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Potassium, 19K
Potassium pearls (in paraffin oil, ~5 mm each)
Appearancesilvery white, faint bluish-purple hue when exposed to air
Standard atomic weight Ar°(K)
Potassium in the periodic table
Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon
Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson


Atomic number (Z)19
Group group 1: hydrogen and alkali metals
Period period 4
Block   s-block
Electron configuration [ Ar ] 4s1
Electrons per shell2, 8, 8, 1
Physical properties
Phase at  STP solid
Melting point 336.7  K (63.5 °C,146.3 °F)
Boiling point 1030.793 K(757.643 °C,1395.757 °F) [3]
Density (near r.t.)0.8590 g/cm3 [4]
when liquid (at m.p.)0.82948 g/cm3 [3]
Critical point 2223 K, 16 MPa [5]
Heat of fusion 2.33  kJ/mol
Heat of vaporization 76.9 kJ/mol
Molar heat capacity 29.6 J/(mol·K)
Atomic properties
Oxidation states −1, +1 (a strongly basic oxide)
Electronegativity Pauling scale: 0.82
Ionization energies
  • 1st: 418.8 kJ/mol
  • 2nd: 3052 kJ/mol
  • 3rd: 4420 kJ/mol
  • (more)
Atomic radius empirical:227  pm
Covalent radius 203±12 pm
Van der Waals radius 275 pm
Potassium spectrum visible.png
Spectral lines of potassium
Other properties
Natural occurrence primordial
Crystal structure body-centered cubic (bcc)(cI2)
Lattice constant
a = 532.69 pm (at 20 °C) [4]
Thermal expansion 83.3 µm/(m⋅K)(at 25 °C)
Thermal conductivity 102.5 W/(m⋅K)
Electrical resistivity 72 nΩ⋅m(at 20 °C)
Magnetic ordering paramagnetic [6]
Molar magnetic susceptibility +20.8×10−6 cm3/mol(298 K) [7]
Young's modulus 3.53 GPa
Shear modulus 1.3 GPa
Bulk modulus 3.1 GPa
Speed of sound thin rod2000 m/s(at 20 °C)
Mohs hardness 0.4
Brinell hardness 0.363 MPa
CAS Number 7440-09-7
Discovery and first isolation Humphry Davy (1807)
Symbol"K": from New Latin kalium
Isotopes of potassium
Main isotopes Decay
abun­dance half-life (t1/2) mode pro­duct
39K93.3% stable
40K 0.0120%1.248×109 y β 40Ca
ε 40Ar
β+ 40Ar
Symbol category class.svg  Category: Potassium
| references

Potassium is a chemical element; it has symbol K (from Neo-Latin kalium) and atomic number  19. It is a silvery white metal that is soft enough to easily cut with a knife. [8] 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 metals, all of which have a single valence electron in the outer electron shell, which is easily removed to create an ion with a positive charge (which combines with anions 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-colored flame. It is found dissolved in seawater (which is 0.04% potassium by weight), [9] [10] and occurs in many minerals such as orthoclase, a common constituent of granites and other igneous rocks. [11]


Potassium is chemically very similar to sodium, the previous element in 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, [12] which was demonstrated in 1807 when elemental potassium was first isolated via electrolysis. Naturally occurring potassium is composed of three isotopes, of which 40
is radioactive. Traces of 40
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 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. [13]


The English name for the element potassium comes from the word potash , [14] 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 Arabic : القَلْيَهal-qalyah 'plant ashes'. In 1797, the German chemist 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. [15] In 1807, Humphry Davy produced the element via electrolysis: in 1809, Ludwig Wilhelm Gilbert proposed the name Kalium for Davy's "potassium". [16] In 1814, the Swedish chemist Berzelius advocated the name kalium for potassium, with the chemical symbol K. [17]

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. [18] The "Gold Book" of the International Union of Pure and Applied Chemistry has designated the official chemical symbol as K. [19]



The flame test of potassium. FlammenfarbungK.png
The flame test of potassium.

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. [20] In a flame test, potassium and its compounds emit a lilac color with a peak emission wavelength of 766.5 nanometers. [21]

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 kJ/mol, the potassium atom is much more likely to lose the last electron and acquire a positive charge, although negatively charged alkalide K ions are not impossible. [22] In contrast, the second ionization energy is very high (3052 kJ/mol).


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 desiccants, although they are no longer used as such. [23]


Structure of solid potassium superoxide (
KO2). Potassium-superoxide-unit-cell-3D-ionic.png
Structure of solid potassium superoxide (KO2).

Four oxides of potassium are well studied: potassium oxide (K2O), potassium peroxide (K2O2), potassium superoxide (KO2) [24] and potassium ozonide (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 kg of KOH can dissolve in a single liter of water. [25] [26] Anhydrous KOH is rarely encountered. KOH reacts readily with carbon dioxide (CO2) to produce potassium carbonate (K2CO3), and in principle could be used to remove traces of the gas from air. Like the closely related sodium hydroxide, KOH reacts with fats to produce soaps.

In general, potassium compounds are ionic and, owing to the high hydration energy of the K+ ion, have excellent water solubility. The main species in water solution are the aquo complexes [K(H2O)n]+ where n = 6 and 7. [27]

Potassium heptafluorotantalate (K2[TaF7]) is an intermediate in the purification of tantalum from the otherwise persistent contaminant of niobium. [28]

Organopotassium compounds illustrate nonionic compounds of potassium. They feature highly polar covalent K–C bonds. Examples include benzyl potassium KCH2C6H5. Potassium intercalates into graphite to give a variety of graphite intercalation compounds, including KC8.


There are 25 known isotopes of potassium, three of which occur naturally: 39
(93.3%), 40
(0.0117%), and 41
(6.7%) (by mole fraction). Naturally occurring 40
has a half-life of 1.250×109 years. It decays to stable 40
by electron capture or positron emission (11.2%) or to stable 40
by beta decay (88.8%). [29] The decay of 40
to 40
is the basis of a common 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 (40
) was quantitatively retained. Minerals are dated by measurement of the concentration of potassium and the amount of radiogenic 40
that has accumulated. The minerals best suited for dating include biotite, muscovite, metamorphic hornblende, and volcanic feldspar; whole rock samples from volcanic flows and shallow instrusives can also be dated if they are unaltered. [29] [30] Apart from dating, potassium isotopes have been used as tracers in studies of weathering and for nutrient cycling studies because potassium is a macronutrient required for life [31] on Earth.

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. 40
is the radioisotope with the largest abundance in the human body. In healthy animals and people, 40
represents the largest source of radioactivity, greater even than 14
. In a human body of 70 kg, about 4,400 nuclei of 40
decay per second. [32] The activity of natural potassium is 31 Bq/g. [33]



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, [12] and Henri Louis Duhamel du Monceau was able to prove this difference in 1736. [34] 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. [35] [36] For a long time the only significant applications for potash were the production of glass, bleach, soap and gunpowder as potassium nitrate. [37] 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. [13] The discovery by Justus Liebig in 1840 that potassium is a necessary element for plants and that most types of soil lack potassium [38] 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. [39] [40] [41] Other potash deposits were discovered, and by the 1960s Canada became the dominant producer. [42] [43]


Sir Humphry Davy Sir Humphry Davy, Bt by Thomas Phillips.jpg
Sir Humphry Davy
Pieces of potassium metal Potassium.JPG
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. [44] 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. [35] [36] [45] [46] 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. [36]

Because of the sensitivity of potassium to water and air, air-free techniques are normally employed for handling the element. It is unreactive toward nitrogen and saturated hydrocarbons such as mineral oil or kerosene. [47] It readily dissolves in liquid ammonia, up to 480 g per 1000 g of ammonia at 0 °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 KNH
, but this reaction is accelerated by minute amounts of transition metal salts. [48] 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 method. [49] Illustrative is the preparation of magnesium:

MgCl2 + 2 K → Mg + 2 KCl


Potassium in feldspar PotassiumFeldsparUSGOV.jpg
Potassium in feldspar

Potassium is formed in supernovae by nucleosynthesis from lighter atoms. Potassium is principally created in Type II supernovae via an explosive oxygen-burning process. [50] These are nuclear fusion reactions, not to be confused with chemical burning of potassium in oxygen. 40
is also formed in s-process nucleosynthesis and the neon burning process. [51]

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. [52] The potassium concentration in seawater is 0.39 g/L [9] (0.039 wt/v%), about one twenty-seventh the concentration of sodium. [53] [54]


Elemental potassium does not occur in nature because of its high reactivity. It reacts violently with water [47] 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 (KCl·MgCl2·6H2O), kainite (MgSO4·KCl·3H2O) and langbeinite (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. [54] 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. [55]

Commercial production


Sylvite from New Mexico Museo de La Plata - Silvita.jpg
Sylvite from New Mexico
Monte Kali, a potash mining and beneficiation waste heap in Hesse, Germany, consisting mostly of sodium chloride. Wintershall Monte Kali 12.jpg
Monte Kali, a potash mining and beneficiation waste heap in 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 seabeds, [53] 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. [56] [57] [58] 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 1,000 meters (3,300 feet) 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. [59] 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. [57] [58]

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 heaps. 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. [54]

Pure potassium metal can be isolated by electrolysis of its 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 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. [54] [60]

2 KF + CaC2 → 2 K + CaF2 + 2 C

Reagent-grade potassium metal costs about $10.00/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 detonates when scratched. The resulting explosion often starts a fire difficult to extinguish. [61] [62]

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

3 K+ + Na3[Co(NO2)6] → K3[Co(NO2)6] + 3 Na+

The potassium cobaltinitrite is obtained as a yellow solid.

Commercial uses


Potassium sulfate/magnesium sulfate fertilizer Patentkali (Potassium sulfate with magnesium).jpg
Potassium sulfate/magnesium sulfate fertilizer

Potassium ions are an essential component of plant nutrition and are found in most soil types. [13] They are used as a fertilizer in agriculture, horticulture, and hydroponic culture in the form of chloride (KCl), sulfate (K2SO4), or nitrate (KNO3), representing the 'K' in 'NPK'. Agricultural fertilizers consume 95% of global potassium chemical production, and about 90% of this potassium is supplied as KCl. [13] The potassium content of most plants ranges from 0.5% to 2% of the harvested weight of crops, conventionally expressed as amount of K2O. Modern high-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 (MgSO4·KCl·3H2O) and langbeinite (MgSO4·K2SO4). Only a very few fertilizers contain potassium nitrate. [63] In 2005, about 93% of world potassium production was consumed by the fertilizer industry. [58] Furthermore, potassium can play a key role in nutrient cycling by controlling litter composition. [64]

Medical use

Potassium citrate

Potassium citrate is used to treat a kidney stone condition called renal tubular acidosis. [65]

Potassium chloride

Potassium, in the form of potassium chloride is used as a medication to treat and prevent low blood potassium. [66] Low blood potassium may occur due to vomiting, diarrhea, or certain medications. [67] It is given by slow injection into a vein or by mouth. [68]

Food additives

Potassium sodium tartrate (KNaC4H4O6, Rochelle salt) is a main constituent of some varieties of baking powder; it is also used in the silvering of mirrors. Potassium bromate (KBrO3) is a strong oxidizer (E924), used to improve dough strength and rise height. 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. [69] [70]


Major potassium chemicals are potassium hydroxide, potassium carbonate, potassium sulfate, and potassium chloride. Megatons of these compounds are produced annually. [71]

KOH is a strong base, which is used in industry to neutralize strong and weak acids, to control pH and to manufacture potassium salts. It is also used to saponify fats and oils, in industrial cleaners, and in hydrolysis reactions, for example of esters. [72] [73]

Potassium nitrate (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 complexes. Its applications include gold mining, electroplating, and electroforming of these metals; it is also used in organic synthesis to make nitriles. Potassium carbonate (K2CO3 or potash) is used in the manufacture of glass, soap, color TV tubes, fluorescent lamps, textile dyes and pigments. [74] Potassium permanganate (KMnO4) is an oxidizing, bleaching and purification substance and is used for production of saccharin. Potassium chlorate (KClO3) is added to matches and explosives. Potassium bromide (KBr) was formerly used as a sedative and in photography. [13]

While potassium chromate (K2CrO4) is used in the manufacture of a host of different commercial products such as inks, dyes, wood stains (by reacting with the tannic acid in wood), explosives, fireworks, fly paper, and safety matches, [75] 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. [76]

Niche uses

There are thousands of uses of various potassium compounds. One example is potassium superoxide, KO2, an orange solid that acts as a portable source of oxygen and a carbon dioxide absorber. It is widely used in respiration systems in mines, submarines and spacecraft as it takes less volume than the gaseous oxygen. [77] [78]

4 KO2 + 2 CO2 → 2 K2CO3 + 3 O2

Another example is potassium cobaltinitrite, K3[Co(NO2)6], which is used as artist's pigment under the name of Aureolin or Cobalt Yellow. [79]

The stable isotopes of potassium can be laser cooled and used to probe fundamental and technological problems in quantum physics. The two bosonic isotopes possess convenient Feshbach resonances to enable studies requiring tunable interactions, while 40
is one of only two stable fermions amongst the alkali metals. [80]

Laboratory uses

An alloy of sodium and potassium, NaK is a liquid used as a heat-transfer medium and a desiccant for producing dry and air-free solvents. It can also be used in reactive distillation. [81] The ternary alloy of 12% Na, 47% K and 41% Cs has the lowest melting point of −78 °C of any metallic compound. [20]

Metallic potassium is used in several types of magnetometers. [82]

Biological role

Potassium is the eighth or ninth most common element by mass (0.2%) in the human body, so that a 60 kg adult contains a total of about 120 g of potassium. [83] 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). [84] Potassium ions are present in a wide variety of proteins and enzymes. [85]

Biochemical function

Potassium levels influence multiple physiological processes, including [86] [87] [88]


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. [91] Levels outside this range are associated with an increasing rate of death from multiple causes, [92] and some cardiac, kidney, [93] and lung diseases progress more rapidly if serum potassium levels are not maintained within the normal range.

An average meal of 40–50 mmol presents the body with more potassium than is present in all plasma (20–25 mmol). This surge causes the plasma potassium to rise up to 10% before clearance by renal and extrarenal mechanisms. [94]

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 (diuresis). [95] Deficiency symptoms include muscle weakness, paralytic ileus, ECG abnormalities, decreased reflex response; and in severe cases, respiratory paralysis, alkalosis, and cardiac arrhythmia. [96]

Control mechanisms

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

  1. a reactive negative-feedback system,
  2. a reactive feed-forward system,
  3. a predictive or circadian system, and
  4. an internal or cell membrane transport system.

Collectively, the first three are sometimes termed the "external potassium homeostasis system"; [97] and the first two, the "reactive potassium homeostasis system".

  • The reactive negative-feedback system refers to the system that induces renal secretion of potassium in response to a rise in the plasma potassium (potassium ingestion, shift out of cells, or intravenous infusion.)
  • The reactive feed-forward system refers to an incompletely understood system that induces renal potassium secretion in response to potassium ingestion prior to any rise in the plasma potassium. This is probably initiated by gut cell potassium receptors that detect ingested potassium and trigger vagal afferent signals to the pituitary gland.
  • The predictive or circadian system increases renal secretion of potassium during mealtime hours (e.g. daytime for humans, nighttime for rodents) independent of the presence, amount, or absence of potassium ingestion. It is mediated by a circadian oscillator in the suprachiasmatic nucleus of the brain (central clock), which causes the kidney (peripheral clock) to secrete potassium in this rhythmic circadian fashion.
    The action of the sodium-potassium pump is an example of primary active transport. The two carrier proteins embedded in the cell membrane on the left are using ATP to move sodium out of the cell against the concentration gradient; The two proteins on the right are using secondary active transport to move potassium into the cell. This process results in reconstitution of ATP. Scheme sodium-potassium pump-en.svg
    The action of the sodium-potassium pump is an example of primary active transport. The two carrier proteins embedded in the cell membrane on the left are using ATP to move sodium out of the cell against the concentration gradient; The two proteins on the right are using secondary active transport to move potassium into the cell. This process results in reconstitution of ATP.
  • 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. [98] This ion pump uses 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 tetramers) are crucial for hyperpolarization inside neurons 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. [99]

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 mmol/L, 4.8 g/L), while sodium is the major cation of extracellular fluid (150 mmol/L, 3.345 g/L). In the kidneys, about 180 liters of plasma is filtered through the glomeruli and into the renal tubules per day. [100] This filtering involves about 600 mg of sodium and 33 mg of potassium. Since only 1–10 mg of sodium and 1–4 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. [101] Sodium pumps 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 [102] [103] in counter flow to sodium in response to an apparent (but not actual) Donnan equilibrium, [104] 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. [105] 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.[ citation needed ]

With no potassium intake, it is excreted at about 200 mg per day until, in about a week, potassium in the serum declines to a mildly deficient level of 3.0–3.5 mmol/L. [106] If potassium is still withheld, the concentration continues to fall until a severe deficiency causes eventual death. [107]

The potassium moves passively through pores in the cell membrane. When ions move through ion transporters (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 channels have only one gate, and there only one kind of ion can stream through, at 10 million to 100 million ions per second. [108] Calcium is required to open the pores, [109] although calcium may work in reverse by blocking at least one of the pores. [110] Carbonyl groups inside the pore on the amino acids mimic the water hydration that takes place in water solution [111] by the nature of the electrostatic charges on four carbonyl groups inside the pore. [112]


Dietary recommendations

The U.S. National Academy of Medicine (NAM), on behalf of both the U.S. and Canada, sets Dietary Reference Intakes, including Estimated Average Requirements (EARs) and Recommended Dietary Allowances (RDAs), or Adequate Intakes (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 mg of potassium for 0 to 6-month-old infants, 860 mg of potassium for 7 to 12-month-old infants, 2,000 mg of potassium for 1 to 3-year-old children, and 2,300 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 mg of potassium for 9 to 13-year-old males, 3,000 mg of potassium for 14 to 18-year-old males, and 3,400 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 mg of potassium for 9 to 18-year-old females, and 2,600 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 mg of potassium for 14 to 18-year-old pregnant females, 2,900 mg for pregnant females that are 19 years of age and older; furthermore, 2,500 mg of potassium for 14 to 18-year-old lactating females, and 2,800 mg for lactating females that are 19 years of age and older. As for safety, the NAM also sets tolerable upper intake levels (ULs) for vitamins and minerals, but for potassium the evidence was insufficient, so no UL was established. [113] [114]

As of 2004, most Americans adults consume less than 3,000 mg. [115]

Likewise, in the European Union, in particular in Germany, and Italy, insufficient potassium intake is somewhat common. [116] The British National Health Service recommends a similar intake, saying that adults need 3,500 mg per day and that excess amounts may cause health problems such as stomach pain and diarrhea. [117]

In 2019, the National Academies of Sciences, Engineering, and Medicine revised the Adequate Intake for potassium to 2,600 mg/day for females 19 years of age and older who are not pregnant or lactating, and 3,400 mg/day for males 19 years of age and older. [118] [119]

Food sources

Potassium is present in all fruits, vegetables, meat and fish. Foods with high potassium concentrations include yam, parsley, dried apricots, milk, chocolate, all nuts (especially almonds and pistachios), potatoes, bamboo shoots, bananas, avocados, coconut water, soybeans, and bran. [120]

The United States Department of Agriculture also lists tomato paste, orange juice, beet greens, white beans, 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. [121]

Deficient intake

Diets low in potassium can lead to hypertension [122] and hypokalemia.


Supplements of potassium are most widely used in conjunction with diuretics that block reabsorption of sodium and water upstream from the distal tubule (thiazides and loop diuretics), because this promotes increased distal tubular potassium secretion, with resultant increased potassium excretion.[ medical citation needed ] A variety of prescription and over-the counter supplements are available.[ citation needed ] Potassium chloride may be dissolved in water, but the salty/bitter taste makes liquid supplements unpalatable. [123] Typical doses range from 10 mmol (400 mg), to 20 mmol (800 mg).[ medical citation needed ] 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. [68] [ medical citation needed ] For this reason, non-prescription potassium pills are limited by law in the US to a maximum of 99 mg of potassium.[ citation needed ]

Potassium supplements may be employed to mitigate the impact of hypertension, thereby reducing cardiovascular risk. 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. 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. [124]

A meta-analysis concluded that a 1640 mg increase in the daily intake of potassium was associated with a 21% lower risk of stroke. [125] Potassium chloride and potassium bicarbonate may be useful to control mild hypertension. [126] In 2020, potassium was the 33rd most commonly prescribed medication in the U.S., with more than 17 million prescriptions. [127] [128]

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. [123] [129]


GHS labelling:
GHS-pictogram-flamme.svg GHS-pictogram-acid.svg
H260, H314
P223, P231+P232, P280, P305+P351+P338, P370+P378, P422 [130]
NFPA 704 (fire diamond)
NFPA 704.svgHealth 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasFlammability 3: Liquids and solids that can be ignited under almost all ambient temperature conditions. Flash point between 23 and 38 °C (73 and 100 °F). E.g. gasolineInstability 2: Undergoes violent chemical change at elevated temperatures and pressures, reacts violently with water, or may form explosive mixtures with water. E.g. white phosphorusSpecial hazard W: Reacts with water in an unusual or dangerous manner. E.g. sodium, sulfuric acid

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

2 K(s) + 2 H2O(l) → 2 KOH(aq) + H2(g)↑
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 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 Class D dry powder extinguishers designed for metal fires are also effective. These agents deprive the fire of oxygen and cool the potassium metal. [131]

During storage, potassium forms peroxides and superoxides. These peroxides may react violently with organic compounds such as oils. Both peroxides and superoxides may react explosively with metallic potassium. [132]

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. [133]

Ingestion of large amounts of potassium compounds can lead to hyperkalemia, strongly influencing the cardiovascular system. [134] [135] Potassium chloride is used in the U.S. for lethal injection executions. [134]

See also

Related Research Articles

<span class="mw-page-title-main">Alkali metal</span> Group of highly reactive chemical elements

The alkali metals consist of the chemical elements lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs), and francium (Fr). Together with hydrogen they constitute group 1, which lies in the s-block of the periodic table. All alkali metals have their outermost electron in an s-orbital: this shared electron configuration results in their having very similar characteristic properties. Indeed, the alkali metals provide the best example of group trends in properties in the periodic table, with elements exhibiting well-characterised homologous behaviour. This family of elements is also known as the lithium family after its leading element.

In chemistry, an alkali is a basic, ionic salt of an alkali metal or an alkaline earth metal. An alkali can also be defined as a base that dissolves in water. A solution of a soluble base has a pH greater than 7.0. The adjective alkaline, and less often, alkalescent, is commonly used in English as a synonym for basic, especially for bases soluble in water. This broad use of the term is likely to have come about because alkalis were the first bases known to obey the Arrhenius definition of a base, and they are still among the most common bases.

<span class="mw-page-title-main">Caesium</span> Chemical element, symbol Cs and atomic number 55

Caesium is a chemical element; it has symbol Cs and atomic number 55. It is a soft, silvery-golden alkali metal with a melting point of 28.5 °C, which makes it one of only five elemental metals that are liquid at or near room temperature. Caesium has physical and chemical properties similar to those of rubidium and potassium. It is pyrophoric and reacts with water even at −116 °C (−177 °F). It is the least electronegative element, with a value of 0.79 on the Pauling scale. It has only one stable isotope, caesium-133. Caesium is mined mostly from pollucite. Caesium-137, a fission product, is extracted from waste produced by nuclear reactors. It has the largest atomic radius of all elements whose radii have been measured or calculated, at about 260 picometers.

<span class="mw-page-title-main">Lithium carbonate</span> Chemical compound

Lithium carbonate is an inorganic compound, the lithium salt of carbonic acid with the formula Li
. This white salt is widely used in processing metal oxides. It is on the World Health Organization's List of Essential Medicines for its efficacy in the treatment of mood disorders such as bipolar disorder.

<span class="mw-page-title-main">Nitrate</span> Polyatomic ion (NO₃, charge –1) found in explosives and fertilisers

Nitrate is a polyatomic ion with the chemical formula NO
. Salts containing this ion are called nitrates. Nitrates are common components of fertilizers and explosives. Almost all inorganic nitrates are soluble in water. An example of an insoluble nitrate is bismuth oxynitrate.

<span class="mw-page-title-main">Rubidium</span> Chemical element, symbol Rb and atomic number 37

Rubidium is a chemical element; it has symbol Rb and atomic number 37. It is a very soft, whitish-grey solid in the alkali metal group, similar to potassium and caesium. Rubidium is the first alkali metal in the group to have a density higher than water. On Earth, natural rubidium comprises two isotopes: 72% is a stable isotope 85Rb, and 28% is slightly radioactive 87Rb, with a half-life of 48.8 billion years—more than three times as long as the estimated age of the universe.

<span class="mw-page-title-main">Sodium</span> Chemical element, symbol Na and atomic number 11

Sodium is a chemical element; it has symbol Na and atomic number 11. It is a soft, silvery-white, highly reactive metal. Sodium is an alkali metal, being in group 1 of the periodic table. Its only stable isotope is 23Na. The free metal does not occur in nature and must be prepared from compounds. Sodium is the sixth most abundant element in the Earth's crust and exists in numerous minerals such as feldspars, sodalite, and halite (NaCl). Many salts of sodium are highly water-soluble: sodium ions have been leached by the action of water from the Earth's minerals over eons, and thus sodium and chlorine are the most common dissolved elements by weight in the oceans.

<span class="mw-page-title-main">Salt (chemistry)</span> Chemical compound involving ionic bonding

In chemistry, a salt or ionic compound is a chemical compound consisting of an ionic assembly of positively charged cations and negatively charged anions, which results in a neutral compound with no net electric charge. The constituent ions are held together by electrostatic forces termed ionic bonds.

The term chloride refers either to a chloride ion, which is a negatively charged chlorine atom, or a non-charged chlorine atom covalently bonded to the rest of the molecule by a single bond. Many inorganic chlorides are salts. Many organic compounds are chlorides. The pronunciation of the word "chloride" is.

Fluoride is an inorganic, monatomic anion of fluorine, with the chemical formula F
, whose salts are typically white or colorless. Fluoride salts typically have distinctive bitter tastes, and are odorless. Its salts and minerals are important chemical reagents and industrial chemicals, mainly used in the production of hydrogen fluoride for fluorocarbons. Fluoride is classified as a weak base since it only partially associates in solution, but concentrated fluoride is corrosive and can attack the skin.

<span class="mw-page-title-main">Sodium carbonate</span> Chemical compound

Sodium carbonate is the inorganic compound with the formula Na2CO3 and its various hydrates. All forms are white, odourless, water-soluble salts that yield alkaline solutions in water. Historically, it was extracted from the ashes of plants grown in sodium-rich soils, and because the ashes of these sodium-rich plants were noticeably different from ashes of wood, sodium carbonate became known as "soda ash". It is produced in large quantities from sodium chloride and limestone by the Solvay process, as well as by carbonating sodium hydroxide which is made using the Chlor-alkali process.

<span class="mw-page-title-main">Potassium chloride</span> Ionic compound (KCl)

Potassium chloride is a metal halide salt composed of potassium and chlorine. It is odorless and has a white or colorless vitreous crystal appearance. The solid dissolves readily in water, and its solutions have a salt-like taste. Potassium chloride can be obtained from ancient dried lake deposits. KCl is used as a fertilizer, in medicine, in scientific applications, domestic water softeners, and in food processing, where it may be known as E number additive E508.

<span class="mw-page-title-main">Potassium in biology</span> Use of Potassium by organisms

Potassium is the main intracellular ion for all types of cells, while having a major role in maintenance of fluid and electrolyte balance. Potassium 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 within plant cells, and in a mixed diet, it is most highly concentrated in fruits. The high concentration of potassium in plants, associated with comparatively very low amounts of sodium there, historically resulted in potassium first being isolated from the ashes of plants (potash), which in turn gave the element its modern name. The high concentration of potassium in plants means that heavy crop production rapidly depletes soils of potassium, and agricultural fertilizers consume 93% of the potassium chemical production of the modern world economy.

<span class="mw-page-title-main">Sodium sulfate</span> Chemical compound with formula Na2SO4

Sodium sulfate (also known as sodium sulphate or sulfate of soda) is the inorganic compound with formula Na2SO4 as well as several related hydrates. All forms are white solids that are highly soluble in water. With an annual production of 6 million tonnes, the decahydrate is a major commodity chemical product. It is mainly used as a filler in the manufacture of powdered home laundry detergents and in the Kraft process of paper pulping for making highly alkaline sulfides.

A bromide ion is the negatively charged form (Br) of the element bromine, a member of the halogens group on the periodic table. Most bromides are colorless. Bromides have many practical roles, being found in anticonvulsants, flame-retardant materials, and cell stains. Although uncommon, chronic toxicity from bromide can result in bromism, a syndrome with multiple neurological symptoms. Bromide toxicity can also cause a type of skin eruption, see potassium bromide. The bromide ion has an ionic radius of 196 pm.

<span class="mw-page-title-main">Radical anion</span> Free radical species

In organic chemistry, a radical anion is a free radical species that carries a negative charge. Radical anions are encountered in organic chemistry as reduced derivatives of polycyclic aromatic compounds, e.g. sodium naphthenide. An example of a non-carbon radical anion is the superoxide anion, formed by transfer of one electron to an oxygen molecule. Radical anions are typically indicated by .

<span class="mw-page-title-main">Carnallite</span> Evaporite mineral

Carnallite (also carnalite) is an evaporite mineral, a hydrated potassium magnesium chloride with formula KCl.MgCl2·6(H2O). It is variably colored yellow to white, reddish, and sometimes colorless or blue. It is usually massive to fibrous with rare pseudohexagonal orthorhombic crystals. The mineral is deliquescent (absorbs moisture from the surrounding air) and specimens must be stored in an airtight container.

Sodium atoms have 11 electrons, one more than the stable configuration of the noble gas neon. As a result, sodium usually forms ionic compounds involving the Na+ cation. Sodium is a reactive alkali metal and is much more stable in ionic compounds. It can also form intermetallic compounds and organosodium compounds. Sodium compounds are often soluble in water.

<span class="mw-page-title-main">Salt substitute</span> Low-sodium table salt alternative

A salt substitute, also known as low-sodium salt, is a low-sodium alternative to edible salt marketed to reduce the risk of high blood pressure and cardiovascular disease associated with a high intake of sodium chloride while maintaining a similar taste.

<span class="mw-page-title-main">Sodium in biology</span> Use of Sodium by organisms

Sodium ions are necessary in small amounts for some types of plants, but sodium as a nutrient is more generally needed in larger amounts by animals, due to their use of it for generation of nerve impulses and for maintenance of electrolyte balance and fluid balance. In animals, sodium ions are necessary for the aforementioned functions and for heart activity and certain metabolic functions. The health effects of salt reflect what happens when the body has too much or too little sodium. Characteristic concentrations of sodium in model organisms are: 10 mM in E. coli, 30 mM in budding yeast, 10 mM in mammalian cell and 100 mM in blood plasma.


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