a chemical element in group I of Mendeleev’s periodic system. Atomic number, 19; atomic weight, 39.098; a silver-white, very light, soft, and low-melting metal. The element consists of two stable isotopes, 39K (93.08 percent) and 41K (6.91 percent), and one slightly radioactive isotope, 40K (0.01 percent), with a half-life of 1.32 × 109 years.
Historical survey. Some potassium compounds (for example, potash prepared from wood ashes) have been known since antiquity. These compounds, however, were not distinguished from those of sodium. The difference between “vegetable caustic” (potash, K2 CO3 ) and “mineral caustic” (soda, Na2 C03) was established only in the 18th century. In 1807, H. Davy electrolyzed slightly moistened solid caustic potash and caustic soda (KOH and NaOH) to obtain the corresponding metals, which he called potassium and sodium. In 1809, L. W. Gilbert proposed the names “kalium” (from the Arabic al-kali, “potash”) and “natronium” (from the Arabic natrun, “natural soda”); the latter name was changed to “natrium” by J. J. Berzelius in 1811. The names “potassium” and “sodium” have been retained in Great Britain, the USA, France, and certain other countries. These names were replaced in Russia during the 1840’s by “kalium” and “natrium,” which had been accepted in Germany, Austria, and the Scandinavian countries.
Distribution in nature. Potassium occurs widely in nature. Its content in the lithosphere is 2.50 weight percent. Potassium, like sodium, is enriched acidic magmas from which granites and other rocks crystallize (average K content, 3.34 percent). Potassium is a constituent of feldspars and micas. Basic and ultrabasic rocks, which are rich in iron and magnesium, contain little potassium. In contrast to sodium, potassium migrates only slightly on the earth’s surface. Weathering of rocks leads to partial transfer of potassium to water, but this is rapidly absorbed by organisms and clays; therefore, river waters are poor in potassium and much less potassium than sodium reaches the oceans. In the ocean, potassium is absorbed by organisms and bottom silts (it is, for example, a component of glauconite). For this reason ocean waters contain only 0.038 percent potassium (25 times less than sodium). Potassium and magnesium salts, such as carnallite (KCl⋅MgCl2⋅6H2O; Solikamsk deposit in the USSR, Stassfurt deposit in the German Democratic Republic), crystallized after the precipitation of NaCl during the later stages of the evaporation of seawater in lagunas that occurred during past geologic epochs (particularly during the Permian, about 200 million years ago). In most soils there are only small amounts of soluble potassium compounds; therefore, cultivated plants require potassium fertilizers.
The radioactive isotope 40K is an important source of underground heat; this was particularly true in past epochs, when the quantities of the isotope were larger. The decay of 40K leads to the formation of 40Ca and 40Ar, which escapes into the atmosphere. Some potassium-containing minerals do not suffer the loss of argon, so that argon content may be used for determining the absolute age of the rock (the potassium-argon method).
Physical and chemical properties. Potassium is a silver-white, very light and soft metal that may be cut with a knife without difficulty. The crystal lattice of potassium is body-centered cubic; α equals 5.33 Å and the atomic radius, 2.36 Å. The ionic radius of K+ is 1.33 Å. Its density is 0.862 g per cu cm (at 20°C). Its melting point is 63.55°C; its boiling point, 760°C. Its coefficient of thermal expansion is 8.33 × 10-5 (0°-50°C). Its thermal conductivity at 21°C is 97.13 watts per (m.°K), or 0.232 cal per (cm.sec⋅°C). Its specific heat at 20°C is 741.2 joules per (kg⋅°K), or 0.177 cal per (g⋅°C); its specific electrical resistivity at 20°C, 7.118 × 10−8 ohm⋅m; and its temperature coefficient of electrical resistivity, 5⋅8 × 10-5 (20°C). The Brinell hardness of potassium is 400 kilo new tons per sq m, or 0.04 kilograms-force per sq mm.
The outer electron-shell configuration of the potassium atom is 4s1, so that the valence of potassium in compounds is always 1. The sole valence electron of the potassium atom is at a greater distance from its nucleus than the corresponding valence electrons in lithium and sodium; therefore, the chemical reactivity of potassium is higher than that of the other two metals. Since potassium is rapidly oxidized in air (particularly in moist air), the metal is stored under gasoline, kerosene, or mineral oil. Potassium reacts at room temperature with the halogens. It unites with sulfur on mild heating; stronger heating is required for reaction with selenium and tellurium. When heated above 200°C in a hydrogen atomosphere, potassium forms the hydride KH, which is spontaneously combustible in air. Potassium does not react with nitrogen even when heated under pressure, but the two elements form potassium nitride K3N and potassium azide KN3 under the influence of an electric discharge. Heating potassium with graphite yields the carbides KC8 (at 300°C) and KC16 (at 360°C). In dry air or oxygen, potassium forms the yellowish-white oxide K2O and the orange peroxide K02; the known peroxides also include K2O2 and K2O3, which are prepared by the action of oxygen on potassium solutions in liquid ammonia.
Potassium reacts extremely vigorously (sometimes explosively) with water to evolve hydrogen (2K + 2H2O = 2KOH + H2) and with aqueous acid solutions to give salts. Potassium dissolves slowly in ammonia; the resulting blue solution is a powerful reducing agent. When heated, potassium removes the oxygen in oxides and the salts of oxyacids to form K2O and free metals or their oxides. Potassium reacts with alcohols to give alkoxides; it accelerates the polymerization of olefins and di-olefins; and it forms potassium alkyls and potassium aryls in reactions with alkyl halides and aryl halides, respectively. The presence of potassium may be readily determined by the violet coloration of its flame.
Production and use. Potassium is produced industrially by using exchange reactions between metallic sodium and KOH or KC1:
KOH + Na = NaOH + K
KC1 + Na = NaCl + K
In the first case the molten hydroxide KOH reacts with liquid sodium in countercurrent in a nickel plate column at 380°–440°C. In the second case, sodium vapor is passed through the molten salt KC1 at 760°–800°C and the evolved potassium vapor is condensed. Potassium production may also be performed by heating (above 200°C) mixtures of potassium chloride with aluminum (or silicon) and lime. The production of potassium by the electrolysis of molten KOH or KC1 is limited because of a low level of efficiency and safety problems.
The main use of metallic potassium is in the preparation of potassium peroxide, which is used for oxygen regeneration in submarines and other closed systems. Alloys of sodium with 40–90 percent potassium, which remain liquid at room temperature, are used as heat transfer agents in nuclear reactors, as reducing agents in titanium production, and as oxygen absorbers. Agriculture is the principal consumer of potassium salts (for use in fertilizers).
REFERENCES“Kalii.” In Kratkaia khimicheskaia entsiklopediia, vol. 2. Moscow, 1963.
Nekrasov, B. V. Osnovy obschchei khimii, vol. 3. Moscow, 1970.
Remy, G. Kurs neorganicheskoi khimii, vol. 1. Moscow, 1963. (Translated from German.)
Potassium ions participate in generating and propagating bioelectric potentials in the nerves and muscles, in regulating the contraction of heart and other muscles, in maintaining the osmotic pressure and hydration of colloids in the cells, and in activating certain enzymes. Potassium metabolism is closely associated with carbohydrate metabolism, and potassium ions affect protein synthesis. In most cases, K+ cannot replace Na+. K+ is selectively concentrated by the cells. Suppression of glycolysis, respiration, and photosynthesis and impairment of the permeability of the outer cell membrane lead to a loss of potassium from cells, frequently in exchange for Na+. Potassium is eliminated from the organism mainly with the urine. The potassium content of the blood and tissues in vertebrates is regulated by hormones of the adrenal glands (the corticosteroids).
Potassium is unevenly distributed within plants; the vegetative organs contain more potassium than do the roots and seeds. Large amounts of potassium are found in beans, beets, potatoes, tobacco leaves, and gramineous feed grains (20–30 g per kg of dry material). Potassium deficiency in the soil retards plant growth and increases the incidence of disease. The amount of potassium used in potassium fertilizers depends on the type of crop and soil.
The trace elements rubidium and cesium accompany potassium in the biosphere. Li+ and Na+ ions are antagonists of K+; for this reason, an optimum K+/Na+ ratio in the cells, as well as in the surrounding medium, is as important as the absolute concentrations of the ions. Almost 90 percent of the natural radioactivity of organisms (gamma radiation) is caused by the presence of the natural radioisotope 40K in the tissues.
REFERENCESKaplanskii, S. Ia. Mineral’nyi obmen. Moscow-Leningrad, 1938.
Vishniakov, S. I. Obmen makroelementov u sel’skokhoziaistvennykh zhivotnykh. Moscow, 1967.
Sutcliffe, J. F. Pogloshchenie mineral’nykh solei rasteniiami. Moscow, 1964. (Translated from English.)
I. A. SKUL’SKII
In medicine, the acetate CH3COOK is used as a diuretic, mainly against edema caused by cardiac insufficiency. The chloride KC1 is used in cases of potassium deficiency in the organism. Such deficiencies develop during treatment with certain hormonal preparations and digitalis, after great fluid losses brought about by vomiting or diarrhea, and with the administration of certain diuretics. The perchlorate KCIO4 inhibits thyroxine production (a hormone of the thyroid gland) and is used in cases of thyrotoxicosis. Potassium permanganate, KCIO4, is used as anantiseptic.