Potassium: Difference between revisions

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{{About|the chemical element|the use of potassium as a medication|Potassium chloride (medical use)|the use of potassium in biology|Potassium in biology}}

{{Infobox potassium}}

'''Potassium''' is a [[chemical element]]; it has [[Symbol (chemistry)|symbol]] '''K''' (from [[Neo-Latin]] {{lang|la|kalium}}) and [[atomic number]]{{nbsp}}19. It is a silvery white metal that is soft enough to easily cut with a knife. Potassium metal reacts rapidly with atmospheric [[oxygen]] to form flaky white [[potassium peroxide]] in only seconds of exposure. It was first isolated from [[potash]], the ashes of plants, from which its name derives. In the [[periodic table]], potassium is one of the [[alkali metal]]s, all of which have a single [[valence electron]] in the outer electron shell, which is easily removed to create [[cation|an&nbsp;ion with a positive charge]] (which combines with [[anion]]s to form [[salts]]). In nature, potassium occurs only in ionic salts. Elemental potassium reacts vigorously with water, generating sufficient heat to ignite [[hydrogen]] emitted in the reaction, and burning with a [[lilac]]-[[flame color|colored flame]]. It is found dissolved in seawater (which is 0.04% potassium by weight), and occurs in many [[mineral]]s such as [[orthoclase]], a common constituent of [[granite]]s and other [[igneous rock]]s.

Potassium is chemically very similar to [[sodium]], the previous element in [[Group (periodic table)|group]] 1 of the periodic table. They have a similar first [[ionization energy]], which allows for each atom to give up its sole outer electron. It was first suggested in 1702 that they were distinct elements that combine with the same anions to make similar salts, which was demonstrated in 1807 when elemental potassium was first isolated via [[electrolysis]]. Naturally occurring potassium is composed of three [[isotope]]s, of which [[potassium-40|{{chem|40|K}}]] is [[radioactive]]. Traces of {{chem|40|K}} are found in all potassium, and it is the most common [[radioisotope]] in the human body.

Potassium ions are vital for the functioning of all living cells. The transfer of potassium ions across nerve cell membranes is necessary for normal nerve transmission; potassium deficiency and excess can each result in numerous signs and symptoms, including an abnormal heart rhythm and various [[electrocardiographic]] abnormalities. Fresh fruits and vegetables are good dietary sources of potassium. The body responds to the influx of dietary potassium, which raises [[serum (blood)|serum]] potassium levels, by shifting potassium from outside to inside cells and increasing potassium excretion by the kidneys.

Most industrial applications of potassium exploit the high [[solubility]] of its compounds in water, such as [[saltwater soap]]. Heavy crop production rapidly depletes the soil of potassium, and this can be remedied with agricultural fertilizers containing potassium, accounting for 95% of global potassium chemical production.

{{TOC limit}}

==Etymology==

The English name for the element ''potassium'' comes from the word ''[[potash]]'', which refers to an early method of extracting various potassium salts: placing in a ''pot'' the ''ash'' of burnt wood or tree leaves, adding water, heating, and evaporating the solution. When [[Humphry Davy]] first isolated the pure element using [[electrolysis]] in 1807, he named it ''potassium'', which he derived from the word ''potash''.

The symbol ''K'' stems from ''kali'', itself from the root word ''[[alkali]]'', which in turn comes from {{lang-ar|القَلْيَه}} ''al-qalyah'' 'plant ashes'. In 1797, the German chemist [[Martin Heinrich Klaproth|Martin Klaproth]] discovered "potash" in the minerals [[leucite]] and [[lepidolite]], and realized that "potash" was not a product of plant growth but actually contained a new element, which he proposed calling ''kali''. In 1807, Humphry Davy produced the element via electrolysis: in 1809, [[Ludwig Wilhelm Gilbert]] proposed the name ''Kalium'' for Davy's "potassium". In 1814, the Swedish chemist [[Jöns Jacob Berzelius|Berzelius]] advocated the name ''kalium'' for potassium, with the chemical symbol ''K''.

The English and French-speaking countries adopted the name ''Potassium'', which was favored by Davy and French chemists [[Joseph Louis Gay-Lussac]] and [[Louis Jacques Thénard]], whereas the other Germanic countries adopted Gilbert and Klaproth's name ''Kalium''. The "Gold Book" of the [[International Union of Pure and Applied Chemistry]] has designated the official chemical symbol as '''K'''.

==Properties==

===Physical===

Potassium is the second least dense metal after [[lithium]]. It is a soft solid with a low [[melting point]], and can be easily cut with a knife. Potassium is silvery in appearance, but it begins to tarnish toward gray immediately on exposure to air. In a [[flame test]], potassium and its compounds emit a [[Lilac (color)|lilac color]] with a peak emission wavelength of 766.5 nanometers.

Neutral potassium atoms have 19 electrons, one more than the configuration of the [[noble gas]] [[argon]]. Because of its low first [[ionization energy]] of 418.8{{nbsp}}kJ/mol, the potassium atom is much more likely to lose the last electron and acquire a positive charge, although negatively charged [[alkalide]] {{chem2|K−}} ions are not impossible. In contrast, the second ionization energy is very high (3052{{nbsp}}kJ/mol).

===Chemical===

Potassium reacts with oxygen, water, and carbon dioxide components in air. With oxygen it forms [[potassium peroxide]]. With water potassium forms [[potassium hydroxide]] (KOH). The reaction of potassium with water can be violently [[exothermic]], especially since the coproduced [[hydrogen]] gas can ignite. Because of this, potassium and the liquid sodium-potassium ([[NaK]]) alloy are potent [[desiccant]]s, although they are no longer used as such.

===Compounds===

[[Image:potassium-superoxide-unit-cell-3D-ionic.png|thumb|left|upright|Structure of solid potassium superoxide ({{chem2|KO2}}).]]

Four oxides of potassium are well studied: [[potassium oxide]] ({{chem2|K2O}}), potassium peroxide ({{chem2|K2O2}}), [[potassium superoxide]] ({{chem2|KO2}}) and [[potassium ozonide]] ({{chem2|KO3}}). The binary potassium-oxygen compounds react with water forming KOH.

KOH is a [[strong base]]. Illustrating its [[hydrophilic]] character, as much as 1.21 [[kilogram|kg]] of KOH can dissolve in a single liter of water. Anhydrous KOH is rarely encountered. KOH reacts readily with [[carbon dioxide]] ({{chem2|CO2}}) to produce [[potassium carbonate]] ({{chem2|K2CO3}}), and in principle could be used to remove traces of the gas from air. Like the closely related [[sodium hydroxide]], KOH reacts with [[fat]]s to produce [[soap]]s.

In general, potassium compounds are ionic and, owing to the high hydration energy of the {{chem2|K+}} ion, have excellent water solubility. The main species in water solution are the [[Metal aquo complex|aquo complexes]] {{chem2|[K(H2O)_{''n''}]+}} where ''n'' = 6 and 7.

[[Potassium heptafluorotantalate]] ({{chem2|K2[TaF7]}}) is an intermediate in the purification of [[tantalum]] from the otherwise persistent contaminant of [[niobium]].

[[Organopotassium compound]]s illustrate nonionic compounds of potassium. They feature highly [[Chemical polarity|polar]] [[covalent]] K–C bonds. Examples include [[benzyl potassium]] {{chem2|KCH2C6H5}}. Potassium [[Intercalation (chemistry)|intercalate]]s into [[graphite]] to give a variety of [[graphite intercalation compounds]], including {{chem2|KC8}}.

===Isotopes===

{{main|Isotopes of potassium}}

There are 25 known [[isotope]]s of potassium, three of which occur naturally: {{chem|39|K}} (93.3%), {{chem|40|K}} (0.0117%), and {{chem|41|K}} (6.7%) (by mole fraction). Naturally occurring [[potassium-40|{{chem|40|K}}]] has a [[half-life]] of {{val|1.250e9}} years. It decays to stable [[Argon|{{chem|40|Ar}}]] by [[electron capture]] or [[positron emission]] (11.2%) or to stable [[Calcium|{{chem|40|Ca}}]] by [[beta decay]] (88.8%). The decay of {{chem|40|K}} to {{chem|40|Ar}} is the basis of a common method for dating rocks. The conventional [[Potassium-argon dating|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 ({{chem|40|Ar}}) was quantitatively retained. [[Mineral]]s are dated by measurement of the concentration of potassium and the amount of radiogenic {{chem|40|Ar}} that has accumulated. The minerals best suited for dating include [[biotite]], [[muscovite]], [[metamorphic]] [[hornblende]], and volcanic [[feldspar]]; [[Petrography|whole rock]] samples from volcanic flows and shallow [[Igneous rock|instrusives]] can also be dated if they are unaltered. Apart from dating, potassium isotopes have been used as [[radioactive tracer|tracers]] in studies of [[weathering]] and for [[nutrient cycling]] studies because potassium is a [[macronutrient (ecology)|macronutrient]] required for [[life]] on Earth.

{{chem|40|K}} 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. {{chem|40|K}} is the radioisotope with the largest abundance [[Composition of the human body|in the human body]]. In healthy animals and people, {{chem|40|K}} represents the largest source of radioactivity, greater even than [[Carbon-14|{{chem|14|C}}]]. In a human body of 70 kg, about 4,400 nuclei of {{chem|40|K}} decay per second. The activity of natural potassium is 31 [[Becquerel|Bq]]/g.

==History==

===Potash===

Potash is primarily a mixture of potassium salts because plants have little or no sodium content, and the rest of a plant's major mineral content consists of calcium salts of relatively low solubility in water. While potash has been used since ancient times, its composition was not understood. [[Georg Ernst Stahl]] obtained experimental evidence that led him to suggest the fundamental difference of sodium and potassium salts in 1702, The exact chemical composition of potassium and sodium compounds, and the status as chemical element of potassium and sodium, was not known then, and thus [[Antoine Lavoisier]] did not include the alkali in his list of chemical elements in 1789. For a long time the only significant applications for potash were the production of glass, bleach, soap and [[gunpowder]] as potassium nitrate. Potassium soaps from animal fats and vegetable oils were especially prized because they tend to be more water-soluble and of softer texture, and are therefore known as soft soaps. The discovery by [[Justus Liebig]] in 1840 that potassium is a necessary element for plants and that most types of soil lack potassium caused a steep rise in demand for potassium salts. Wood-ash from fir trees was initially used as a potassium salt source for fertilizer, but, with the discovery in 1868 of mineral deposits containing [[potassium chloride]] near [[Staßfurt]], Germany, the production of potassium-containing fertilizers began at an industrial scale. Other potash deposits were discovered, and by the 1960s Canada became the dominant producer.

===Metal===

[[File:Sir Humphry Davy, Bt by Thomas Phillips.jpg|thumb|[[Humphry Davy|Sir Humphry Davy]] ]]

[[File:Potassium.JPG|thumb|upright|left|Pieces of potassium metal]]

Potassium ''metal'' was first isolated in 1807 by Humphry Davy, who derived it by electrolysis of molten [[caustic potash]] (KOH) with the newly discovered [[voltaic pile]]. Potassium was the first metal that was isolated by electrolysis. Later in the same year, Davy reported extraction of the metal [[sodium]] from a mineral derivative ([[caustic soda]], NaOH, or lye) rather than a plant salt, by a similar technique, demonstrating that the elements, and thus the salts, are different. Although the production of potassium and sodium metal should have shown that both are elements, it took some time before this view was universally accepted.

Because of the sensitivity of potassium to water and air, [[air-free technique]]s are normally employed for handling the element. It is unreactive toward nitrogen and saturated hydrocarbons such as mineral oil or [[kerosene]]. It readily dissolves in liquid [[ammonia]], up to 480 g per 1000 g of ammonia at 0{{nbsp}}°C. Depending on the concentration, the ammonia solutions are blue to yellow, and their electrical conductivity is similar to that of liquid metals. Potassium slowly reacts with ammonia to form [[Potassium amide|{{chem|KNH|2}}]], but this reaction is accelerated by minute amounts of transition metal salts. Because it can reduce the [[salts]] to the metal, potassium is often used as the reductant in the preparation of finely divided metals from their salts by the [[Rieke metal|Rieke method]]. Illustrative is the preparation of magnesium:

:{{chem2|[[Magnesium chloride|MgCl2]] + 2 K → Mg + 2 KCl}}

==Occurrence==

[[File:PotassiumFeldsparUSGOV.jpg|thumb|left|upright|Potassium in [[feldspar]]]]

Potassium is formed in [[supernova]]e by [[nucleosynthesis]] from lighter atoms. Potassium is principally created in Type II supernovae via an [[Supernova nucleosynthesis|explosive oxygen-burning process]]. These are nuclear [[Nuclear_Fusion|fusion]] reactions, not to be confused with chemical burning of potassium in oxygen. {{chem|40|K}} is also formed in {{Nowrap|[[s-process]]}} nucleosynthesis and the [[neon burning process]].

Potassium is the 20th most abundant element in the solar system and the 17th most abundant element by weight in the Earth. It makes up about 2.6% of the weight of the [[Earth's crust]] and is the seventh most abundant element in the crust. The potassium concentration in seawater is 0.39{{nbsp}}g/L (0.039 wt/v%), about one twenty-seventh the concentration of sodium.

===Geology===

Elemental potassium does not occur in nature because of its high reactivity. It reacts violently with water and also reacts with oxygen. [[Orthoclase]] (potassium feldspar) is a common rock-forming mineral. [[Granite]] for example contains 5% potassium, which is well above the average in the Earth's crust. [[Sylvite]] (KCl), [[carnallite]] ({{chem2|KCl*MgCl2*6H2O}}), [[kainite]] ({{chem2|MgSO4*KCl*3H2O}}) and [[langbeinite]] ({{chem2|MgSO4*K2SO4}}) are the minerals found in large [[evaporite]] deposits worldwide. The deposits often show layers starting with the least soluble at the bottom and the most soluble on top. Deposits of niter ([[potassium nitrate]]) are formed by decomposition of organic material in contact with atmosphere, mostly in caves; because of the good water solubility of niter the formation of larger deposits requires special environmental conditions.

==Commercial production==

===Mining===

[[File:Wintershall Monte Kali 12.jpg|thumb|[[Monte Kali (Heringen)|Monte Kali]], a potash mining and [[beneficiation]] waste heap in [[Hesse|Hesse, Germany]], consisting mostly of [[sodium chloride]].]]

Potassium salts such as [[carnallite]], [[langbeinite]], [[polyhalite]], and [[sylvite]] form extensive [[evaporite]] deposits in ancient lake bottoms and [[seabed]]s, making extraction of potassium salts in these environments commercially viable. The principal source of potassium – potash – is mined in [[Canada]], [[Russia]], [[Belarus]], [[Kazakhstan]], [[Germany]], [[Israel]], the U.S., [[Jordan]], and other places around the world. The first mined deposits were located near Staßfurt, Germany, but the deposits span from [[Great Britain]] over Germany into Poland. They are located in the [[Zechstein]] and were deposited in the Middle to Late [[Permian]]. The largest deposits ever found lie {{convert|1000|m|ft|abbr=off|sp=us}} below the surface of the Canadian province of [[Saskatchewan]]. The deposits are located in the [[Elk Point Group]] produced in the [[Middle Devonian]]. Saskatchewan, where several large mines have operated since the 1960s pioneered the technique of freezing of wet sands (the Blairmore formation) to drive mine shafts through them. The main potash mining company in Saskatchewan until its merge was the [[Potash Corporation of Saskatchewan]], now [[Nutrien]]. The water of the [[Dead Sea]] is used by Israel and Jordan as a source of potash, while the concentration in normal oceans is too low for commercial production at current prices.

===Chemical extraction===

Several methods are used to separate potassium salts from sodium and magnesium compounds. The most-used method is fractional precipitation using the solubility differences of the salts. Electrostatic separation of the ground salt mixture is also used in some mines. The resulting sodium and magnesium waste is either stored underground or piled up in [[slag heap]]s. Most of the mined potassium mineral ends up as potassium chloride after processing. The mineral industry refers to potassium chloride either as potash, muriate of potash, or simply MOP.

Pure potassium metal can be isolated by electrolysis of its [[potassium hydroxide|hydroxide]] in a process that has changed little since it was first used by Humphry Davy in 1807. Although the electrolysis process was developed and used in industrial scale in the 1920s, the thermal method by reacting sodium with [[potassium chloride]] in a chemical equilibrium reaction became the dominant method in the 1950s.

:Na + KCl → NaCl + K

The production of [[NaK|sodium potassium alloys]] is accomplished by changing the reaction time and the amount of sodium used in the reaction. The Griesheimer process employing the reaction of [[potassium fluoride]] with [[calcium carbide]] was also used to produce potassium.

:{{chem2|2 KF + CaC2 → 2 K + CaF2 + 2 C}}

[[reagent|Reagent-grade]] potassium metal costs about $10.00/[[pound (mass)|pound]] ($22/[[kg]]) in 2010 when purchased by the [[tonne]]. Lower purity metal is considerably cheaper. The market is volatile because long-term storage of the metal is difficult. It must be stored in a dry [[inert gas]] atmosphere or [[anhydrous]] [[mineral oil]] to prevent the formation of a surface layer of [[potassium superoxide]], a pressure-sensitive [[explosive]] that [[Detonation|detonates]] when scratched. The resulting explosion often starts a fire difficult to extinguish.

===Cation identification===

Potassium is now quantified by ionization techniques, but at one time it was quantitated by [[gravimetric analysis]].

Reagents used to precipitate potassium salts include [[sodium tetraphenylborate]], [[hexachloroplatinic acid]], and [[sodium cobaltinitrite]] into respectively [[potassium tetraphenylborate]], [[potassium hexachloroplatinate]], and [[potassium cobaltinitrite]].

The reaction with [[sodium cobaltinitrite]] is illustrative:

The potassium cobaltinitrite is obtained as a yellow solid.

==Commercial uses==

===Fertilizer===

Potassium ions are an essential component of [[plant]] nutrition and are found in most [[soil]] types. They are used as a [[fertilizer]] in [[agriculture]], [[horticulture]], and [[hydroponic]] culture in the form of [[potassium chloride|chloride]] (KCl), [[potassium sulfate|sulfate]] ({{chem2|K2SO4}}), or [[potassium nitrate|nitrate]] ({{chem2|KNO3}}), representing the 'K' [[labeling of fertilizer|in 'NPK']]. Agricultural fertilizers consume 95% of global potassium chemical production, and about 90% of this potassium is supplied as KCl. The potassium content of most plants ranges from 0.5% to 2% of the harvested weight of crops, conventionally expressed as amount of {{chem2|K2O}}. Modern high-[[Crop yield|yield]] agriculture depends upon fertilizers to replace the potassium lost at harvest. Most agricultural fertilizers contain potassium chloride, while potassium sulfate is used for chloride-sensitive crops or crops needing higher sulfur content. The sulfate is produced mostly by decomposition of the complex minerals [[kainite]] ({{chem2|MgSO4*KCl*3H2O}}) and [[langbeinite]] ({{chem2|MgSO4*K2SO4}}). Only a very few fertilizers contain potassium nitrate. In 2005, about 93% of world potassium production was consumed by the fertilizer industry. Furthermore, potassium can play a key role in nutrient cycling by controlling litter composition.

====Potassium citrate====

[[Potassium citrate]] is used to treat a [[kidney stone]] condition called [[renal tubular acidosis]].

====Potassium chloride====

{{see also|Potassium chloride (medical use)}}

Potassium, in the form of potassium chloride is used as a medication to treat and prevent [[low blood potassium]]. Low blood potassium may occur due to [[vomiting]], [[diarrhea]], or certain medications. It is given by [[intravenous infusion|slow injection into a vein]] or by mouth.

===Food additives===

Potassium sodium tartrate ({{chem2|KNaC4H4O6}}, [[Rochelle salt]]) is a main constituent of some varieties of [[baking powder]]; it is also used in the [[silvering]] of mirrors. [[Potassium bromate]] ({{chem2|KBrO3}}) is a strong oxidizer (E924), used to improve [[dough]] strength and rise height. [[Potassium bisulfite]] ({{chem2|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 [[leather]]s.

===Industrial===

Major potassium chemicals are potassium hydroxide, potassium carbonate, potassium sulfate, and potassium chloride. Megatons of these compounds are produced annually.

KOH is a strong base, which is used in industry to neutralize strong and weak [[acid]]s, to control [[pH]] and to manufacture potassium [[salt (chemistry)|salts]]. It is also used to [[saponification|saponify]] fats and [[oils]], in industrial cleaners, and in hydrolysis reactions, for example of [[esters]].

[[Potassium nitrate]] ({{chem2|KNO3}}) or saltpeter is obtained from natural sources such as [[guano]] and [[evaporites]] or manufactured via the [[Haber process]]; it is the [[oxidant]] in gunpowder ([[black powder]]) and an important agricultural fertilizer. [[Potassium cyanide]] (KCN) is used industrially to dissolve [[copper]] and precious metals, in particular [[silver]] and [[gold]], by forming [[complex (chemistry)|complexes]]. Its applications include [[gold mining]], [[electroplating]], and [[electroforming]] of these [[metal]]s; it is also used in [[organic synthesis]] to make [[nitriles]]. Potassium carbonate ({{chem2|K2CO3}} or potash) is used in the manufacture of glass, soap, color TV tubes, fluorescent lamps, textile dyes and pigments. [[Potassium permanganate]] ({{chem2|KMnO4}}) is an oxidizing, bleaching and purification substance and is used for production of [[saccharin]]. [[Potassium chlorate]] ({{chem2|KClO3}}) is added to matches and explosives. [[Potassium bromide]] (KBr) was formerly used as a sedative and in photography.

While [[potassium chromate]] ({{chem2|K2CrO4}}) is used in the manufacture of a host of different commercial products such as [[ink]]s, [[dye]]s, wood [[stain]]s (by reacting with the [[tannic acid]] in wood), [[explosive]]s, [[fireworks]], [[fly paper]], and [[safety match]]es, as well as in the tanning of leather, all of these uses are due to the chemistry of the [[chromate ion]] rather than to that of the potassium ion.

====Niche uses====

There are thousands of uses of various potassium compounds. One example is [[potassium superoxide]], {{chem2|KO2}}, an orange solid that acts as a portable source of oxygen and a carbon dioxide absorber. It is widely used in [[Rebreather#Rebreathers using an absorbent that releases oxygen|respiration systems]] in mines, submarines and spacecraft as it takes less volume than the gaseous oxygen.

:{{chem2|4 KO2 + 2 CO2 → 2 K2CO3 + 3 O2}}

Another example is [[potassium cobaltinitrite]], {{chem2|K3[Co(NO2)6]}}, which is used as artist's pigment under the name of [[Aureolin]] or Cobalt Yellow.

The stable isotopes of potassium can be [[Laser cooling|laser cooled]] and used to probe fundamental and [[Quantum technology|technological]] problems in [[Quantum mechanics|quantum physics]]. The two [[boson]]ic isotopes possess convenient [[Feshbach resonance]]s to enable studies requiring tunable interactions, while {{chem|40|K}} is one of only two stable [[fermion]]s amongst the alkali metals.

====Laboratory uses====

An [[alloy]] of sodium and potassium, [[NaK]] is a liquid used as a heat-transfer medium and a [[desiccant]] for producing [[air-free technique|dry and air-free solvents]]. It can also be used in [[reactive distillation]]. The ternary alloy of 12% Na, 47% K and 41% Cs has the lowest melting point of −78{{nbsp}}°C of any metallic compound.

Metallic potassium is used in several types of [[magnetometer]]s.

==Biological role==

{{Main|Potassium in biology}}

Potassium is the eighth or ninth most common element by mass (0.2%) in the human body, so that a 60{{nbsp}}kg adult contains a total of about 120{{nbsp}}g of potassium. The body has about as much potassium as sulfur and chlorine, and only calcium and phosphorus are more abundant (with the exception of the ubiquitous [[CHON]] elements).

===Biochemical function===

Potassium levels influence multiple physiological processes, including

*local cortical monoaminergic norepinephrine, serotonin, and dopamine levels, and through them, sleep/wake balance, and spontaneous activity.

===Homeostasis===

Potassium homeostasis denotes the maintenance of the total body potassium content, plasma potassium level, and the ratio of the intracellular to extracellular potassium concentrations within narrow limits, in the face of pulsatile intake (meals), obligatory renal excretion, and shifts between intracellular and extracellular compartments.

====Plasma levels====

Plasma potassium is normally kept at 3.5 to 5.5 millimoles (mmol) [or milliequivalents (mEq)] per liter by multiple mechanisms. Levels outside this range are associated with an increasing rate of death from multiple causes, and some cardiac, kidney, and lung diseases progress more rapidly if serum potassium levels are not maintained within the normal range.

An average meal of 40–50{{nbsp}}mmol presents the body with more potassium than is present in all plasma (20–25{{nbsp}}mmol). This surge causes the plasma potassium to rise up to 10% before clearance by renal and extrarenal mechanisms.

[[Hypokalemia]], a deficiency of potassium in the plasma, can be fatal if severe. Common causes are increased gastrointestinal loss ([[vomiting]], [[diarrhea]]), and increased renal loss ([[polyuria|diuresis]]). Deficiency symptoms include muscle weakness, [[paralytic ileus]], ECG abnormalities, decreased reflex response; and in severe cases, respiratory paralysis, [[alkalosis]], and [[cardiac arrhythmia]].

====Control mechanisms====

Potassium content in the plasma is tightly controlled by four basic mechanisms, which have various names and classifications. These are:

* The ion transport system moves potassium across the cell membrane using two mechanisms. One is active and pumps sodium out of, and potassium into, the cell. The other is passive and allows potassium to leak out of the cell. Potassium and sodium cations influence fluid distribution between intracellular and extracellular compartments by [[osmotic]] forces. The movement of potassium and sodium through the cell membrane is mediated by the [[Na⁺/K⁺-ATPase]] pump. This [[Ion transporter|ion pump]] uses [[Adenosine triphosphate|ATP]] to pump three sodium ions out of the cell and two potassium ions into the cell, creating an electrochemical gradient and electromotive force across the cell membrane. The highly selective [[potassium ion channels]] (which are [[tetramer]]s) are crucial for [[Hyperpolarization (biology)|hyperpolarization]] inside [[neuron]]s after an action potential is triggered, to cite one example. The most recently discovered potassium ion channel is KirBac3.1, which makes a total of five potassium ion channels (KcsA, KirBac1.1, KirBac3.1, KvAP, and MthK) with a determined structure. All five are from [[prokaryotic]] species.

====Renal filtration, reabsorption, and excretion====

Renal handling of potassium is closely connected to sodium handling. Potassium is the major cation (positive ion) inside animal cells (150{{nbsp}}mmol/L, 4.8{{nbsp}}g/L), while sodium is the major cation of [[extracellular fluid]] (150{{nbsp}}mmol/L, 3.345{{nbsp}}g/L). In the kidneys, about 180{{nbsp}}liters of plasma is filtered through the [[Glomerulus (kidney)|glomeruli]] and into the [[renal tubules]] per day. This filtering involves about 600{{nbsp}}mg of sodium and 33{{nbsp}}mg of potassium. Since only 1–10{{nbsp}}mg of sodium and 1–4{{nbsp}}mg of potassium are likely to be replaced by diet, renal filtering must efficiently reabsorb the remainder from the plasma.

Sodium is reabsorbed to maintain extracellular volume, osmotic pressure, and serum sodium concentration within narrow limits. Potassium is reabsorbed to maintain serum potassium concentration within narrow limits. [[Sodium pump]]s in the renal tubules operate to reabsorb sodium. Potassium must be conserved, but because the amount of potassium in the blood plasma is very small and the pool of potassium in the cells is about 30 times as large, the situation is not so critical for potassium. Since potassium is moved passively in counter flow to sodium in response to an apparent (but not actual) [[Donnan equilibrium]], 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 excreted twice and reabsorbed three times before the urine reaches the collecting tubules. At that point, urine usually has about the same potassium concentration as plasma. At the end of the processing, potassium is secreted one more time if the serum levels are too high.

With no potassium intake, it is excreted at about 200{{nbsp}}mg per day until, in about a week, potassium in the serum declines to a mildly deficient level of 3.0–3.5{{nbsp}}mmol/L. If potassium is still withheld, the concentration continues to fall until a severe deficiency causes eventual death.

The potassium moves passively through pores in the cell membrane. When ions move through [[ion transporter]]s (pumps) there is a gate in the pumps on both sides of the cell membrane and only one gate can be open at once. As a result, approximately 100 ions are forced through per second. [[Ion channel]]s have only one gate, and there only one kind of ion can stream through, at 10 million to 100 million ions per second. Calcium is required to open the pores, although calcium may work in reverse by blocking at least one of the pores. Carbonyl groups inside the pore on the amino acids mimic the water hydration that takes place in water solution by the nature of the electrostatic charges on four carbonyl groups inside the pore.

====Dietary recommendations====

The U.S. [[National Academy of Medicine]] (NAM), on behalf of both the U.S. and Canada, sets [[Dietary Reference Intake]]s, including Estimated Average Requirements (EARs) and Recommended Dietary Allowances (RDAs), or [[Adequate Intake]]s (AIs) for when there is not sufficient information to set EARs and RDAs.

For both males and females under 9 years of age, the AIs for potassium are: 400{{nbsp}}mg of potassium for 0 to 6-month-old infants, 860{{nbsp}}mg of potassium for 7 to 12-month-old infants, 2,000{{nbsp}}mg of potassium for 1 to 3-year-old children, and 2,300{{nbsp}}mg of potassium for 4 to 8-year-old children.

For males 9 years of age and older, the AIs for potassium are: 2,500{{nbsp}}mg of potassium for 9 to 13-year-old males, 3,000{{nbsp}}mg of potassium for 14 to 18-year-old males, and 3,400{{nbsp}}mg for males that are 19 years of age and older.

For females 9 years of age and older, the AIs for potassium are: 2,300{{nbsp}}mg of potassium for 9 to 18-year-old females, and 2,600{{nbsp}}mg of potassium for females that are 19 years of age and older.

For pregnant and lactating females, the AIs for potassium are: 2,600{{nbsp}}mg of potassium for 14 to 18-year-old pregnant females, 2,900{{nbsp}}mg for pregnant females that are 19 years of age and older; furthermore, 2,500{{nbsp}}mg of potassium for 14 to 18-year-old lactating females, and 2,800{{nbsp}}mg for lactating females that are 19 years of age and older. As for safety, the NAM also sets [[tolerable upper intake level]]s (ULs) for vitamins and minerals, but for potassium the evidence was insufficient, so no UL was established.

As of 2004, most Americans adults consume less than 3,000{{nbsp}}mg.

Likewise, in the European Union, in particular in Germany, and Italy, insufficient potassium intake is somewhat common. The [[National Health Service|British National Health Service]] recommends a similar intake, saying that adults need 3,500{{nbsp}}mg per day and that excess amounts may cause health problems such as stomach pain and [[diarrhea]].

In 2019, the [[National Academies of Sciences, Engineering, and Medicine]] revised the Adequate Intake for potassium to 2,600&nbsp;mg/day for females 19 years of age and older who are not pregnant or lactating, and 3,400&nbsp;mg/day for males 19 years of age and older.

====Food sources====

Potassium is present in all fruits, vegetables, meat and fish. Foods with high potassium concentrations include [[Yam (vegetable)|yam]], [[parsley]], dried [[apricot]]s, [[milk]], [[chocolate]], all [[nut (fruit)|nuts]] (especially [[almond]]s and [[pistachio]]s), [[potato]]es, [[bamboo shoot]]s, [[banana]]s, [[avocado]]s, [[coconut water]], [[soybean]]s, and [[bran]].

The  [[United States Department of Agriculture]] also lists [[tomato paste]], [[orange juice]], [[beet greens]], [[white beans]], [[Cooking banana|plantains]], and many other dietary sources of potassium, ranked in descending order according to potassium content. A day's worth of potassium is in 5 plantains or 11 bananas.

====Deficient intake====

Diets low in potassium can lead to [[hypertension]] and [[hypokalemia]].

====Supplementation====

Supplements of potassium are most widely used in conjunction with [[diuretic]]s that block reabsorption of sodium and water upstream from the [[distal tubule]] ([[thiazide]]s and [[loop diuretics]]), because this promotes increased distal tubular potassium secretion, with resultant increased potassium excretion. A variety of prescription and over-the counter supplements are available. Potassium chloride may be dissolved in water, but the salty/bitter taste makes liquid supplements unpalatable. Typical doses range from 10{{nbsp}}mmol (400{{nbsp}}mg), to 20{{nbsp}}mmol (800{{nbsp}}mg). Potassium is also available in tablets or capsules, which are formulated to allow potassium to leach slowly out of a matrix, since very high concentrations of potassium ion that occur adjacent to a solid tablet can injure the gastric or intestinal mucosa. For this reason, non-prescription potassium pills are limited by law in the US to a maximum of 99{{nbsp}}mg of potassium.

Potassium supplementation can also be combined with other metabolites, such as citrate or chloride, to achieve specific clinical effects.

Potassium supplements may be employed to mitigate the impact of hypertension, thereby reducing cardiovascular risk. [[Potassium chloride]] and [[potassium bicarbonate]] may be useful to control mild [[hypertension]]. In 2020, potassium was the 33rd most commonly prescribed medication in the U.S., with more than 17{{nbsp}}million prescriptions. Potassium supplementation has been shown to reduce both systolic and diastolic blood pressure in individuals with essential hypertension.

Additionally, potassium supplements may be employed with the aim of preventing the formation of kidney stones, a condition that can lead to renal complications if left untreated. Low potassium levels can lead to decreased calcium reabsorption in the kidneys, increasing the risk of elevated urine calcium and the formation of kidney stones. By maintaining adequate potassium levels, this risk can be reduced.

The mechanism of action of potassium involves various types of transporters and channels that facilitate its movement across cell membranes. This process can lead to an increase in the pumping of hydrogen ions. This, in turn, can escalate the production of gastric acid, potentially contributing to the development of gastric ulcers.

Potassium has a role in bone health. It contributes to the acid-base equilibrium in the body and helps protect bone tissue. Potassium salts produce an alkaline component that can aid in maintaining bone health.

For individuals with diabetes, potassium supplementation may be necessary, particularly for those with type 2 diabetes. Potassium is essential for the secretion of insulin by pancreatic beta cells, which helps regulate glucose levels. Without sufficient potassium, insulin secretion is compromised, leading to hyperglycemia and worsening diabetes.

Excessive potassium intake can have adverse effects, such as gastrointestinal discomfort and disturbances in heart rhythm.

Potassium supplementation can have side effects on ulceration, particularly in relation to peptic ulcer disease. Potassium channels have the potential to increase gastric acid secretion, which can lead to an increased risk of ulcerations. Medications used for peptic ulcer disease, known as "proton pump inhibitors", work by inhibiting potassium pumps that activate the H/K ATPase. This inhibition helps to reduce the secretion of hydrochloric acid into the parietal cell, thereby decreasing acidic synthesis and lowering the risk of ulcers. Nicorandil, a drug used for the treatment of ischemic heart disease, can stimulate nitrate and potassium ATP channels, and as a result, it has been associated with side effects such as GI, oral, and anal ulcers. Prolonged and chronic use of potassium supplements has been linked to more severe side effects, including ulcers outside of the gastrointestinal (GI) tract. Close monitoring is necessary for patients who are also taking angiotensinogen-converting enzyme inhibitors, angiotensin receptor blockers, or potassium-sparing diuretics.

====Detection by taste buds====

Potassium can be detected by taste because it triggers three of the five types of taste sensations, according to concentration. Dilute solutions of potassium ions 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 solutions makes high-dose potassium supplementation by liquid drinks a palatability challenge.

==Precautions==

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Potassium metal can react violently with water producing KOH and [[hydrogen]] gas.

:{{chem2|2 K(s) + 2 H2O(l) → 2 KOH(aq) + H2(g)↑}}

[[File:Potassium water 20.theora.ogv|thumb|left|alt=A piece of potassium metal is dropped into a clear container of water and skates 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 produced, and with potassium vapor, burns with a pink or lilac flame. Strongly alkaline potassium hydroxide is formed in solution.]]

This reaction is exothermic and releases sufficient heat to ignite the resulting hydrogen in the presence of oxygen. Finely powdered potassium ignites in air at room temperature. The bulk metal ignites in air if heated. Because its density is 0.89{{nbsp}}g/cm³, burning potassium floats in water that exposes it to atmospheric oxygen. Many common fire extinguishing agents, including water, either are ineffective or make a potassium fire worse. [[Nitrogen]], [[argon]], [[sodium chloride]] (table salt), [[sodium carbonate]] (soda ash), and [[silicon dioxide]] (sand) are effective if they are dry. Some [[Fire extinguisher|Class D]] dry powder extinguishers designed for metal fires are also effective. These agents deprive the fire of oxygen and cool the potassium metal.

During storage, potassium forms peroxides and superoxides. These peroxides may react violently with [[organic compound]]s such as oils. Both peroxides and superoxides may react explosively with metallic potassium.

Because potassium reacts with water vapor in the air, it is usually stored under anhydrous mineral oil or kerosene. Unlike lithium and sodium, potassium should not be stored under oil for longer than six months, unless in an inert (oxygen-free) atmosphere, or under vacuum. After prolonged storage in air dangerous shock-sensitive peroxides can form on the metal and under the lid of the container, and can detonate upon opening.

Ingestion of large amounts of potassium compounds can lead to [[hyperkalemia]], strongly influencing the cardiovascular system. Potassium chloride is used in the U.S. for [[lethal injection]] executions.

==See also==

{{Portal bar|Contents|Science|Medicine|Chemistry|Cornwall|Physics|Geophysics}}

==Bibliography==

* {{cite book|doi = 10.1002/14356007.a22_031.pub2|title = Ullmann's Encyclopedia of Industrial Chemistry|date = 2006|ref=Burkhardt|last = Burkhardt |first=Elizabeth R.|chapter = Potassium and Potassium Alloys|isbn = 978-3-527-30673-2|volume=A22|pages=31–38 }}

* [https://fdc.nal.usda.gov/fdc-app.html#/ National Nutrient Database] {{Webarchive|url=https://web.archive.org/web/20140810230234/http://ndb.nal.usda.gov/ndb/search/list |date=2014-08-10 }} at [[USDA]] Website

==External links==

{{Sister project links|auto=1|wikt=potassium|n=y|s=y|v=Potassium atom|b=Wikijunior:The Elements/Potassium}}

* {{pubchem|Potassium}}

{{Periodic table (navbox)}}

{{Potassium compounds}}

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[[Category:Potassium| ]]