an element in nature, because it is so reactive.

Potassium has two stable isotopes, with 20 or 22 neutrons. Its atomic mass is 39.098.



Physical properties

Potassium is a soft gray metal. It can be cut easily with a knife. Its melting point is 63 degrees celsius (145.4 degrees Fahrenheit). It melts at a very low temperature. It is an alkali metal. It is the second lightest metal, after lithium.

Chemical properties

File:Potassium water 20.theora.ogv
A 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.[1]

Potassium reacts in many chemical reactions similar to sodium and other alkali metals. It tarnishes in air to produce a whitish oxidized layer on the surface. This is why it is stored in oil. It also reacts very fast with water, which is another reason for its storage in oil. The hydrogen produced during its reaction with water can burst into flames when a large amount of potassium is added to water. Potassium hydroxide is also produced. Potassium also burns in air easily, to make the peroxide or the superoxide.

Chemical compounds

File:Flame Test
Potassium chloride in a flame

Potassium compounds are only in one oxidation state: +1. Potassium ions are colorless and similar to sodium ions. Potassium chloride can be used as a substitute for table salt. Potassium hydroxide is used in the electrolyte of alkaline cells. Most potassium compounds are nontoxic. If they are toxic, it is because of the anion. Potassium chromate is colored because of the chromate, not the potassium. Potassium chromate is toxic because of the chromate, not the potassium.


The word potassium comes from the word "potash". Potash is a chemical that is a mixture of potassium carbonate and potassium hydroxide that has been used for a very long time. It is used to make fertilizer, soap, and glass.


[[File:|thumb|Sylvite, a natural potassium chloride mineral]] Potassium does not occur in nature because it is too reactive. It is found in minerals, though. It is extracted from them by electrolysis of potassium hydroxide or potassium chloride. The potassium hydroxide or potassium chloride has to be melted at a very high temperature.

Use as element

It is used to absorb water from solvents. It is also used in some scientific instruments.

Use as compounds

Main article: Potassium compounds

Potassium compounds are used in soap, fertilizer, explosives, and matches.

Use by living organisms

Potassium ions are very important to organisms. They the body send messages from cells to other cells. It helps biological membranes depolarize. This means go from a negative to a positive electrical charge. This is needed for muscles to contract (get shorter and move things.) It is needed for the heart to beat (push blood through blood vessels.) If the potassium level in the blood is too high or too low it can cause death because the heart stops. A few good sources of potassium are bananas, apricots and raisins.


Potassium metal is very dangerous and can form an explosive coating if it is kept in air. It also reacts violently with water, spewing corrosive liquid all over. Potassium compounds are not normally dangerous, unless they contain a toxic anion like chromate or chlorate.

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