Jump to content

Potassium

Edit links
This is a good article. Click here for more information.
Page semi-protected
From Wikipedia, the free encyclopedia
(Redirected from K (element))

This article is about the chemical element. For the use of potassium as a medication, see Potassium chloride (medical use). For the use of potassium in biology, see Potassium in biology.

Potassium, 19K
Potassium pearls (in paraffin oil, ~5 mm each)
Potassium
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
Na
↑
K
↓
Rb
argon ← potassium → calcium
Atomic number (Z)19
Groupgroup 1: hydrogen and alkali metals
Periodperiod 4
Block  s-block
Electron configuration[Ar] 4s1
Electrons per shell2, 8, 8, 1
Physical properties
Phase at STPsolid
Melting point336.7 K ​(63.5 °C, ​146.3 °F)
Boiling point1030.793 K ​(757.643 °C, ​1395.757 °F)[3]
Density (at 20° C)0.8590 g/cm3 [4]
when liquid (at m.p.)0.82948 g/cm3[3]
Critical point2223 K, 16 MPa[5]
Heat of fusion2.33 kJ/mol
Heat of vaporization76.9 kJ/mol
Molar heat capacity29.6 J/(mol·K)
Atomic properties
Oxidation statescommon: +1
−1?
ElectronegativityPauling scale: 0.82
Ionization energies
  • 1st: 418.8 kJ/mol
  • 2nd: 3052 kJ/mol
  • 3rd: 4420 kJ/mol
  • (more)
Atomic radiusempirical: 227 pm
Covalent radius203±12 pm
Van der Waals radius275 pm
Color lines in a spectral range
Spectral lines of potassium
Other properties
Natural occurrenceprimordial
Crystal structure ​body-centered cubic (bcc) (cI2)
Lattice constant
Body-centered cubic crystal structure for potassium
a = 532.69 pm (at 20 °C)[4]
Thermal expansion77.37×10−6/K (at 20 °C)[4]
Thermal conductivity102.5 W/(m⋅K)
Electrical resistivity72 nΩ⋅m (at 20 °C)
Magnetic orderingparamagnetic[6]
Molar magnetic susceptibility+20.8×10−6 cm3/mol (298 K)[7]
Young's modulus3.53 GPa
Shear modulus1.3 GPa
Bulk modulus3.1 GPa
Speed of sound thin rod2000 m/s (at 20 °C)
Mohs hardness0.4
Brinell hardness0.363 MPa
CAS Number7440-09-7
History
Discovery and first isolationHumphry Davy (1807)
Symbol"K": from New Latin kalium
Isotopes of potassium
Main isotopes Decay
abun­dance half-life (t1/2) mode pro­duct
39K 93.3% stable
40K 0.0120% 1.248×109 y β− 40Ca
ε 40Ar
β+ 40Ar
41K 6.73% stable
 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
K is radioactive. Traces of 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 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]

Etymology

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: القَلْيَه al-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]

Properties

Physical

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).

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

Compounds

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.

Isotopes

Main article: Isotopes of potassium

There are 25 known isotopes of potassium, three of which occur naturally: 39
K (93.3%), 40
K (0.0117%), and 41
K (6.7%) (by mole fraction). Naturally occurring 40
K has a half-life of 1.250×109 years. It decays to stable 40
Ar by electron capture or positron emission (11.2%) or to stable 40
Ca by beta decay (88.8%).[29] The decay of 40
K to 40
Ar 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
Ar) was quantitatively retained. Minerals are dated by measurement of the concentration of potassium and the amount of radiogenic 40
Ar 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.

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

History

Potash

Main article: 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,[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]

Metal

Sir Humphry Davy
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
2, 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

Occurrence

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
K 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]

Geology

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

Mining

Sylvite from New Mexico
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

Fertilizer

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

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.[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]

Industrial

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
K 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

Main article: Potassium in biology

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]

  • resting cellular-membrane potential and the propagation of action potentials in neuronal, muscular, and cardiac tissue. Due to the electrostatic and chemical properties, K+ ions are larger than Na+ ions, and ion channels and pumps in cell membranes can differentiate between the two ions, actively pumping or passively passing one of the two ions while blocking the other.[89]
  • hormone secretion and action
  • vascular tone
  • systemic blood pressure control
  • gastrointestinal motility
  • acid–base homeostasis
  • glucose and insulin metabolism
  • mineralocorticoid action
  • renal concentrating ability
  • fluid and electrolyte balance
  • local cortical monoaminergic norepinephrine, serotonin, and dopamine levels, and through them, sleep/wake balance, and spontaneous activity.[90]

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.[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.
  • 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]

Nutrition

Dietary recommendations

North America

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]

Europe

Likewise, in the European Union, in particular in Germany, and Italy, insufficient potassium intake is somewhat common.[116]

The National Health Service in the United Kingdom recommends a similar intake, saying that "adults (19 to 64 years) need need 3500 mg per day" and that excess amounts may cause health problems such as stomach pain and diarrhea.[117]

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.[118]

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.[119]

Deficient intake

Main article: Hypokalemia

Mild hypokalemia does not cause distinct symptoms acting instead as a risk factor for hypertension[120][121] and cardiac arrhythmia.[122] Severe hypokalemia usually presents with hypertension, arrhythmia, muscle cramps, fatigue, weakness and constipation.[122] Causes of hypokalemia include vomiting, diarrhea, medications like furosemide and steroids, dialysis, diabetes insipidus, hyperaldosteronism, hypomagnesemia.[122]

Supplementation

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.[123] A variety of prescription and over-the counter supplements are available.[124] Potassium chloride may be dissolved in water, but the salty/bitter taste makes liquid supplements unpalatable.[125][126] 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][127] For this reason, non-prescription potassium pills are limited by law in the US to a maximum of 99 mg of potassium.[128]

Potassium supplementation can also be combined with other metabolites, such as citrate or chloride, to achieve specific clinical effects.[129]

Potassium supplements may be employed to mitigate the impact of hypertension, thereby reducing cardiovascular risk.[130] Potassium chloride and potassium bicarbonate may be useful to control mild hypertension.[131] In 2020, potassium was the 33rd most commonly prescribed medication in the U.S., with more than 17 million prescriptions.[132][133] Potassium supplementation has been shown to reduce both systolic and diastolic blood pressure in individuals with essential hypertension.[129]

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.[129]

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.[129]

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.[129]

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.[129]

Excessive potassium intake can have adverse effects, such as gastrointestinal discomfort and disturbances in heart rhythm.[129]

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.[129]


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.[125][134]

Precautions

Potassium
Hazards
GHS labelling:
GHS02: FlammableGHS05: Corrosive
Danger
H260, H314
P223, P231+P232, P280, P305+P351+P338, P370+P378, P422[135]
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 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
3
3
2
W

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.[136]

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.[137]

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.[138]

Ingestion of large amounts of potassium compounds can lead to hyperkalemia, strongly influencing the cardiovascular system.[139][140] Potassium chloride is used in the U.S. for lethal injection executions.[139]

See also

References

  1. ^ "Standard Atomic Weights: Potassium". CIAAW. 1979.
  2. ^ Prohaska T, Irrgeher J, Benefield J, Böhlke JK, Chesson LA, Coplen TB, Ding T, Dunn PJ, Gröning M, Holden NE, Meijer HA (2022-05-04). "Standard atomic weights of the elements 2021 (IUPAC Technical Report)". Pure and Applied Chemistry. doi:10.1515/pac-2019-0603. ISSN 1365-3075.
  3. ^ a b Aitken F, Volino F (January 2022). "New equations of state describing both the dynamic viscosity and self-diffusion coefficient for potassium and thallium in their fluid phases". Physics of Fluids. 34 (1): 017112. doi:10.1063/5.0079944.
  4. ^ a b c Arblaster JW (2018). Selected Values of the Crystallographic Properties of Elements. Materials Park, Ohio: ASM International. ISBN 978-1-62708-155-9.
  5. ^ Haynes WM, ed. (2011). CRC Handbook of Chemistry and Physics (92nd ed.). Boca Raton, FL: CRC Press. p. 4.122. ISBN 1-4398-5511-0.
  6. ^ Magnetic susceptibility of the elements and inorganic compounds, in Lide DR, ed. (2005). CRC Handbook of Chemistry and Physics (86th ed.). Boca Raton (FL): CRC Press. ISBN 0-8493-0486-5.
  7. ^ Weast R (1984). CRC, Handbook of Chemistry and Physics. Boca Raton, Florida: Chemical Rubber Company Publishing. pp. E110. ISBN 0-8493-0464-4.
  8. ^ Augustyn A. "Potassium/ Chemical element". Encyclopedia Britannica. Archived from the original on 2019-07-09. Retrieved 2019-04-17. Potassium Physical properties
  9. ^ a b Webb DA (April 1939). "The Sodium and Potassium Content of Sea Water" (PDF). The Journal of Experimental Biology (2): 183. Archived (PDF) from the original on 2019-09-24. Retrieved 2017-07-23.
  10. ^ Anthoni J (2006). "Detailed composition of seawater at 3.5% salinity". seafriends.org.nz. Archived from the original on 2019-01-18. Retrieved 2011-09-23.
  11. ^ Halperin ML, Kamel KS (1998-07-11). "Potassium". The Lancet. 352 (9122): 135–140. doi:10.1016/S0140-6736(98)85044-7. ISSN 0140-6736. PMID 9672294. S2CID 208790031. Archived from the original on 2021-06-07. Retrieved 2021-06-07.
  12. ^ a b Marggraf AS (1761). Chymische Schriften. p. 167.
  13. ^ a b c d e Greenwood, p. 73
  14. ^ Davy H (1808). "On some new phenomena of chemical changes produced by electricity, in particular the decomposition of the fixed alkalies, and the exhibition of the new substances that constitute their bases; and on the general nature of alkaline bodies". Philosophical Transactions of the Royal Society. 98: 32. doi:10.1098/rstl.1808.0001.
  15. ^ Klaproth, M. (1797) "Nouvelles données relatives à l'histoire naturelle de l'alcali végétal" (New data regarding the natural history of the vegetable alkali), Mémoires de l'Académie royale des sciences et belles-lettres (Berlin), pp. 9–13 ; see p. 13. Archived 2020-01-24 at the Wayback Machine From p. 13: "Cet alcali ne pouvant donc plus être envisagé comme un produit de la végétation dans les plantes, occupe une place propre dans la série des substances primitivement simples du règne minéral, &I il devient nécessaire de lui assigner un nom, qui convienne mieux à sa nature.
    La dénomination de Potasche (potasse) que la nouvelle nomenclature françoise a consacrée comme nom de tout le genre, ne sauroit faire fortune auprès des chimistes allemands, qui sentent à quel point la dérivation étymologique en est vicieuse. Elle est prise en effet de ce qu'anciennement on se servoit pour la calcination des lessives concentrées des cendres, de pots de fer (pott en dialecte de la Basse-Saxe) auxquels on a substitué depuis des fours à calciner.
    Je propose donc ici, de substituer aux mots usités jusqu'ici d'alcali des plantes, alcali végétal, potasse, &c. celui de kali, & de revenir à l'ancienne dénomination de natron, au lieu de dire alcali minéral, soude &c."
    (This alkali [i.e., potash] — [which] therefore can no longer be viewed as a product of growth in plants — occupies a proper place in the originally simple series of the mineral realm, and it becomes necessary to assign it a name that is better suited to its nature.
    The name of "potash" (potasse), which the new French nomenclature has bestowed as the name of the entire species [i.e., substance], would not find acceptance among German chemists, who feel to some extent [that] the etymological derivation of it is faulty. Indeed, it is taken from [the vessels] that one formerly used for the roasting of washing powder concentrated from cinders: iron pots (pott in the dialect of Lower Saxony), for which roasting ovens have been substituted since then.
    Thus I now propose to substitute for the until now common words of "plant alkali", "vegetable alkali", "potash", etc., that of kali ; and to return to the old name of natron instead of saying "mineral alkali", "soda", etc.)
  16. ^ Davy, Humphry (1809). "Ueber einige neue Erscheinungen chemischer Veränderungen, welche durch die Electricität bewirkt werden; insbesondere über die Zersetzung der feuerbeständigen Alkalien, die Darstellung der neuen Körper, welche ihre Basen ausmachen, und die Natur der Alkalien überhaupt" [On some new phenomena of chemical changes that are achieved by electricity; particularly the decomposition of flame-resistant alkalis [i.e., alkalies that cannot be reduced to their base metals by flames], the preparation of new substances that constitute their [metallic] bases, and the nature of alkalies generally]. Annalen der Physik. 31 (2): 113–175. Bibcode:1809AnP....31..113D. doi:10.1002/andp.18090310202. p. 157: In unserer deutschen Nomenclatur würde ich die Namen Kalium und Natronium vorschlagen, wenn man nicht lieber bei den von Herrn Erman gebrauchten und von mehreren angenommenen Benennungen Kali-Metalloid and Natron-Metalloid, bis zur völligen Aufklärung der chemischen Natur dieser räthzelhaften Körper bleiben will. Oder vielleicht findet man es noch zweckmässiger fürs Erste zwei Klassen zu machen, Metalle und Metalloide, und in die letztere Kalium und Natronium zu setzen. — Gilbert. (In our German nomenclature, I would suggest the names Kalium and Natronium, if one would not rather continue with the appellations Kali-metalloid and Natron-metalloid which are used by Mr. Erman [i.e., German physics professor Paul Erman (1764–1851)] and accepted by several [people], until the complete clarification of the chemical nature of these puzzling substances. Or perhaps one finds it yet more advisable for the present to create two classes, metals and metalloids, and to place Kalium and Natronium in the latter — Gilbert.)
  17. ^ Berzelius, J. Jacob (1814) Försök, att, genom användandet af den electrokemiska theorien och de kemiska proportionerna, grundlägga ett rent vettenskapligt system för mineralogien [Attempt, by the use of electrochemical theory and chemical proportions, to found a pure scientific system for mineralogy]. Stockholm, Sweden: A. Gadelius., p. 87.
  18. ^ 19. Kalium (Potassium) – Elementymology & Elements Multidict Archived 2019-02-18 at the Wayback Machine. vanderkrogt.net
  19. ^ McNaught, A. D. and Wilkinson,A. eds. (1997). Compendium of Chemical Terminology, 2nd ed. (the "Gold Book"). IUPAC. Blackwell Scientific Publications, Oxford.
  20. ^ a b Greenwood, p. 76
  21. ^ Greenwood, p. 75
  22. ^ Dye JL (1979). "Compounds of Alkali Metal Anions". Angewandte Chemie International Edition. 18 (8): 587–598. doi:10.1002/anie.197905871.
  23. ^ Williams DB, Lawton M (2010). "Drying of Organic Solvents: Quantitative Evaluation of the Efficiency of Several Desiccants". The Journal of Organic Chemistry. 75 (24): 8351–8354. doi:10.1021/jo101589h. PMID 20945830. S2CID 17801540.
  24. ^ Lide DR (1998). Handbook of Chemistry and Physics (87th ed.). Boca Raton, Florida, United States: CRC Press. pp. 477, 520. ISBN 978-0-8493-0594-8.
  25. ^ Lide DR, ed. (2005). CRC Handbook of Chemistry and Physics (86th ed.). Boca Raton (FL): CRC Press. p. 4–80. ISBN 0-8493-0486-5.
  26. ^ Schultz, p. 94
  27. ^ Lincoln, S. F.; Richens, D. T. and Sykes, A. G. "Metal Aqua Ions" in J. A. McCleverty and T. J. Meyer (eds.) Comprehensive Coordination Chemistry II Archived 2019-04-19 at the Wayback Machine, Vol. 1, pp. 515–555, ISBN 978-0-08-043748-4.
  28. ^ Anthony Agulyanski (2004). "Fluorine chemistry in the processing of tantalum and niobium". In Anatoly Agulyanski (ed.). Chemistry of Tantalum and Niobium Fluoride Compounds (1st ed.). Burlington: Elsevier. ISBN 978-0-08-052902-8.
  29. ^ a b Audi G, Bersillon O, Blachot J, Wapstra AH (2003), "The NUBASE evaluation of nuclear and decay properties", Nuclear Physics A, 729: 3–128, Bibcode:2003NuPhA.729....3A, doi:10.1016/j.nuclphysa.2003.11.001
  30. ^ Bowen R, Attendorn HG (1988). "Theory and Assumptions in Potassium–Argon Dating". Isotopes in the Earth Sciences. Springer. pp. 203–8. ISBN 978-0-412-53710-3.
  31. ^ Anaç, D., Martin-Prével, P. (1999). Improved crop quality by nutrient management. Springer. pp. 290–. ISBN 978-0-7923-5850-3.
  32. ^ "Radiation and Radioactive Decay. Radioactive Human Body". Harvard Natural Sciences Lecture Demonstrations. Archived from the original on May 28, 2023. Retrieved July 2, 2016.
  33. ^ Winteringham, F. P. W, Effects, F.A.O. Standing Committee on Radiation, Land And Water Development Division, Food and Agriculture Organization of the United Nations (1989). Radioactive fallout in soils, crops and food: a background review. Food & Agriculture Org. p. 32. ISBN 978-92-5-102877-3.{{cite book}}: CS1 maint: multiple names: authors list (link)
  34. ^ du Monceau HL (1702–1797). "Sur la Base de Sel Marin". Mémoires de l'Académie Royale des Sciences (in French): 65–68. Archived from the original on 2019-08-21. Retrieved 2011-05-09.
  35. ^ a b Weeks ME (1932). "The discovery of the elements. IX. Three alkali metals: Potassium, sodium, and lithium". Journal of Chemical Education. 9 (6): 1035. Bibcode:1932JChEd...9.1035W. doi:10.1021/ed009p1035.
  36. ^ a b c Siegfried R (1963). "The Discovery of Potassium and Sodium, and the Problem of the Chemical Elements". Isis. 54 (2): 247–258. doi:10.1086/349704. JSTOR 228541. PMID 14147904. S2CID 38152048.
  37. ^ Browne CA (1926). "Historical notes upon the domestic potash industry in early colonial and later times". Journal of Chemical Education. 3 (7): 749–756. Bibcode:1926JChEd...3..749B. doi:10.1021/ed003p749.
  38. ^ Liebig, Justus von (1840). Die organische Chemie in ihrer Anwendung auf Agricultur und Physiologie (in German). F. Vieweg und Sohn.
  39. ^ Cordel, Oskar (1868). Die Stassfurter Kalisalze in der Landwirthschalt: Eine Besprechung ... (in German). L. Schnock.
  40. ^ Birnbaum K (1869). Die Kalidüngung in ihren Vortheilen und Gefahren (in German).
  41. ^ United Nations Industrial Development Organization and Int'l Fertilizer Development Center (1998). Fertilizer Manual. Springer. pp. 46, 417. ISBN 978-0-7923-5032-3.
  42. ^ Miller H (1980). "Potash from Wood Ashes: Frontier Technology in Canada and the United States". Technology and Culture. 21 (2): 187–208. doi:10.2307/3103338. JSTOR 3103338. S2CID 112819807.
  43. ^ Rittenhouse PA (1979). "Potash and politics". Economic Geology. 74 (2): 353–7. Bibcode:1979EcGeo..74..353R. doi:10.2113/gsecongeo.74.2.353.
  44. ^ Enghag P (2004). "11. Sodium and Potassium". Encyclopedia of the elements. Wiley-VCH Weinheim. ISBN 978-3-527-30666-4.
  45. ^ Davy H (1808). "On some new phenomena of chemical changes produced by electricity, in particular the decomposition of the fixed alkalies, and the exhibition of the new substances that constitute their bases; and on the general nature of alkaline bodies". Philosophical Transactions of the Royal Society. 98: 1–44. doi:10.1098/rstl.1808.0001.
  46. ^ Shaposhnik VA (2007). "History of the discovery of potassium and sodium (on the 200th anniversary of the discovery of potassium and sodium)". Journal of Analytical Chemistry. 62 (11): 1100–2. doi:10.1134/S1061934807110160. S2CID 96141217.
  47. ^ a b c Holleman AF, Wiberg E, Wiberg N (1985). "Potassium". Lehrbuch der Anorganischen Chemie (in German) (91–100 ed.). Walter de Gruyter. ISBN 978-3-11-007511-3.
  48. ^ Burkhardt, p. 32
  49. ^ Rieke, R. D. (1989). "Preparation of Organometallic Compounds from Highly Reactive Metal Powders". Science. 246 (4935): 1260–4. Bibcode:1989Sci...246.1260R. doi:10.1126/science.246.4935.1260. PMID 17832221. S2CID 92794.
  50. ^ Shimansky V, Bikmaev IF, Galeev AI, Shimanskaya NN, Ivanova DV, Sakhibullin NA, Musaev FA, Galazutdinov GA (September 2003). "Observational constraints on potassium synthesis during the formation of stars of the Galactic disk". Astronomy Reports. 47 (9): 750–762. Bibcode:2003ARep...47..750S. doi:10.1134/1.1611216. S2CID 120396773.
  51. ^ The LS, Eid MF, Meyer BS (2000). "A New Study of s-Process Nucleosynthesis in Massive Stars". The Astrophysical Journal. 533 (2): 998. arXiv:astro-ph/9812238. Bibcode:2000ApJ...533..998T. doi:10.1086/308677. ISSN 0004-637X. S2CID 7698683.
  52. ^ Greenwood, p. 69
  53. ^ a b Micale G, Cipollina A, Rizzuti L (2009). Seawater Desalination: Conventional and Renewable Energy Processes. Springer. p. 3. ISBN 978-3-642-01149-8.
  54. ^ a b c d Prud'homme M, Krukowski ST (2006). "Potash". Industrial minerals & rocks: commodities, markets, and uses. Society for Mining, Metallurgy, and Exploration. pp. 723–740. ISBN 978-0-87335-233-8.
  55. ^ Ross WH (1914). "The Origin of Nitrate Deposits". Popular Science. Bonnier Corporation. pp. 134–145.
  56. ^ Garrett DE (1995-12-31). Potash: deposits, processing, properties and uses. Springer. ISBN 978-0-412-99071-7.
  57. ^ a b Ober JA. "Mineral Commodity Summaries 2008:Potash" (PDF). United States Geological Survey. Archived (PDF) from the original on 2019-01-11. Retrieved 2008-11-20.
  58. ^ a b c Ober JA. "Mineral Yearbook 2006:Potash" (PDF). United States Geological Survey. Archived (PDF) from the original on 2008-12-17. Retrieved 2008-11-20.
  59. ^ Wishart DJ (2004). Encyclopedia of the Great Plains. U of Nebraska Press. p. 433. ISBN 978-0-8032-4787-1.
  60. ^ Chiu KW (2000). "Potassium". Kirk-Othmer Encyclopedia of Chemical Technology. John Wiley & Sons, Inc. doi:10.1002/0471238961.1615200103080921.a01.pub2. ISBN 978-0-471-23896-6.
  61. ^ Burkhardt, p. 34
  62. ^ Delahunt J, Lindeman T (2007). "Review of the safety of potassium and potassium oxides, including deactivation by introduction into water". Journal of Chemical Health and Safety. 14 (2): 21–32. doi:10.1016/j.jchas.2006.09.010.
  63. ^ Roy AH (2007). Kent and Riegel's handbook of industrial chemistry and biotechnology. Springer. pp. 1135–57. Bibcode:2007karh.book....... ISBN 978-0-387-27843-8.
  64. ^ Ochoa-Hueso R, Delgado-Baquerizo M, King P, Benham M, Arca V, Power SA (2019). "Ecosystem type and resource quality are more important than global change drivers in regulating early stages of litter decomposition". Soil Biology and Biochemistry. 129: 144–152. Bibcode:2019SBiBi.129..144O. doi:10.1016/j.soilbio.2018.11.009. hdl:10261/336676. S2CID 92606851.
  65. ^ "Potassium Uses, Side Effects & Interactions". Drugs.com. Archived from the original on 2022-04-30. Retrieved 2022-04-30.
  66. ^ World Health Organization (2009). Stuart MC, Kouimtzi M, Hill SR (eds.). WHO Model Formulary 2008. World Health Organization. p. 491. hdl:10665/44053. ISBN 978-92-4-154765-9.
  67. ^ "Potassium chloride medical facts from Drugs.com". www.drugs.com. Archived from the original on 18 January 2017. Retrieved 14 January 2017.
  68. ^ a b British national formulary : BNF 69 (69 ed.). British Medical Association. 2015. pp. 680, 684. ISBN 978-0-85711-156-2.
  69. ^ Figoni, Paula I (2010). "Bleaching and Maturing Agents". How Baking Works: Exploring the Fundamentals of Baking Science. John Wiley and Sons. p. 86. ISBN 978-0-470-39267-6.
  70. ^ Chichester, C. O. (July 1986). "Uses and Exposure to Sulfites in Food". Advances in food research. Academic Press. pp. 4–6. ISBN 978-0-12-016430-1.
  71. ^ Schultz
  72. ^ Toedt, John, Koza, Darrell, Cleef-Toedt, Kathleen Van (2005). "Personal Cleansing Products: Bar Soap". Chemical composition of everyday products. Greenwood Publishing Group. ISBN 978-0-313-32579-3.
  73. ^ Schultz, p. 95
  74. ^ Schultz, p. 99
  75. ^ Siegel RS (1940). "Ignition of the safety match". Journal of Chemical Education. 17 (11): 515. Bibcode:1940JChEd..17..515S. doi:10.1021/ed017p515.
  76. ^ Anger G, Halstenberg J, Hochgeschwender K, Scherhag C, Korallus U, Knopf H, Schmidt P, Ohlinger M. "Chromium Compounds". Ullmann's Encyclopedia of Industrial Chemistry. Vol. 9. Weinheim: Wiley-VCH. p. 178. doi:10.1002/14356007.a07_067. ISBN 978-3527306732.
  77. ^ Greenwood, p. 74
  78. ^ Marx RF (1990). The history of underwater exploration. Courier Dover Publications. p. 93. ISBN 978-0-486-26487-5.
  79. ^ Gettens, Rutherford John, Stout, George Leslie (1966). Painting materials: A short encyclopaedia. Courier Dover Publications. pp. 109–110. ISBN 978-0-486-21597-6. Archived from the original on 2023-01-19. Retrieved 2016-01-08.
  80. ^ Modugno G, Benkő C, Hannaford P, Roati G, Inguscio M (1999-11-01). "Sub-Doppler laser cooling of fermionic ${}^{40}\mathrm{K}$ atoms". Physical Review A. 60 (5): R3373–R3376. arXiv:cond-mat/9908102. Bibcode:1999PhRvA..60.3373M. doi:10.1103/PhysRevA.60.R3373. S2CID 119001675.
  81. ^ Jackson CB, Werner RC (1957). "Ch. 18: The Manufacture of Potassium and NaK". Handling and uses of the alkali metals. Advances in Chemistry. Vol. 19. pp. 169–173. doi:10.1021/ba-1957-0019.ch018. ISBN 978-0-8412-0020-3.
  82. ^ Kearey, Philip, Brooks, M, Hill, Ian (2002). "Optical Pumped Magnetometer". An introduction to geophysical exploration. Wiley-Blackwell. p. 164. ISBN 978-0-632-04929-5.
  83. ^ Abdel-Wahab M, Youssef S, Aly A, el-Fiki S, el-Enany N, Abbas M (1992). "A simple calibration of a whole-body counter for the measurement of total body potassium in humans". International Journal of Radiation Applications and Instrumentation A. 43 (10): 1285–9. doi:10.1016/0883-2889(92)90208-V. PMID 1330980.
  84. ^ Chang, Raymond (2007). Chemistry. McGraw-Hill Higher Education. p. 52. ISBN 978-0-07-110595-8.
  85. ^ Vašák M, Schnabl J (2016). "Chapter 8. Sodium and Potassium Ions in Proteins and Enzyme Catalysis". In Astrid S, Helmut S, Roland KO S (eds.). The Alkali Metal Ions: Their Role in Life. Metal Ions in Life Sciences. Vol. 16. Springer. pp. 259–290. doi:10.1007/978-3-319-21756-7_8. ISBN 978-3-319-21755-0. PMID 26860304.
  86. ^ Weiner ID, Linus S, Wingo CS (2014). "Disorders of potassium metabolism". In Freehally J, Johnson RJ, Floege J (eds.). Comprehensive clinical nephrology (5th ed.). St. Louis: Saunders. p. 118. ISBN 978-0-323-24287-5.
  87. ^ Malnic G, Giebisch G, Muto S, Wang W, Bailey MA, Satlin LM (2013). "Regulation of K+ excretion". In Alpern RJ, Caplan MJ, Moe OW (eds.). Seldin and Giebisch's the kidney: physiology and pathophysiology (5th ed.). London: Academic Press. pp. 1659–1716. ISBN 978-0-12-381463-0.
  88. ^ Mount DB, Zandi-Nejad K (2011). "Disorders of potassium balance". In Taal MW, Chertow GM, Marsden PA, Skorecki KL, Yu AS, Brenner BM (eds.). The kidney (9th ed.). Philadelphia: Elsevier. pp. 640–688. ISBN 978-1-4557-2304-1.
  89. ^ Lockless SW, Zhou M, MacKinnon R (2007). "Structural and thermodynamic properties of selective ion binding in a K+ channel". PLOS Biol. 5 (5): e121. doi:10.1371/journal.pbio.0050121. PMC 1858713. PMID 17472437.
  90. ^ Dietz AG, Weikop P, Hauglund N, Andersen M, Petersen NC, Rose L, Hirase H, Nedergaard M (2023). "Local extracellular K + in cortex regulates norepinephrine levels, network state, and behavioral output". Proceedings of the National Academy of Sciences. 120 (40): e2305071120. Bibcode:2023PNAS..12005071D. doi:10.1073/pnas.2305071120. ISSN 0027-8424. PMC 10556678. PMID 37774097.
  91. ^ Wei KY, Gritter M, Vogt L, de Borst MH, Rotmans JI, Hoorn EJ (2020-09-02). "Dietary potassium and the kidney: lifesaving physiology". Clinical Kidney Journal. 13 (6). Oxford University Press (OUP): 952–968. doi:10.1093/ckj/sfaa157. ISSN 2048-8513. PMC 7769543. PMID 33391739.
  92. ^ Goyal A, Spertus JA, Gosch K, Venkitachalam L, Jones PG, Van den Berghe G, Kosiborod M (2012). "Serum Potassium Levels and Mortality in Acute Myocardial Infarction". JAMA. 307 (2): 157–164. doi:10.1001/jama.2011.1967. PMID 22235086.
  93. ^ Smyth A, Dunkler D, Gao P, et al. (2014). "The relationship between estimated sodium and potassium excretion and subsequent renal outcomes". Kidney Int. 86 (6): 1205–1212. doi:10.1038/ki.2014.214. PMID 24918156.
  94. ^ Moore-Ede MC (1986). "Physiology of the circadian timing system: predictive versus reactive homeostasis". Am J Physiol. 250 (5 Pt 2): R737–R752. doi:10.1152/ajpregu.1986.250.5.R737. PMID 3706563.
  95. ^ Slonim AD, Pollack MM (2006). "Potassium". Pediatric critical care medicine. Lippincott Williams & Wilkins. p. 812. ISBN 978-0-7817-9469-5.
  96. ^ Visveswaran K (2009). "hypokalemia". Essentials of Nephrology (2nd ed.). BI Publications. p. 257. ISBN 978-81-7225-323-3.
  97. ^ Gumz ML, Rabinowitz L, Wingo CS (2015-07-02). "An Integrated View of Potassium Homeostasis". The New England Journal of Medicine. 373 (1): 60–72. doi:10.1056/NEJMra1313341. ISSN 0028-4793. PMC 5675534. PMID 26132942.
  98. ^ Campbell N (1987). Biology. Menlo Park, California: Benjamin/Cummings Pub. Co. p. 795. ISBN 978-0-8053-1840-1.
  99. ^ Hellgren M, Sandberg L, Edholm O (2006). "A comparison between two prokaryotic potassium channels (KirBac1.1 and KcsA) in a molecular dynamics (MD) simulation study". Biophysical Chemistry. 120 (1): 1–9. doi:10.1016/j.bpc.2005.10.002. PMID 16253415.
  100. ^ Potts, W. T. W., Parry, G. (1964). Osmotic and ionic regulation in animals. Pergamon Press.
  101. ^ Lans H', Stein IF, Meyer KA (1952). "The relation of serum potassium to erythrocyte potassium in normal subjects and patients with potassium deficiency". American Journal of the Medical Sciences. 223 (1): 65–74. doi:10.1097/00000441-195201000-00011. PMID 14902792.
  102. ^ Bennett CM, Brenner BM, Berliner RW (1968). "Micropuncture study of nephron function in the rhesus monkey". Journal of Clinical Investigation. 47 (1): 203–216. doi:10.1172/JCI105710. PMC 297160. PMID 16695942.
  103. ^ Solomon AK (1962). "Pumps in the living cell". Scientific American. 207 (2): 100–8. Bibcode:1962SciAm.207b.100S. doi:10.1038/scientificamerican0862-100. PMID 13914986.
  104. ^ Kernan RP (1980). Cell potassium (Transport in the life sciences). New York: Wiley. pp. 40, 48. ISBN 978-0-471-04806-0.
  105. ^ Wright F' (1977). "Sites and mechanisms of potassium transport along the renal tubule". Kidney International. 11 (6): 415–432. doi:10.1038/ki.1977.60. PMID 875263.
  106. ^ Squires RD, Huth EJ (1959). "Experimental potassium depletion in normal human subjects. I. Relation of ionic intakes to the renal conservation of potassium". Journal of Clinical Investigation. 38 (7): 1134–48. doi:10.1172/JCI103890. PMC 293261. PMID 13664789.
  107. ^ Fiebach, Nicholas H., Barker, Lee Randolph, Burton, John Russell, Zieve, Philip D. (2007). Principles of ambulatory medicine. Lippincott Williams & Wilkins. pp. 748–750. ISBN 978-0-7817-6227-4.
  108. ^ Gadsby DC (2004). "Ion transport: spot the difference". Nature. 427 (6977): 795–7. Bibcode:2004Natur.427..795G. doi:10.1038/427795a. PMID 14985745. S2CID 5923529.; for a diagram of the potassium pores are viewed, see Miller, C (2001). "See potassium run". Nature. 414 (6859): 23–24. Bibcode:2001Natur.414...23M. doi:10.1038/35102126. PMID 11689922. S2CID 4423041.
  109. ^ Jiang Y, Lee A, Chen J, Cadene M, Chait B', Mackinnon R (2002). "Crystal structure and mechanism of a calcium-gated potassium channel" (PDF). Nature. 417 (6888): 515–22. Bibcode:2002Natur.417..515J. doi:10.1038/417515a. PMID 12037559. S2CID 205029269. Archived from the original (PDF) on 2009-04-24.
  110. ^ Shi N, Ye S, Alam A, Chen L, Jiang Y (2006). "Atomic structure of a Na+- and K+-conducting channel". Nature. 440 (7083): 570–4. Bibcode:2006Natur.440..570S. doi:10.1038/nature04508. PMID 16467789. S2CID 4355500; includes a detailed picture of atoms in the pump.
  111. ^ Zhou Y, Morais-Cabral JH, Kaufman A, MacKinnon R (2001). "Chemistry of ion coordination and hydration revealed by a K+ channel-Fab complex at 2.0 A resolution" (PDF). Nature. 414 (6859): 43–48. Bibcode:2001Natur.414...43Z. doi:10.1038/35102009. PMID 11689936. S2CID 205022645. Archived from the original (PDF) on Oct 17, 2021.
  112. ^ Noskov SY, Bernèche S, Roux B (2004). "Control of ion selectivity in potassium channels by electrostatic and dynamic properties of carbonyl ligands" (PDF). Nature. 431 (7010): 830–4. Bibcode:2004Natur.431..830N. doi:10.1038/nature02943. PMID 15483608. S2CID 4414885. Archived (PDF) from the original on Mar 26, 2023.
  113. ^ National Academies of Sciences, Engineering and Medicine (2019). "Potassium: Dietary Reference Intakes for Adequacy". In Stallings VA, Harrison M, Oria M (eds.). Dietary Reference Intakes for Sodium and Potassium. Washington, DC: The National Academies Press. doi:10.17226/25353. ISBN 978-0-309-48834-1. PMID 30844154. Archived from the original on 2019-05-13. Retrieved 2019-05-13.
  114. ^ Stallings VA, Harrison M, Oria M, eds. (March 5, 2019). Dietary Reference Intakes for Sodium and Potassium – Publication. National Academies of Sciences, Engineering and Medicine. doi:10.17226/25353. ISBN 978-0-309-48834-1. PMID 30844154. S2CID 104464967. Archived from the original on May 9, 2019. Retrieved May 13, 2019. {{cite book}}: |website= ignored (help)
  115. ^ Panel on Dietary Reference Intakes for Electrolytes and Water, Standing Committee on the Scientific Evaluation of Dietary Reference Intakes, Food and Nutrition (2004). DRI, dietary reference intakes for water, potassium, sodium, chloride, and sulfate. Washington, D.C.: National Academies Press. ISBN 978-0-309-53049-1. Archived from the original on 2011-10-06.{{cite book}}: CS1 maint: multiple names: authors list (link)
  116. ^ Karger S (2004). "Energy and nutrient intake in the European Union". Annals of Nutrition and Metabolism. 48 (2 (suppl)): 1–16. doi:10.1159/000083041.
  117. ^ "Vitamins and minerals". National Health Service (NHS). 18 November 2021. Archived from the original on 3 April 2019. Retrieved 13 November 2022.
  118. ^ "Potassium Food Charts". Asia Pacific Journal of Clinical Nutrition. Archived from the original on 2021-04-29. Retrieved 2011-05-18.
  119. ^ "Potassium Content of Selected Foods per Common Measure, sorted by nutrient content" (PDF). USDA National Nutrient Database for Standard Reference, Release 20. Archived from the original (PDF) on December 17, 2008.
  120. ^ Whelton PK, He J, Cutler JA, Brancati FL, Appel LJ, Follmann D, Klag MJ (1997). "Effects of oral potassium on blood pressure. Meta-analysis of randomized controlled clinical trials". JAMA. 277 (20): 1624–32. doi:10.1001/jama.1997.03540440058033. PMID 9168293. S2CID 25937399.
  121. ^ Krishna GG, Miller E, Kapoor S (1989). "Increased blood pressure during potassium depletion in normotensive men". The New England Journal of Medicine. 320 (18): 1177–82. doi:10.1056/NEJM198905043201804. PMID 2624617.
  122. ^ a b c Soar J, Perkins GD, Abbas G, Alfonzo A, Barelli A, Bierens JJ, Brugger H, Deakin CD, Dunning J, Georgiou M, Handley AJ, Lockey DJ, Paal P, Sandroni C, Thies KC, Zideman DA, Nolan JP (October 2010). "European Resuscitation Council Guidelines for Resuscitation 2010 Section 8. Cardiac arrest in special circumstances: Electrolyte abnormalities, poisoning, drowning, accidental hypothermia, hyperthermia, asthma, anaphylaxis, cardiac surgery, trauma, pregnancy, electrocution". Resuscitation. 81 (10): 1400–33. doi:10.1016/j.resuscitation.2010.08.015. PMID 20956045.
  123. ^ Horisberger JD, Giebisch G (1987). "Potassium-Sparing Diuretics". Kidney and Blood Pressure Research. 10 (3–4): 198–220. doi:10.1159/000173130. ISSN 1420-4096. PMID 2455308.
  124. ^ Ng K, Lee CS (February 16, 2017). "Updated Treatment Options in the Management of Hyperkalemia". US Pharmacist. Retrieved 2024-07-19.
  125. ^ a b Institute of Medicine (U.S.). Committee on Optimization of Nutrient Composition of Military Rations for Short-Term, High-Stress Situations, Institute of Medicine (U.S.). Committee on Military Nutrition Research (2006). Nutrient composition of rations for short-term, high-intensity combat operations. National Academies Press. pp. 287–. ISBN 978-0-309-09641-6.
  126. ^ Benge CD, Burka AT (2020). "Oral Liquid Potassium Chloride Dosing Pathway in a Tertiary Care Veteran Affairs Academic Medical Center". Critical Pathways in Cardiology: A Journal of Evidence-Based Medicine. 19 (1): 18–21. doi:10.1097/HPC.0000000000000197. ISSN 1535-282X. PMID 31478945.
  127. ^ Gueta I, Markovits N, Halkin H, Loebstein R (2021). "Concomitant oral potassium chloride and anticholinergic therapy is associated with upper gastrointestinal bleeding: A cohort study". British Journal of Clinical Pharmacology. 87 (4): 2064–2069. doi:10.1111/bcp.14616. ISSN 0306-5251. PMID 33068044.
  128. ^ "Potassium - Fact Sheet for Health Professionals". National Institutes of Health (NIH) Office of Dietary Supplements (ODS). June 2, 2022. Retrieved 2024-07-19.
  129. ^ a b c d e f g h Keller CL, Jones NT, Abadie RB, Barham W, Behara R, Patil S, Paladini A, Ahmadzadeh S, Shekoohi S, Varrassi G, Kaye AD (January 2024). "Non-steroidal Anti-inflammatory Drug (NSAID)-, Potassium Supplement-, Bisphosphonate-, and Doxycycline-Mediated Peptic Ulcer Effects: A Narrative Review". Cureus. 16 (1): e51894. doi:10.7759/cureus.51894. PMC 10849936. PMID 38333496.
  130. ^ D'Elia L, Barba G, Cappuccio F, Strazzullo (2011). "Potassium Intake, Stroke, and Cardiovascular Disease: A Meta-Analysis of Prospective Studies". J Am Coll Cardiol. 57 (10): 1210–9. doi:10.1016/j.jacc.2010.09.070. PMID 21371638.
  131. ^ He FJ, Marciniak M, Carney C, Markandu ND, Anand V, Fraser WD, Dalton RN, Kaski JC, MacGregor GA (2010). "Effects of potassium chloride and potassium bicarbonate on endothelial function, cardiovascular risk factors, and bone turnover in mild hypertensives". Hypertension. 55 (3): 681–8. doi:10.1161/HYPERTENSIONAHA.109.147488. PMID 20083724.
  132. ^ "The Top 300 of 2020". ClinCalc. Archived from the original on 12 February 2021. Retrieved 7 October 2022.
  133. ^ "Potassium Chloride - Drug Usage Statistics". ClinCalc. Archived from the original on 8 October 2022. Retrieved 7 October 2022.
  134. ^ Shallenberger RS (1993). Taste chemistry. Springer. pp. 120–. ISBN 978-0-7514-0150-9.
  135. ^ "Potassium 244856". Sigma Aldrich. Archived from the original on 2018-10-01. Retrieved 2018-10-01.
  136. ^ Solomon RE (2002). Fire and Life Safety Inspection Manual. Jones & Bartlett Learning. p. 459. ISBN 978-0-87765-472-8.
  137. ^ "DOE Handbook-Alkali Metals Sodium, Potassium, NaK, and Lithium". Hss.doe.gov. Archived from the original on 2010-09-28. Retrieved 2010-10-16.
  138. ^ Wray TK. "Danger: peroxidazable chemicals" (PDF). Environmental Health & Public Safety, North Carolina State University. Archived from the original on 2016-07-29.{{cite web}}: CS1 maint: bot: original URL status unknown (link)
  139. ^ a b Schonwald S (2004). "Potassium Chloride and Potassium Permanganate". Medical toxicology. Lippincott Williams & Wilkins. pp. 903–5. ISBN 978-0-7817-2845-4.
  140. ^ Markovchick VJ, Pons PT (2003). Emergency medicine secrets. Elsevier Health Sciences. p. 223. ISBN 978-1-56053-503-4.

Bibliography

Retrieved from "https://en.wikipedia.org/w/index.php?title=Potassium&oldid=1252623230"