Potassium

Potassium (pronounced /pɵˈtæsiəm/ po-TAS-ee-əm) is the chemical element with the symbol K (Latin: kalium, from Arabic: القَلْيَه‎ al-qalyah "plant ashes" cf. Alkali from the same root, more commonly known in Modern Standard Arabic as بوتاسيوم ‹bwtasywm›), 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.

Occurrence

See also: Category:Potassium minerals

Elemental potassium does not occur in nature because it reacts violently with water (see section Precautions below). 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,[when?] 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]

Isotopes
Main article: isotopes of potassium

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.
Properties
Physical
The flame-test colour 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 lilac 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.

Chemical

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
Main article: Action potential

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
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.

Applications

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

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

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.
Precautions
A piece of potassium metal is dropped into a clear container containing water and slides around, burning with a bright pinkish or lilac flame for a short time until finishing with a pop and splash.

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.
Wiki letter w.svg This section requires expansion with:
additional citations.

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

* Category:Potassium compounds
* Potassium in biology

References

1. ^ a b c Mark Winter. "Potassium: Key Information". Webelements. http://www.webelements.com/webelements/elements/text/K/key.html.
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. http://books.google.com/books?id=Kg9GAAAAMAAJ.
4. ^ Ober, Joyce A.. "Mineral Commodity Summaries 2008:Potash". United States Geological Survey. http://minerals.usgs.gov/minerals/pubs/commodity/potash/mcs-2008-potas.pdf. Retrieved 2008-11-20.
5. ^ a b Ober, Joyce A.. "Mineral Yearbook 2006:Potash". United States Geological Survey. http://minerals.usgs.gov/minerals/pubs/commodity/potash/myb1-2006-potas.pdf. Retrieved 2008-11-20.
6. ^ Potassium Metal 98.50% Purity
7. ^ 004 – Potassium Metal
8. ^ "background radiation – potassium-40 – γ radiation". http://www.fas.harvard.edu/~scdiroff/lds/QuantumRelativity/RadioactiveHumanBody/RadioactiveHumanBody.html.
9. ^ Anne Marie Helmenstine. "Qualitative Analysis – Flame Tests". About.com. http://chemistry.about.com/library/weekly/aa110401a.htm.
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". http://www.foodnavigator.com/Science-Nutrition/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. http://www.ncbi.nlm.nih.gov/pubmed/17472437. 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". http://www.pamf.org/patients/pdf/potassium_count.pdf. 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. http://www.nal.usda.gov/fnic/foodcomp/Data/SR20/nutrlist/sr20w306.pdf.
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). http://www.ncsu.edu/ehs/www99/right/handsMan/lab/Peroxide.pdf.