Potassium

Introduction

Potassium was first isolated as a free metal in 1807 by Sir Humphry Davy. It was the first alkali metal to be discovered, produced by electrolysis of potassium carbonate (potash). The element was earlier called Kalium, derived from the Arabic word qili, meaning grass wort, the ash of which was a source of potash. The element derived its symbol K from Kalium. The English name potassium came from potash (pot ash), the carbonate salt of the metal.
Potassium is distributed widely in nature. The metal is too reactive to occur in native elemental form. It is the seventh most abundant element on earth, constituting 2.40% by weight of the earth’s crust. It is abundantly present in sea water. Oceans contain 0.07% (wt to volume) potassium chloride.
Potassium occurs in many igneous rocks, such as, feldspar (potassium aluminum silicate), KAlSi3O8 (leucite) and mica, KH2Al3(SiO4)3. Disintegration of these rocks adds potassium to soil and water. Deposits of potassium chloride are found in practically all salt beds, associated with sodium chloride. Some important potassium minerals are leucite, KAlSi2O6; glauconite (a complex silicoaluminate structure of varying compositions); sylvite, KCl; carnallite, KCl•MgCl2•6H2O; langbeinite, K2SO4•2MgSO4; and polyhalite, K2SO4•MgSO4•2CaSO4•2H2O.
Potassium, along with nitrogen and phosphorus, is an essential element needed for plant growth. In plants, it occurs mostly as K+ ion in cell juice. It is found in fruit or seed. Deficiency can cause curling leaves, yellow or brown coloration of leaves, weak stalk and diminished root growth. Potassium deficiency has been associated with several common animal ailments. Potassium is in extracellular fluid in animals at lower concentrations than sodium.

Potassium and health

Potassium in its ionic form, K+, is the most abundant positive ion in human and animal cells. As an electrolytic solution, K+ ions are pumped through the blood to all vital organs. Potassium's importance to the physiological system cannot be overstated: It plays a crucial role in electrical pulse transmission along nerve fibers; protein synthesis; acid-base balance; and formation of collagen, elastin, and muscle.

Potassium is highly soluble in water and regulates flow across semipermeable membranes like cell walls. It is this feature that makes a deficiency or an excess of potassium hazardous to health. Either extreme can have undesirable and even disastrous consequences.

Physical Properties

Silvery metal; body-centered cubic structure; imparts crimson-red color to flame; density 0.862g/cm3 at 20ºC; melts at 63.25ºC; density of liquid potassium at 100ºC is 0.819 g/cm3 and 0.771g/cm3 at 300ºC; vaporizes at 760ºC; vapor pressure 123 torr at 587ºC; electrical resistivity 6.1 microhm-cm at 0ºC and 15.31 microhm-cm at 100ºC; viscosity 0.25 centipoise at 250ºC; surface tension 86 dynes/cm at 100ºC; thermal neutron absorption cross section 2.07 barns; reacts violently with water and acids; reacts with alcohol; dissolves in liquid ammonia and mercury.

Production

Potassium can be produced by several methods that may be classified under three distinct types: (1) electrolysis, (2) chemical reduction, and (3) thermal decomposition.
Electrolysis processes have been known since Davy first isolated the metal in 1807. Electrolysis, however, suffers from certain disadvantages. A major problem involves miscibility of the metal with its fused salts. Because of this molten potassium chloride, unlike sodium chloride, cannot be used to produce the metal. Fused mixtures of potassium hydroxide and potassium carbonate or chloride have been used as electrolytes with limited success. Chemical reduction processes are employed nowadays in commercial, as well as, laboratory preparation of potassium. In one such process, molten potassium chloride is reduced with sodium at 760 to 880ºC and the free metal is separated by fractionation:
KCl + Na → K + NaCl
Potassium is obtained at over 99.5% purity. The metal, alternatively, may be alloyed with sodium for further applications.
Reduction of potassium fluoride with calcium carbide at 1,000 to 1,100ºC (Greisheim process) is an effective production method (Greer, J.S., Madaus, J.H and J.W. Mausteller. 1982. In Kirk-Othmer Encyclopedia of Chemical Technology, 3rd ed. p. 914, New York: Wiley Interscience):
2KF + CaC2 → CaF2 + 2C + 2K
Some other chemical reduction methods that may be applied for laboratory generation of small quantities of potassium from its salts at high temperatures require a suitable reducing agent such as carbon, calcium, or calcium carbide:
K2CO3 + 2C → 3CO +2K
2KCl + Ca → CaCl2 + 2K
2KCl + CaC2 → CaCl2 + 2C + 2K
2K2CO3 + +3Si + 3CaO → 4K + 2C + 3CaSiO3
2K2SiO3 + Si + 3 CaO → 4K + 3CaSiO3
Potassium can be produced by thermal decomposition of potassium azide:
2KN3 → 2K + 3N2
High purity metal may be produced by distillation of technical grade metal. Potassium (technical grade) may be packed under nitrogen. Argon should be used for packing high purity metal. Metal is shipped in stainless steel or carbon containers. In small quantities potassium is transported in glass or metal ampules.

Reactions

Potassium reacts with oxygen or air forming three oxides: potassium monoxide, K2O; potassium peroxide, K2O2; and potassium superoxide, KO2. The nature of the product depends on oxygen supply. In limited supply of oxygen potassium monoxide is formed, while in excess oxygen, superoxide is obtained:
4K + O2 → 2K2O
2K + O2 → K2O2
K + O2 → KO2
Potassium reacts violently with water, forming potassium hydroxide:
2K + 2H2O → 2KOH + H2
Potassium reacts with hydrogen at about 350ºC to form potassium hydride:
2K + H2 → 2KH
Reactions with halogens, fluorine, chlorine and bromine occur with explosive violence. Thus, in contact with liquid bromine it explodes forming potassium bromide:
2K + Br2 → 2KBr
Potassium ignites in iodine vapor forming potassium iodide.
Violent reactions can occur with many metal halides. For example, with zinc halides or iron halides, single replacement reactions take place. Such potassium-metal halide mixtures can react violently when subjected to mechanical shock.
At ordinary temperatures, potassium does not combine with nitrogen but with an electric charge, potassium azide is formed.
Reaction with carbon (graphite) at above 400ºC produces a series of carbides, such as KC4, KC8, and KC24. With carbon monoxide, an unstable explosive carbonyl forms:
K + CO → KCO
Potassium reduces carbon dioxide to carbon, carbon monoxide and potassium carbonate:
6K + 5CO2 → CO + C + 3K2CO3
Potassium reacts with ammonia gas to form potassium amide with liberation of hydrogen:
2K + 2NH3 → 2KNH2 + H2
Reactions with phosphorus, arsenic and antimony form phosphide, arsenide, and antimonide of potassium, respectively:
K + As → K3As
Reaction with sulfur forms three sulfides. When reactants are in molten state, the product is K2S, but in liquid ammonia K2S2 and KS2 are the main products.
Potassium reacts explosively with sulfuric acid, forming potassium sulfate with evolution of hydrogen:
K + H2SO4 → K2SO4 + H2
Potassium liberates hydrogen from ethanol forming potassium ethoxide:
2K + 2C2H5OH → 2C2H5OK + H2
Reaction with potassium nitrate yields potassium monoxide and nitrogen:
10K + 2KNO3 → 6K2O + N2

Hazard

Potassium metal can be dangerous to handle if proper precautions are not taken. Many of its reactions at ordinary temperatures can proceed to explosive violence (see Reactions). Also, it liberates flammable hydrogen gas when combined with water, acids, and alcohols.

Description

Potassium has atomic number 19 and the chemical symbol K, which is derived from its Latin name kalium . Potassium was first isolated from potash, which is potassium carbonate (K₂CO₃). Potassium occurs in nature only in the form of its ion (K⁺) either dissolved in the ocean or coordinated in minerals because elemental potassium reacts violently with water . Potassium ions are essential for the human body and are also present in plants. The major use of K⁺ can be found in fertilisers, which contains a variety of potassium salts such as potassium chloride (KCl), potassium sulfate (K₂SO₄) and potassium nitrate (KNO₃).

Chemical Properties

Potassium is a soft silvery metal, tarnishing upon exposure to air.

Physical properties

Elemental potassium is a soft, butter-like silvery metal whose cut surface oxidizes in dryair to form a dark gray potassium superoxide (KO₂) coating. KO₂ is an unusual compound,in that it reacts with both water and carbon dioxide to produce oxygen gas. It appears morelike a hard wax than a metal. Its density (specific gravity) is 0.862 g/cm³, its melting point is63.25°C, and its boiling point is 760°C. It has an oxidation state of +1 and reacts explosivelywith room temperature air or water to form potassium hydroxide as follows: 2K + 2 H₂O→ 2KOH + H₂. This is an endothermic reaction, which means the heat generated is greatenough to ignite the liberated hydrogen gas. Potassium metal must be stored in a non-oxygen,non-aqueous environment such as kerosene or naphtha.

Isotopes

A total of 18 isotopes of potassium have been discovered so far. Just two ofthem are stable: K-39 makes up 93.2581% of potassium found in the Earth’s crust, andK-41 makes up 6.7301% of the remainder of potassium found on Earth. All the other16 potassium isotopes are unstable and radioactive with relatively short half-lives, and asthey decay, they produce beta particles. The exception is K-40, which has a half-life of1.25×10⁹ years.

Origin of Name

Its symbol “K” is derived from the Latin word for alkali, kalium, but it is commonly called “potash” in English.

Occurrence

Potassium is the eighth most abundant element in the Earth’s crust, which contains about2.6% potassium, but not in natural elemental form. Potassium is slightly less abundant thansodium. It is found in almost all solids on Earth, in soil, and in seawater, which contains 380ppm of potassium in solution. Some of the potassium ores are sylvite, carnallite, and polyhalite. Ore deposits are found in New Mexico, California, Salt Lake in Utah, Germany, Russia,and Israel. Potassium metal is produced commercially by two processes. One is thermochemical distillation, which uses hot vapors of gaseous NaCl (sodium chloride) and KCl (potassiumchloride); the potassium is cooled and drained off as molten potassium, and the sodium chloride is discharged as a slag. The other procedure is an electrolytic process similar to that used toproduce lithium and sodium, with the exception that molten potassium chloride (which meltsat about 770°C) is used to produce potassium metal at the cathode.

Characteristics

Because its outer valence electrons are at a greater distance from its nuclei, potassium ismore reactive than sodium or lithium. Even so, potassium and sodium are very similar in theirchemical reactions. Due to potassium’s high reactivity, it combines with many elements, particularly nonmetals. Like the other alkali metals in group 1, potassium is highly alkaline (caustic) with a relatively high pH value. When given the flame test, it produces a violet color.

History

Discovered in 1807 by Davy, who obtained it from caustic potash (KOH); this was the first metal isolated by electrolysis. The metal is the seventh most abundant and makes up about 2.4% by weight of the Earth’s crust. Most potassium minerals are insoluble and the metal is obtained from them only with great difficulty. Certain minerals, however, such as sylvite, carnallite, langbeinite, and polyhalite are found in ancient lake and sea beds and form rather extensive deposits from which potassium and its salts can readily be obtained. Potash is mined in Germany, New Mexico, California, Utah, and elsewhere. Large deposits of potash, found at a depth of some 1000 m in Saskatchewan, promise to be important in coming years. Potassium is also found in the ocean, but is present only in relatively small amounts compared to sodium. The greatest demand for potash has been in its use for fertilizers. Potassium is an essential constituent for plant growth and it is found in most soils. Potassium is never found free in nature, but is obtained by electrolysis of the hydroxide, much in the same manner as prepared by Davy. Thermal methods also are commonly used to produce potassium (such as by reduction of potassium compounds with CaC2, C, Si, or Na). It is one of the most reactive and electropositive of metals. Except for lithium, it is the lightest known metal. It is soft, easily cut with a knife, and is silvery in appearance immediately after a fresh surface is exposed. It rapidly oxidizes in air and should be preserved in a mineral oil. As with other metals of the alkali group, it decomposes in water with the evolution of hydrogen. It catches fire spontaneously on water. Potassium and its salts impart a violet color to flames. Twenty-one isotopes, one of which is an isomer, of potassium are known. Ordinary potassium is composed of three isotopes, one of which is 40^K (0.0117%), a radioactive isotope with a half-life of 1.26 × 109 years. The radioactivity presents no appreciable hazard. An alloy of sodium and potassium (NaK) is used as a heat-transfer medium. Many potassium salts are of utmost importance, including the hydroxide, nitrate, carbonate, chloride, chlorate, bromide, iodide, cyanide, sulfate, chromate, and dichromate. Metallic potassium is available commercially for about $1200/ kg (98% purity) or $75/g (99.95% purity).

Uses

Some of the most common compounds in 19th century photography were made with this silvery metallic element discovered by Sir Humphrey Davy in 1807. There is not enough room in this work to list all of these compounds, but the following represent a reasonable sampling.

Uses

Liquid potassium, when mixed with liquid sodium (NaK), is an alloy used as a heatexchange substance to cool nuclear reactors. Potassium is an important reagent (something that is used in chemical reactions to analyze other substances) that forms many compounds used in chemical and industrial laboratories. It is used to manufacture both hard and soft soaps, as a bleaching agent, and where a highly caustic chemical is required. Potassium is essential to all living organisms. It is a trace element required for a healthy diet and is found in many foods. One natural source is bananas.

Uses

Potassium is used in the manufacture ofmany reactive potassium salts, in organicsynthesis, and as a heat exchange fluid whenalloyed with sodium.

Uses

In synthesis of inorganic potassium Compounds; in organic syntheses involving condensation, dehalogenation, reduction, and polymerization reactions. As heat transfer medium together with sodium: Chem. Eng. News 33, 648 (1955). Radioactive decay of 40K to 40Ar used as tool for geological dating.

Preparation

Potassium metal is not produced commercially by a fused salt electrolysis of the chloride —as is sodium—for several reasons: the metal is too soluble in the molten chloride to separate and float on top of the bath; potassium metal vapors may also issue from the molten bath, thus creating hazardous conditions; and potassium superoxide may form in the cell and react explosively with potassium metal. Consequently, the established method of preparing potassium metal commercially involves the reduction of molten potassium chloride by metallic sodium at elevated temperatures (850°C). Molten potassium chloride is fed into the midpoint of a steel vessel provided with a fractionating tower packed with stainless steel rings. Sodium is vaporized at the bottom and rises countercurrent to the molten potassium chloride with which it reacts according to the equilibrium expression.
Although the left-hand side of the equation is favored thermodynamically, the escape of the potassium vapors causes the reaction to proceed very efficiently to the right. The potassium vapors are condensed and the product normally contains sodium metal as the only major impurity up to about 1 % by weight. This product is sometimes purified by fractionating it in a 38 ft high 316 stainless steel tower equipped with a reflux return reservoir. The condensate is potassium metal of 99.99 % purity.

Production Methods

Potassium superoxide (KO2) can create oxygen from water vapor (H2O) and carbon dioxide (CO2) and is used in respiratory equipment and is produced by burning potassium metal in dry air.

Definition

potassium: Symbol K. A soft silverymetallic element belonging to group1 (formerly IA) of the periodic table(see alkali metals); a.n. 19; r.a.m.39.098; r.d. 0.86; m.p. 63.7°C; b.p.774°C. The element occurs in seawaterand in a number of minerals,such as sylvite (KCl), carnallite(KCl·MgCl2·6H2O), and kainite(MgSO4·KCl·3H2O). It is obtained byelectrolysis. The metal has few usesbut potassium salts are used for awide range of applications. Potassiumis an essential element for livingorganisms. The potassium ion,K+, is the most abundant cation inplant tissues, being absorbed throughthe roots and being used in suchprocesses as protein synthesis. In animalsthe passage of potassium andsodium ions across the nerve-cellmembrane is responsible for thechanges of electrical potential thataccompany the transmission of impulses.Chemically, it is highly reactive,resembling sodium in itsbehaviour and compounds. It alsoforms an orange-coloured superoxide,KO₂, which contains the O2^- ion.Potassium was discovered by SirHumphry Davy in 1807.

General Description

Potassium is potassium mixed with some other metal, usually sodium. Potassium is a liquid under normal conditions. Potassium reacts vigorously with water to form potassium hydroxide, a corrosive material and hydrogen, a flammable gas. The heat from this reaction may be sufficient to ignite the hydrogen. Potassium alloy may ignite spontaneously in contact with air. Once ignited, potassium burns quite violently. Potassium is used as a heat exchange fluid.

Air & Water Reactions

Reacts vigorously with oxygen. Reacts vigorously with water even at less than 100°C [Merck, 11th ed., 1989]. Water (caustic solution, H2) The oxidation of potassium in air is so rapid that the heat generated by the reaction melts and ignites the metal. This is particularly the case when pressure is applied at ordinary temperatures [Sidgwick 1. 1950]. Potassium burns in moist air at room temperature [Mellor 2:468. 1946-47]. The higher oxides of potassium, formed in air, react explosively with pure potassium, sodium, sodium-potassium alloys, and organic matter [Mellor 2, Supp. 3:1559. 1963].

Reactivity Profile

Boron trifluoride reacts with incandescence when heated with alkali metals or alkaline earth metals except magnesium [Merck 11th ed. 1989]. Maleic anhydride decomposes explosively in the presence of alkali metals . Sodium peroxide oxidizes antimony, arsenic, copper, potassium, tin, and zinc with incandescence . Alkali metal hydroxides, acids, anhydrous chlorides of iron, tin, and aluminum, pure oxides of iron and aluminum, and metallic potassium are some of the catalysts that may cause ethylene oxide to rearrange and polymerize, liberating heat . Explosions occur, although infrequently, from the combination of ethylene oxide and alcohols or mercaptans [Chem. Eng. News 20:1318. 1942]. A mixture of potassium and any of the following metallic halides produces a strong explosion on impact: aluminum chloride, aluminum fluoride, ammonium fluorocuprate, antimony tribromide, antimony trichloride, antimony triiodide, cadmium bromide, cadmium chloride, cadmium iodide, chromium tetrachloride, cupric bromide, cupric chloride, cuprous bromide cuprous chloride, cuprous iodide, manganese chloride, mercuric bromide, mercuric chloride, mercuric fluoride, mercuric iodide, mercurous chloride, nickel bromide, nickel chloride, nickel iodide, silicon tetrachloride, silver fluoride, stannic chloride, stannic iodide (with silver), stannous chloride, sulfur dibromide, thallous bromide, vanadium pentachloride, zinc bromide, zinc chloride, and zinc iodide [Mellor 2, Supp. 3:1571. 1963]. A mixture of potassium and any of the following compounds produces a weak explosion on impact: ammonium bromide, ammonium iodide, cadmium fluoride, chromium trifluoride, manganous bromide, manganous iodide, nickel fluoride, potassium chlorocuprate, silver chloride, silver iodide, strontium iodide, thallous chloride, and zinc fluoride [Mellor 2, Supp. 3:1571. 1963]. A mixture of potassium and any of the following compounds may explode on impact: boric acid, copper oxychloride, lead oxychloride, lead peroxide, lead sulfate, silver iodate, sodium iodate, and vanadium oxychloride [Mellor 2, Supp. 3:1571. 1963]. A mixture of potassium with any of the following compounds produces a very violent explosion on impact: boron tribromide, carbon tetrachloride, cobaltous bromide, cobaltous chloride, ferric bromide, ferric chloride, ferrous bromide, ferrous chloride, ferrous iodide, phosphorus pentachloride, phosphorus tribromide, and sulfur dichloride [Mellor 2, Supp. 3:1571. 1963]. Mixture of solid potassium and carbon dioxide(as dry ice) explodes when subjected to shock [Mellor 2, Supp. 3:1568. 1963]. Potassium and its alloys form explosive mixtures with chlorinated hydrocarbons [Chem. Eng. News 26:2604. 1948]. Ethylene oxide is dangerously reactive with metallic potassium [Chemical Safety Data Sheet SD-38:11. 1951]. Potassium in contact with the following oxides causes an explosive reaction: potassium ozonide, potassium peroxide, or potassium superoxide [Mellor 2, Supp. 3:1577. 1963].

Hazard

Elemental potassium as a metal is not found in its pure form in nature, but is derived fromits numerous compounds. The metal is very dangerous to handle. It can ignite while you areholding it with your hands or as you cut it. The metal must be stored in an inert gas atmosphereor in oil. Potassium fires cannot be extinguished with water—it only makes matters worsebecause it results in the formation of potassium hydroxide and hydrogen gas with enough heatto ignite the hydrogen. Dry chemicals such as soda ash, graphite, or dry sand can be used.
A particular hazard, which has been with humans since the beginning of time, is theradioactive isotope potassium-40 (K-40). Less than 1% of all potassium atoms on Earth arein the form of this radioactive isotope. It has a half-life of 1.25 billion years. Its decay process ends with the formation of the noble gas argon, which can then be analyzed to determine theage of rocks. This system (K-40 → argon) has been used to establish that the oldest rocks onEarth were formed about 3.8 billion years ago. Every living thing needs some potassium inits diet, including humans, who cannot escape this source of radiation, given that the humanbody cannot distinguish the radioactive potassium from the nonradioactive form. Along withcosmic rays and other naturally radioactive elements in the Earth’s crust, potassium-40 contributesto the normal lifetime accumulation of radiation. It makes up almost one-fourth ofthe total radiation the human body receives during a normal life span.

Health Hazard

Potassium reacts with the moisture on skin and other tissues to form highly corrosive potassium hydroxide. Contact of metallic potassium with the skin, eyes, or mucous membranes causes severe burns; thermal burns may also occur due to ignition of the metal and liberated hydrogen.

Flammability and Explosibility

Potassium metal may ignite spontaneously on contact with air at room temperature. Potassium reacts explosively with water to form potassium hydroxide; the heat liberated generally ignites the hydrogen formed and can initiate the combustion of potassium metal itself. Potassium fires must be extinguished with a class D dry chemical extinguisher or by the use of sand, ground limestone, dry clay or graphite, or "Met-L-X" type solids. Water or CO2, extinguishers must never be used on potassium fires.

reaction suitability

reagent type: reductant

Agricultural Uses

Since the beginning of the 19^th century, potassium has been recognized as an essential element and a major nutrient for plant growth, needed in large quantities. The exact function of potassium is not fully understood. Potassium makes plants more resistant to fimgal diseases and insect attacks. It is good for healthy root development and crop quality. For instance, potassium improves the (a) texture, color and combustibility of tobacco leaf, (b) sugar, starch and oil content in many plants, and (c) taste, size and keeping quality of fruits. Potato, tobacco and sugar use potassium, especially during their early growth stages. A small quantity of potassium is essential near young seedlings, while an excessive quantity causes salt damage.
The requirement of potassium varies in growing plants. Most seeds contain 0.1 to 10% potassium, which is sufficient for germination and early growth. The vegetative growth is characterized by a progressive increase in the absorption of inorganic elements like potassium. In tobacco, potassium is absorbed at the rate of 0.1 kglhalday from the 2l^st day of transplanting; a maximum uptake of 2 kg/ha/day occurs 49 days after transplanting. The minimum level of readily available potassium in the soil is around 175 kg/ha.
Potassium is present in the cell sap solution or plasma, and is almost fully extractable with water from plant tissues. It accumulates at the site of cell division, and helps in maintaining the physiological state of the swelling of plasma colloids which is necessary for all normal metabolic processes. It maintains the balance of anabolism, respiration and transpiration of a plant or leaf, and keeps the plant's water economy in equilibrium (in turn, reducing the plant's tendency to wilt.)
Potassium has a very important role to play in plant energy metabolism. Its liberal use helps to harden the supporting tissues which, in turn, improves the keeping qualities of fruits, and consequently leads to a stronger structure.
Potassium does not become a part of the plant structure as P, S, Ca and Mg do. But it helps in carbon dioxide assimilation, translocation of proteins and sugars, enzyme activity, cell division, reduction of nitrates and fat synthesis. The influence of potassium in these activities is now well established.
The levels of potassium and nitrogen are closely related in most plants. Nitrogen stimulates the rapid growth of soft tissues, whereas potassium promotes the growth of soft tissues. If sufficient potassium is unavailable, nitrogen level increases in the outer leaves of cabbage and in the upper stems and leaves of tomato. In the sheath tissue of sugar cane, the relationship of potassium to nitrogen depends on their respective concentrations.
Ammonium has a greater depressing effect on potassium in soil-grown plants than in solution, because ammonium interferes with the diffusion of potassium from the clay lattice. Potassium influences the uptake of the two forms of nitrogen. The relative presence of K, Ca and Mg influences the concentration of each individual cation within the plant. In this, potassium seems to be the most active. In plants, magnesium has a greater depressing effect on the content of potassium than that of calcium.
Because potassium ions (K+)an d sodium ions (Na+) are similar in size and chemical properties, sodium may replace potassium in several essential roles. However, potassium is an essential element, whereas sodium is not. Therefore, use of sodium may compensate for the potassium shortage to some extent, but sodium will not produce healthy plants in a situation when potassium deficiency is large.
There is a close relationship between carbohydrates and the potassium level. When soil potassium concentration is insufficient for optimum growth, it is commonly transported from more mature tissues to the meristems, so that older leaves exhibit early deficiency symptoms. Chlorosis appears first around the edges and tips of the leaves, and then spreads to the mid rib, followed finally by necrosis.
In many crops, potassium deficiency is characterized by a contrast between chlorosis, necrosis and healthy green areas of leaves. In the advanced stages of potassium starvation in corn, leaf edges become necrotic, the tissue disintegrates, and the leaf gets a ragged appearance. This condition is called leaf scorch.
Potassium deficiency in alfalfa is seen as white spots on the leaf edges, whereas chlorosis and necrosis of leaf edges are observed in other grasses. Potassium deficiency can also occur among young upper leaves in some high-yielding , fast-maturing crops like cotton and wheat. Insufficient potassium weakens the straw in grain crops, causes lodging in small grains and stalk breakage in corn and sorghum. Potassium deficiencies greatly reduce crop yield. A phenomenon in which deficiency symptoms are not visible is called hidden hunger. Potassium stress increases the degree of crop damage by bacterial and fungal diseases, insect and mite infestation, and nematode and virus infection. Lack of potassium in wetland rice greatly increases the sensitivity of foliar diseases such as stem rot, sheath blight and brown leaf spot.
Soil humus is a major source of sulphur, but not of potassium. Potassium ion is a highly soluble cation in solution, but it moves slowly in soils (unlike sulphur which is soluble and a readily mobile sulphate ion). Diffusion and mass flow of potassium to plant roots account for a large portion of absorbed potassium. In decaying humus, the potassium ion is fust leached into the soil solution and then to cation exchange sites on the humus and clay particles. A non-decomposed organic mass added to the soil replaces large amounts of potassium which flows with the water to the roots. In plant cells, potassium is the most abundant metal cation. On decomposition, fresh plant residues give all the potassium the plant needs for growth as a mobile soluble ion. Soluble potassium can be immobilized into the bodies of microbes, lost in leaching waters, or held between layers of hydrous mica and similar clays during drying. High yielding crop plants take potassium ions from a small reservoir of readily available potassium, namely the exchangeable source. For a good crop, at least 170 to 200 kg/ha potassium is considered essential. Soluble potassium may suffice if the soil is neutral or basic.
Using potassium fertilizers in excess, or too frequently, may result in an excess uptake of potassium by plants and in lowering their potassium-magnesium absorption. The effectiveness of the soil solutionpotassium for crop uptake is influenced by the presence of other cations, especially Na, Ca, Mg and Al. The absorption of potassium, in excess of that required for optimum growth, results in the accumulation of the nutrient without a corresponding increase in the growth, and is known as luxury consumption. The exchangeable or water-soluble potassium is converted by the potassium furation process to a form, not easily exchangeable from the adsorption complex, by a cation of a neutral salt solution.

Safety Profile

The toxicity of potassium compounds is almost always that of the anion, not of potassium. A dangerous fire hazard. Metallic potassium reacts with moisture to form potassium hydroxide and hydrogen. The reaction evolves much heat, causing the potassium to melt and spatter. The reaction also ignites the hydrogen, which burns, or if there is any confinement, may explode. It can ignite spontaneously in moist air. Store under mineral oil. Potassium metal wdl form the peroxide (K2O2) and the superoxide (KO3 or K2O4) at room temperature even when stored under mineral oil. These oxides can explode on contact with organic materials. Metal that has oxidized on storage under oil may explode violently when handled or cut. Oxide-coated potassium should be destroyed by burning. Danger: burning potassium is difficult to extinguish; dry powdered soda ash or graphte or special mixtures of dry chemical are recommended. A violent explosion hazard with the following materials under required conditions of temperature, pressure, and state of division: acetylene, air, moist air, alcohols (e.g., n-propanol through n-octanol, benzyl alcohol, cyclohexanol), AlBr3, ammonium nitrate + ammonium sulfate, ammonium chlorocuprate, NHdi, NH41, antimony halides, arsenic hahdes, AsH3 + NH3, Bi203, boric acid, BBr3, carbon disulfide (impact-sensitive), solid carbon dioxide, carbon monoxide, chlorinated hydrocarbons (e.g., chloroethane, dichloroethane, dchloromethane, trichloroethane, chloroform, pentachloro- ethane, carbon tetrachloride, tetrachloro- ethane), halocarbons (e.g., bromoform, dbromomethane, diiodomethane) , iodme (impact-sensitive), interhalogens (e.g., chlorine trifluoride, iodine bromide, iodine chloride, iodine pentafluoride, iodme trichloride), ClO, CrO3, Cu2OCl2, CuO, ethylene oxide, fluorine, graphite, graphte + air, graphite + K2O2, hydrogen iodide, H2O2, hydrogen chloride, hydrazine, Pb2OCl2, PbO2, PbSO4, maleic anhydride, metal halides (e.g., calcium bromide, iron(Ⅲ) bromide, iron(Ⅲ) chloride, iron(Ⅱ) chloride, iron(Ⅱ) bromide, iron(Ⅱ) iodide, cobalt(Ⅱ) chloride, chromium tetrachloride, silver fluoride, mercury(Ⅱ) bromide, mercury(Ⅱ) chloride, mercury(Ⅱ) fluoride, mercury(Ⅱ) iodide, copper0 chloride,copper(Ⅰ) iodde, copper(Ⅱ) bromide, copper(Ⅱ) chloride, ammonium tetrachlorocuprate, zinc chlorides, bromides, or ioddes, cadmium chlorides, bromides or iodides, aluminum fluorides, chlorides, or bromides, thalliump) bromide, tin chlorides, tin iodide, arsenic trichloride, arsenic triiodde, antimony tribromides, trichlorides or triiodides, bismuth tribromides, trichlorides, or triioddes, vanadiumo chloride, manganese(Ⅰ) chloride, nickel bromide, chloride, or iodide), metal oxides (e.g., lead peroxide, mercury(Ⅰ) oxide, MoO3, nitric acid, nitrogen-containing explosives (e.g., ammonium nitrate, picric acid, nitrobenzene), nonmetal halides (e.g., diselenium dichloride, seleninyl chloride, seleninyl bromide, sulfur dichloride, sulfur dibromide, phosphorus tribromide, phosphorus trichloride, phosgene, disulfur dichloride), nonmetal oxides (e.g., dichlorine oxide, dinitrogen tetraoxide, dinitrogen pentaoxide, NO2, P2O5), oxalyl dibromide, oxalyl dichloride, P2NF, peroxides, COCl2, PH3 + NH3, phosphorus, PCl5, PBr3, potassium chlorocuprate, potassium oxides (e.g., KO3, K2O2, KO2), selenium, SeOCl2, SiCl4, AglO3, NalO3, NH3 + NaNO2, Na2O2, SnI4 + S, SnO2, S, sulfuric acid, tellurium, thiophosphoryl fluoride, VOCl2, water. Other hazardous reactions may occur with carbon (e.g., soot, graphte, activated charcoal), dimethyl sulfoxide, ethylene oxide, chlorine, bromine vapor, hydrogen bromide, potassium iodide + magnesium bromide, chloride or iodide, maleic anhydride, mercury, copper(Ⅱ) oxide, mercury(Ⅱ) oxide, tin(Ⅳ) oxide, molybdenum(Ⅲ) oxide, bismuth trioxide, phosphorus trichloride, sulfur dioxide, chromium trioxide. toxic fumes of K2O. When heated to decomposition it emits

Potential Exposure

Used as a reagent and in sodiumpotassium alloys which are used as high-temperature heat transfer media.

Environmental Fate

Potassium metal in the environment will react with air, oxidizing the exposed surfaces, and reacts violently with water, yielding potassium hydroxide and hydrogen gas, which reacts with oxygen in air, producing flame.

storage

Safety glasses, impermeable gloves, and a fire-retardant laboratory coat should be worn at all times when working with potassium, and the metal should be handled under the surface of an inert liquid such as mineral oil, xylene, or toluene. Potassium should be used only in areas free of ignition sources and should be stored under mineral oil in tightly sealed metal containers under an inert gas such as argon. Potassium metal that has formed a yellow oxide coating should be disposed of immediately; do not attempt to cut such samples with a knife since the oxide coating may be explosive.

Shipping

UN2257Potassium, Hazard Class: 4.3; Labels: 4.3-Dangerous when wet material. UN1420 Potassium, metal alloys and metal alloys, liquid, Hazard Class: 4.3; Labels: 4.3-Dangerous when wet material. UN3089 Metal powder, flammable, n.o.s. Hazard Class: 4.2; Labels: 4.2-Spontaneously combustible material.

Toxicity evaluation

Potassium is a cofactor and activates a large variety of enzymes, including glycerol dehydrogenase, pyruvate kinase, L-threonine dehydrase, and ATPase. Its acute toxicity is primarily due to its action as an electrolyte. Excessive or diminished potassium levels can disrupt membrane excitability and influence muscle cell contractility and neuronal excitability.

Incompatibilities

Air contact causes spontaneous ignition. Violent reaction with water, forming heat, spattering, corrosive potassium hydroxide and explosive hydrogen. The heat from the reaction can ignite the hydrogen that is generated. A powerful reducing agent. Violent reaction with oxidizers, organic materials; carbon dioxide; heavy metal compounds; carbon tetrachloride; halogenated hydrocarbons; easily oxidized materials; and many other substances. Store under nitrogen, mineral oil, or kerosene. Oxidizes and forms unstable peroxides under storage conditions. Potassium metal containing an oxide coating is an extremely dangerous explosion hazard and should be removed by an expert and destroyed.

Waste Disposal

Excess potassium and waste material containing this substance should be placed in an appropriate container under an inert atmosphere, clearly labeled, and handled according to your institution's waste disposal guidelines. Experienced personnel can destroy small scraps of potassium by carefully adding t-butanol or nbutanol to a beaker containing the metal scraps covered in an inert solvent such as xylene or toluene.