Калий химический элемент как пишется

Potassium, ₁₉K

Potassium pearls (in paraffin oil, ~5 mm each)
Potassium
Pronunciation	​(pə-TASS-ee-əm)
Appearance	silvery white, faint bluish-purple hue when exposed to air
Standard atomic weight Ar°(K)
	39.0983±0.0001 39.098±0.001 (abridged)[1]
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
Group	group 1: hydrogen and alkali metals
Period	period 4
Block	s-block
Electron configuration	[Ar] 4s1
Electrons per shell	2, 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)[2]
Density (near r.t.)	0.89 g/cm3
when liquid (at m.p.)	0.82948 g/cm3[2]
Critical point	2223 K, 16 MPa[3]
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
Spectral lines of potassium
Other properties
Natural occurrence	primordial
Crystal structure	​body-centered cubic (bcc)
Speed of sound thin rod	2000 m/s (at 20 °C)
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[4]
Molar magnetic susceptibility	+20.8×10−6 cm3/mol (298 K)[5]
Young’s modulus	3.53 GPa
Shear modulus	1.3 GPa
Bulk modulus	3.1 GPa
Mohs hardness	0.4
Brinell hardness	0.363 MPa
CAS Number	7440-09-7
History
Discovery and first isolation	Humphry Davy (1807)
Symbol	«K»: from New Latin kalium
Isotopes of potassium v e
Main isotopes Decay abun­dance half-life (t1/2) mode pro­duct 39K 93.258% stable 40K 0.012% 1.248×109 y β− 40Ca ε 40Ar β+ 40Ar 41K 6.730% stable
Category: Potassium view talk edit | references

Potassium is the chemical element with the 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.^[6] 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),^[7][8] and occurs in many minerals such as orthoclase, a common constituent of granites and other igneous rocks.^[9]

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 suspected in 1702 that they were distinct elements that combine with the same anions to make similar salts,^[10] and this was proven in 1807 through using 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.^[11]

Etymology

The English name for the element potassium comes from the word potash,^[12] 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.^[13] In 1807, Humphry Davy produced the element via electrolysis: in 1809, Ludwig Wilhelm Gilbert proposed the name Kalium for Davy’s «potassium».^[14] In 1814, the Swedish chemist Berzelius advocated the name kalium for potassium, with the chemical symbol K.^[15]

The English and French-speaking countries adopted Davy and Gay-Lussac/Thénard’s name Potassium, whereas the Germanic countries adopted Gilbert/Klaproth’s name Kalium.^[16] The «Gold Book» of the International Union of Pure and Applied Chemistry has designated the official chemical symbol as K.^[17]

Properties

Physical

Potassium is the second least dense metal after lithium. It is a soft solid with a low melting point, and can be easily cut with a knife. Freshly cut potassium is silvery in appearance, but it begins to tarnish toward gray immediately on exposure to air.^[18] In a flame test, potassium and its compounds emit a lilac color with a peak emission wavelength of 766.5 nanometers.^[19]

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.^[20] 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.^[21]

Compounds

Four oxides of potassium are well studied: potassium oxide (K₂O), potassium peroxide (K₂O₂), potassium superoxide (KO₂)^[22] 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.^[23][24] 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.^[25]

Potassium heptafluorotantalate (K₂[TaF₇]) is an intermediate in the purification of tantalum from the otherwise persistent contaminant of niobium.^[26]

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

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%).^[27] 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.^[27][28] 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^[29] 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.^[30] The activity of natural potassium is 31 Bq/g.^[31]

Cosmic formation and distribution

Potassium is formed in supernovae by nucleosynthesis from lighter atoms. Potassium is principally created in Type II supernovae via an explosive oxygen-burning process.^[32] (These are fusion reactions; do not confuse with chemical burning between potassium and oxygen.) ⁴⁰
K is also formed in s-process nucleosynthesis and the neon burning process.^[33]

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.^[34] The potassium concentration in seawater is 0.39 g/L^[7] (0.039 wt/v%), about one twenty-seventh the concentration of sodium.^[35][36]

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,^[10] and Henri Louis Duhamel du Monceau was able to prove this difference in 1736.^[37] 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.^[38][39] For a long time the only significant applications for potash were the production of glass, bleach, soap and gunpowder as potassium nitrate.^[40] 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.^[11] The discovery by Justus Liebig in 1840 that potassium is a necessary element for plants and that most types of soil lack potassium^[41] 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.^[42][43][44] Other potash deposits were discovered, and by the 1960s Canada became the dominant producer.^[45][46]

Metal

Potassium metal was first isolated in 1807 by Humphry Davy, who derived it by electrolysis of molten KOH with the newly discovered voltaic pile. Potassium was the first metal that was isolated by electrolysis.^[47] 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.^{[38][39][48][49]} 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.^[39]

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.^[50] 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.^[51] 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.^[52] Illustrative is the preparation of magnesium:

MgCl₂ + 2 K → Mg + 2 KCl

Geology

Elemental potassium does not occur in nature because of its high reactivity. It reacts violently with water (see section Precautions below)^[50] 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.^[36] 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.^[53]

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.^[54] 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).^[55] Potassium ions are present in a wide variety of proteins and enzymes.^[56]

Biochemical function

Potassium levels influence multiple physiological processes, including^[57][58][59]

• 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.^[60]
• 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

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.^[61] Levels outside this range are associated with an increasing rate of death from multiple causes,^[62] and some cardiac, kidney,^[63] 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.^[64]

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).^[65] Deficiency symptoms include muscle weakness, paralytic ileus, ECG abnormalities, decreased reflex response; and in severe cases, respiratory paralysis, alkalosis, and cardiac arrhythmia.^[66]

Control mechanisms

Potassium content in the plasma is tightly controlled by four basic mechanisms, which have various names and classifications. The four 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»;^[67] 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.^[68] 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.^[69]

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)], while sodium is the major cation of extracellular fluid [150 mmol/L, (3.345 g)]. In the kidneys, about 180 liters of plasma is filtered through the glomeruli and into the renal tubules per day.^[70] This filtering involves about 600 g of sodium and 33 g of potassium. Since only 1–10 g of sodium and 1–4 g 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.^[71] 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^[72][73] in counter flow to sodium in response to an apparent (but not actual) Donnan equilibrium,^[74] 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.^[75] 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.^[76] If potassium is still withheld, the concentration continues to fall until a severe deficiency causes eventual death.^[77]

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 channel have only one gate, and there only one kind of ion can stream through, at 10 million to 100 million ions per second.^[78] Calcium is required to open the pores,^[79] although calcium may work in reverse by blocking at least one of the pores.^[80] Carbonyl groups inside the pore on the amino acids mimic the water hydration that takes place in water solution^[81] by the nature of the electrostatic charges on four carbonyl groups inside the pore.^[82]

Nutrition

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 0-6-month-old infants, 860 mg of potassium for 7-12-month-old infants, 2,000 mg of potassium for 1-3-year-old children, and 2,300 mg of potassium for 4-8-year-old children.

For males 9 years of age and older, the AIs for potassium are: 2,500 mg of potassium for 9-13-year-old males, 3,000 mg of potassium for 14-18-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 9-18-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 14-18-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 14-18-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.^[83][84]

As of 2004, most Americans adults consume less than 3,000 mg.^[85]

Likewise, in the European Union, in particular in Germany, and Italy, insufficient potassium intake is somewhat common.^[86] 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.^[87]

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.^[88][89]

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.^[90]

The United States Department of Agriculture lists tomato paste, orange juice, beet greens, white beans, potatoes, plantains, bananas, apricots, 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.^[91]

Deficient intake

Diets low in potassium can lead to hypertension^[92] and hypokalemia.

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.^{[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.^[93] 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.^{[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]}

A meta-analysis concluded that a 1640 mg increase in the daily intake of potassium was associated with a 21% lower risk of stroke.^[94] Potassium chloride and potassium bicarbonate may be useful to control mild hypertension.^[95] In 2020, potassium was the 33rd most commonly prescribed medication in the U.S., with more than 17 million prescriptions.^[96][97]

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.^[93][98]

Commercial production

Mining

Potassium salts such as carnallite, langbeinite, polyhalite, and sylvite form extensive evaporite deposits in ancient lake bottoms and seabeds,^[35] 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.^{[99][100][101]} 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.^[102] 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.^[100][101]

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.^[36]

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.^[36][103]

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.^[104][105]

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.^[50]
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 ions are an essential component of plant nutrition and are found in most soil types.^[11] 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.^[11] 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.^[106] In 2005, about 93% of world potassium production was consumed by the fertilizer industry.^[101] Furthermore, potassium can play a key role in nutrient cycling by controlling litter composition.^[107]

Medical use

Potassium citrate

Potassium citrate is used to treat a kidney stone condition called renal tubular acidosis.^[108]

Potassium chloride

Potassium, in the form of potassium chloride is used as a medication to treat and prevent low blood potassium.^[109] Low blood potassium may occur due to vomiting, diarrhea, or certain medications.^[110] It is given by slow injection into a vein or by mouth.^[111]

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.^[112][113]

Industrial

Major potassium chemicals are potassium hydroxide, potassium carbonate, potassium sulfate, and potassium chloride. Megatons of these compounds are produced annually.^[114]

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.^[115][116]

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.^[117] 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.^[11]

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,^[118] 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.^[119]

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.^[120][121]

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.^[122]

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.^[123]

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.^[124] The ternary alloy of 12% Na, 47% K and 41% Cs has the lowest melting point of −78 °C of any metallic compound.^[18]

Metallic potassium is used in several types of magnetometers.^[125]

Precautions

Potassium

Hazards
GHS labelling:
Pictograms
Signal word	Danger
Hazard statements	H260, H314
Precautionary statements	P223, P231+P232, P280, P305+P351+P338, P370+P378, P422[126]
NFPA 704 (fire diamond)	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)↑

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.^[127]

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.^[128]

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.^[129]

Ingestion of large amounts of potassium compounds can lead to hyperkalemia, strongly influencing the cardiovascular system.^[130][131] Potassium chloride is used in the U.S. for lethal injection executions.^[130]

See also

References

Bibliography

• Burkhardt, Elizabeth R. (2006). «Potassium and Potassium Alloys». Ullmann’s Encyclopedia of Industrial Chemistry. Vol. A22. pp. 31–38. doi:10.1002/14356007.a22_031.pub2. ISBN 978-3-527-30673-2.
• Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
• Holleman, Arnold F.; Wiberg, Egon; Wiberg, Nils (2007). «Potassium». Lehrbuch der Anorganischen Chemie (in German) (91–100 ed.). Walter de Gruyter. ISBN 978-3110177701.
• Schultz, H.; et al. (2006). «Potassium compounds». Ullmann’s Encyclopedia of Industrial Chemistry. Vol. A22. pp. 39–103. doi:10.1002/14356007.a22_031.pub2. ISBN 978-3-527-30673-2.
• National Nutrient Database Archived 2014-08-10 at the Wayback Machine at USDA Website

External links

• «Potassium». Drug Information Portal. U.S. National Library of Medicine.

Potassium, ₁₉K

Potassium pearls (in paraffin oil, ~5 mm each)
Potassium
Pronunciation	​(pə-TASS-ee-əm)
Appearance	silvery white, faint bluish-purple hue when exposed to air
Standard atomic weight Ar°(K)
	39.0983±0.0001 39.098±0.001 (abridged)[1]
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
Group	group 1: hydrogen and alkali metals
Period	period 4
Block	s-block
Electron configuration	[Ar] 4s1
Electrons per shell	2, 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)[2]
Density (near r.t.)	0.89 g/cm3
when liquid (at m.p.)	0.82948 g/cm3[2]
Critical point	2223 K, 16 MPa[3]
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
Spectral lines of potassium
Other properties
Natural occurrence	primordial
Crystal structure	​body-centered cubic (bcc)
Speed of sound thin rod	2000 m/s (at 20 °C)
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[4]
Molar magnetic susceptibility	+20.8×10−6 cm3/mol (298 K)[5]
Young’s modulus	3.53 GPa
Shear modulus	1.3 GPa
Bulk modulus	3.1 GPa
Mohs hardness	0.4
Brinell hardness	0.363 MPa
CAS Number	7440-09-7
History
Discovery and first isolation	Humphry Davy (1807)
Symbol	«K»: from New Latin kalium
Isotopes of potassium v e
Main isotopes Decay abun­dance half-life (t1/2) mode pro­duct 39K 93.258% stable 40K 0.012% 1.248×109 y β− 40Ca ε 40Ar β+ 40Ar 41K 6.730% stable
Category: Potassium view talk edit | references

Potassium is the chemical element with the 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.^[6] 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),^[7][8] and occurs in many minerals such as orthoclase, a common constituent of granites and other igneous rocks.^[9]

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 suspected in 1702 that they were distinct elements that combine with the same anions to make similar salts,^[10] and this was proven in 1807 through using 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.^[11]

Etymology

The English name for the element potassium comes from the word potash,^[12] 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.^[13] In 1807, Humphry Davy produced the element via electrolysis: in 1809, Ludwig Wilhelm Gilbert proposed the name Kalium for Davy’s «potassium».^[14] In 1814, the Swedish chemist Berzelius advocated the name kalium for potassium, with the chemical symbol K.^[15]

The English and French-speaking countries adopted Davy and Gay-Lussac/Thénard’s name Potassium, whereas the Germanic countries adopted Gilbert/Klaproth’s name Kalium.^[16] The «Gold Book» of the International Union of Pure and Applied Chemistry has designated the official chemical symbol as K.^[17]

Properties

Physical

Potassium is the second least dense metal after lithium. It is a soft solid with a low melting point, and can be easily cut with a knife. Freshly cut potassium is silvery in appearance, but it begins to tarnish toward gray immediately on exposure to air.^[18] In a flame test, potassium and its compounds emit a lilac color with a peak emission wavelength of 766.5 nanometers.^[19]

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.^[20] 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.^[21]

Compounds

Four oxides of potassium are well studied: potassium oxide (K₂O), potassium peroxide (K₂O₂), potassium superoxide (KO₂)^[22] 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.^[23][24] 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.^[25]

Potassium heptafluorotantalate (K₂[TaF₇]) is an intermediate in the purification of tantalum from the otherwise persistent contaminant of niobium.^[26]

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

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%).^[27] 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.^[27][28] 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^[29] 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.^[30] The activity of natural potassium is 31 Bq/g.^[31]

Cosmic formation and distribution

Potassium is formed in supernovae by nucleosynthesis from lighter atoms. Potassium is principally created in Type II supernovae via an explosive oxygen-burning process.^[32] (These are fusion reactions; do not confuse with chemical burning between potassium and oxygen.) ⁴⁰
K is also formed in s-process nucleosynthesis and the neon burning process.^[33]

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.^[34] The potassium concentration in seawater is 0.39 g/L^[7] (0.039 wt/v%), about one twenty-seventh the concentration of sodium.^[35][36]

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,^[10] and Henri Louis Duhamel du Monceau was able to prove this difference in 1736.^[37] 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.^[38][39] For a long time the only significant applications for potash were the production of glass, bleach, soap and gunpowder as potassium nitrate.^[40] 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.^[11] The discovery by Justus Liebig in 1840 that potassium is a necessary element for plants and that most types of soil lack potassium^[41] 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.^[42][43][44] Other potash deposits were discovered, and by the 1960s Canada became the dominant producer.^[45][46]

Metal

Potassium metal was first isolated in 1807 by Humphry Davy, who derived it by electrolysis of molten KOH with the newly discovered voltaic pile. Potassium was the first metal that was isolated by electrolysis.^[47] 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.^{[38][39][48][49]} 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.^[39]

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.^[50] 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.^[51] 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.^[52] Illustrative is the preparation of magnesium:

MgCl₂ + 2 K → Mg + 2 KCl

Geology

Elemental potassium does not occur in nature because of its high reactivity. It reacts violently with water (see section Precautions below)^[50] 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.^[36] 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.^[53]

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.^[54] 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).^[55] Potassium ions are present in a wide variety of proteins and enzymes.^[56]

Biochemical function

Potassium levels influence multiple physiological processes, including^[57][58][59]

• 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.^[60]
• 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

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.^[61] Levels outside this range are associated with an increasing rate of death from multiple causes,^[62] and some cardiac, kidney,^[63] 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.^[64]

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).^[65] Deficiency symptoms include muscle weakness, paralytic ileus, ECG abnormalities, decreased reflex response; and in severe cases, respiratory paralysis, alkalosis, and cardiac arrhythmia.^[66]

Control mechanisms

Potassium content in the plasma is tightly controlled by four basic mechanisms, which have various names and classifications. The four 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»;^[67] 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.^[68] 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.^[69]

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)], while sodium is the major cation of extracellular fluid [150 mmol/L, (3.345 g)]. In the kidneys, about 180 liters of plasma is filtered through the glomeruli and into the renal tubules per day.^[70] This filtering involves about 600 g of sodium and 33 g of potassium. Since only 1–10 g of sodium and 1–4 g 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.^[71] 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^[72][73] in counter flow to sodium in response to an apparent (but not actual) Donnan equilibrium,^[74] 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.^[75] 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.^[76] If potassium is still withheld, the concentration continues to fall until a severe deficiency causes eventual death.^[77]

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 channel have only one gate, and there only one kind of ion can stream through, at 10 million to 100 million ions per second.^[78] Calcium is required to open the pores,^[79] although calcium may work in reverse by blocking at least one of the pores.^[80] Carbonyl groups inside the pore on the amino acids mimic the water hydration that takes place in water solution^[81] by the nature of the electrostatic charges on four carbonyl groups inside the pore.^[82]

Nutrition

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 0-6-month-old infants, 860 mg of potassium for 7-12-month-old infants, 2,000 mg of potassium for 1-3-year-old children, and 2,300 mg of potassium for 4-8-year-old children.

For males 9 years of age and older, the AIs for potassium are: 2,500 mg of potassium for 9-13-year-old males, 3,000 mg of potassium for 14-18-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 9-18-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 14-18-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 14-18-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.^[83][84]

As of 2004, most Americans adults consume less than 3,000 mg.^[85]

Likewise, in the European Union, in particular in Germany, and Italy, insufficient potassium intake is somewhat common.^[86] 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.^[87]

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.^[88][89]

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.^[90]

The United States Department of Agriculture lists tomato paste, orange juice, beet greens, white beans, potatoes, plantains, bananas, apricots, 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.^[91]

Deficient intake

Diets low in potassium can lead to hypertension^[92] and hypokalemia.

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.^{[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.^[93] 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.^{[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]}

A meta-analysis concluded that a 1640 mg increase in the daily intake of potassium was associated with a 21% lower risk of stroke.^[94] Potassium chloride and potassium bicarbonate may be useful to control mild hypertension.^[95] In 2020, potassium was the 33rd most commonly prescribed medication in the U.S., with more than 17 million prescriptions.^[96][97]

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.^[93][98]

Commercial production

Mining

Potassium salts such as carnallite, langbeinite, polyhalite, and sylvite form extensive evaporite deposits in ancient lake bottoms and seabeds,^[35] 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.^{[99][100][101]} 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.^[102] 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.^[100][101]

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.^[36]

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.^[36][103]

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.^[104][105]

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.^[50]
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 ions are an essential component of plant nutrition and are found in most soil types.^[11] 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.^[11] 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.^[106] In 2005, about 93% of world potassium production was consumed by the fertilizer industry.^[101] Furthermore, potassium can play a key role in nutrient cycling by controlling litter composition.^[107]

Medical use

Potassium citrate

Potassium citrate is used to treat a kidney stone condition called renal tubular acidosis.^[108]

Potassium chloride

Potassium, in the form of potassium chloride is used as a medication to treat and prevent low blood potassium.^[109] Low blood potassium may occur due to vomiting, diarrhea, or certain medications.^[110] It is given by slow injection into a vein or by mouth.^[111]

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.^[112][113]

Industrial

Major potassium chemicals are potassium hydroxide, potassium carbonate, potassium sulfate, and potassium chloride. Megatons of these compounds are produced annually.^[114]

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.^[115][116]

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.^[117] 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.^[11]

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,^[118] 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.^[119]

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.^[120][121]

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.^[122]

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.^[123]

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.^[124] The ternary alloy of 12% Na, 47% K and 41% Cs has the lowest melting point of −78 °C of any metallic compound.^[18]

Metallic potassium is used in several types of magnetometers.^[125]

Precautions

Potassium

Hazards
GHS labelling:
Pictograms
Signal word	Danger
Hazard statements	H260, H314
Precautionary statements	P223, P231+P232, P280, P305+P351+P338, P370+P378, P422[126]
NFPA 704 (fire diamond)	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)↑

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.^[127]

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.^[128]

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.^[129]

Ingestion of large amounts of potassium compounds can lead to hyperkalemia, strongly influencing the cardiovascular system.^[130][131] Potassium chloride is used in the U.S. for lethal injection executions.^[130]

See also

References

Bibliography

• Burkhardt, Elizabeth R. (2006). «Potassium and Potassium Alloys». Ullmann’s Encyclopedia of Industrial Chemistry. Vol. A22. pp. 31–38. doi:10.1002/14356007.a22_031.pub2. ISBN 978-3-527-30673-2.
• Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
• Holleman, Arnold F.; Wiberg, Egon; Wiberg, Nils (2007). «Potassium». Lehrbuch der Anorganischen Chemie (in German) (91–100 ed.). Walter de Gruyter. ISBN 978-3110177701.
• Schultz, H.; et al. (2006). «Potassium compounds». Ullmann’s Encyclopedia of Industrial Chemistry. Vol. A22. pp. 39–103. doi:10.1002/14356007.a22_031.pub2. ISBN 978-3-527-30673-2.
• National Nutrient Database Archived 2014-08-10 at the Wayback Machine at USDA Website

External links

• «Potassium». Drug Information Portal. U.S. National Library of Medicine.

К Калий является химическим элементом таблицы Менделеева с атомным номером 19 и обозначением K. Калий представляет собой мягкий щелочной металл серебристо-белого цвета.

Как был открыт Калий

Открытие такого химического элемента как калий принадлежит английскоиу химику Хэмфри Дэви. 19 ноября 1807 года ученый сообщил перед Королевским обществом в Лондоне, что обнаружил два новых химических элемента. Первым являлся Калий, а вторым, который был обнаружен чуть позже был Натрий. Калий Дэви получил методом электролиза из карбоната калия. Само название калий химический элемент получил из перевода слова щелочь с английского языка. На английском слово щелочь пишется как «alkali», которое в свою очередь переводится с арабского как «растения пепла». Другое название «Kalium» предложил в 1809 году немецкий химик Людвиг Вильгельм Гилберт. По итогу на сегодняшний день существует два латинских названия для этого химического элемента. В то же время «Золотая книга» из Международного союза теоретической и прикладной химии установила обозначение для калия буквой K.

Где и как добывают Калий

Такой химический элемент как калий добывается во многих местах по всему миру. Его ежегодная добыча составляет около 90 млн тонн. Лидерами по добыче калия являются Канада, Россия, США, Израиль, Германия, Казахстан и Белорусь. Этот металл добывают в большей своей части в виде минералов. К таким минералам относят карналлит, лангбейнит, полигалит и сильвин. Самые большие залежи калия предположительно находятся в канадской провинции Саскачеван. После успешной добычи калиевой руды необходимо отчистить калий от других химических элементов. От чистоты калия (в принципе как и любого элемента) будет зависеть его дальнейшая цена. Цена за 1 кг калия хорошей чистоты варьируется в районе 22 долларов США.

На сегодняшний день существует множество способов отделения калиевых солей от натрия и магния с последующим их применением. Наиболее часто используемый метод – это фракционное осаждение с использованием различий растворимости солей. Следующим по популярности идет электростатическое разделение солевой смеси из калиевых растворов. Это позволяет экономить и время и средства.

Чистый калий также добывается в больших количествах. Метод его очистки практически ничем не изменился со времен открытия Хэмфри Дэви. Метод электролиза являлся самым эффективным и распространенным в мире до 1950-х годов. Хотя некоторые предприятия используют этот метод и по сей день. С 1950-х годов популярность захватил метод термической реакции натрия с хлоридом калия. Сегодня этот метод является самым популярным.

Распространенность Калия

Калий является довольно распространенным химическим элементом. Он занимает 20 место по распространенности среди других элементов в нашей солнечной системе и 17 место по массе на Земле.  В тоже время он замыкает топ-7 элементов по массе в земной коре. Его процентное отношение к общей массе земной коры составляет около 2,6% по оценкам ученых.

Известно, что калий образуется в сверхновых звездах в результате нуклеосинтеза из более легких атомов. Так же калий образуется в сверхновых типа II в результате взрывного процесса сжигания кислорода и в результате горения неона. Что же касается Земли, то калий встречается сдесь только в виде катиона в соединениях калия. Это связано с тем, что он имеет только один внешний электрон и очень легко его высвобождает. Калий входит в состав очень многих минеральных пород. Самыми важными из них являются сильвин, сильвинит, карналлит, каинит, счелит, полигалит, ортоклаз и мусковит. Так же не нужно забывать про морскую и океаничечкую воду как источник калия. Содержание в морской воде калия ориентировочно находится в соотношении 27 частей на тысячу. В подтверждение стоит отметить, что у одного из крупнейших добытчиков калия (Израиль) источником добычи является Мертвое море.

Применение Калия

Применение калия является очень неоднозначным. С одной стороны калий добывается относительно в очень больших количествах. С другой стороны 95% добытого калия отправляется для переработки в различного рода удобрения. Оставшиеся 5% добытого калия распределяется в разных частях на такие отрасли как медицина, пищевая промышленность, военная и химическая промышленность. Ионы калия являются важным компонентом питания растений и встречаются почти во всех видах почв.

Стоит заметить, что 90% от общего количества калия выделяемого для производства удобрений поступает в виде соединения KCl. В медицине это же соединение используется для лечения недостатка калия в организме человека. Оно применяется в виде внутривенной инъекции или орально в виде таблеток.

В пищевой промышленности калий применяется исключительно в виде разлличного рода пищевых добавок. В пример можно привести тартрат калий-натрия, который является основополагающим компонентом разрыхлителей. Бромат калия выступает в качестве сильного окислителя. Его пищевая нуменклатура — Е924, и используется в основном для увеличения прочности теста и регуляции величины его подъема. Калий используется также при добыче драгоценных металлов, таких как золото и серебро. Его химические соединения используются и военной промышленности. Так к примеру супероксид калия KO₂ используется в качестве портативного источника кислорода и поглотителя диоксида углерода. В основном это системы вентиляции шахт, бункеры, а также системы жизнеоьеспечения подводных лодок и космических кораблей.

Интересные факты

Интересных фактов связанных с калием достаточно много. Стоит начать с того, что калий играет огромнейшую роль в организме человека. Его недостаток может вызывать огромное количество проблем со здоровьем у человека. Он оказывает влияние на тонус сосудов, артериальное давление, секрецию гормонов, желудочно-кишечную моторику, метаболизм глюкозы и инсулина, а также почечную концентрирующую способность. Калий в чистом виде является очень взрывоопасным веществом. Взрыв может произойти от малейшей царапины и его очень трудно потом потушить.

Еще одним интересным фактом является то, что цианистый калий используется спецслужбами(информация не подтвержденаа). Вероятно, на сегодняшний день, каждый видел фильм с иностранными разведчиками, которые в случае попадания в плен раскалывают зубами ампулу с цианистым калией и мгновенно умирают. В этом есть доля истины. Ведь цианистый калий является сильнейшим неорганическим ядом. И смертельная доза его составляет 1.7 мг/кг веса. Только принцип его действия основывается не на разъедании половины челюсти как в фильмах, а на неспособности усваивать клетками организма кислорода. И неизвестно конечно через какое время наступает смерть и возможно ли спасти человека после приема смертельной ампулы. То есть человек просто задыхается. Еще одним интересным фактом является то, что при определенном экстремальном давлении из твердого металлического калия начинает вытекать жидкая составляющая калия.

История калия

Калий был открыт осенью 1807 года английским химиком Дэви при электролизе твёрдого едкого кали. Увлажнив едкий кали, ученый выделил металл, которому дал название потассий, намекая на производство поташа (необходимого ингредиента для изготовления моющих средств) из золы. Своё привычное название металл получил через два года, в 1809г, инициатором переименования вещества стал Л.В. Гильберт, предложивший название калий (от арабского аль-кали – поташ).

Общая характеристика калия

Калий (лат. Kalium) является мягким щелочным металлом, элементом главной подгруппы I группы, IV периода периодической системы химических элементов Д.И. Менделеева, имеет атомный номер 19 и обозначение – К.

Нахождение в природе

Калий в свободном состоянии в природе не встречается, он входит в состав всех клеток. Достаточно распространённый металл, занимает 7-е место по содержанию в земной коре (calorizator). Основными поставщиками калия являются Канада, Белоруссия и Россия, имеющие крупные месторождения данного вещества.

Физические и химические свойства

Калий – легкоплавкий металл серебристо-белого цвета. Имеет свойство окрашивать открытый огонь в яркий фиолетово-розовый цвет.

Калий имеет высокую химическую активность, это сильный восстановитель. При реакции с водой происходит взрыв, при длительном нахождении на воздухе полностью разрушается. Поэтому калий требует определённых условий для хранения – его заливают слоем керосина, силикона или бензина, для исключения вредных для металла контактов с водой и атмосферой.

Продукты питания богатые калием

Основными пищевыми источниками калия являются сушёные абрикосы, дыня, бобы, киви, картофель, авокадо, бананы, брокколи, печень, молоко, ореховое масло, цитрусовые, виноград, все зелёные овощи с листьями, листья мяты, семечки подсолнуха. Калия достаточно много в рыбе и молочных продуктах. Вообще, калий входит в состав почти всех растений. Яблочный уксус и мёд – чемпионы по содержанию калия.

Суточная потребность в калии

Суточная потребность организма человека в калии зависит от возраста, физического состояния и даже места проживания. Взрослым здоровым людям нужно 2,5г калия, беременным женщинам – 3,5г, спортсменам – до 5-ти грамм калия ежедневно. Количество необходимого калия для подростков рассчитывается по весу – 20 мг калия на 1 кг массы тела.

Полезные свойства калия и его влияние на организм

Калий вместе с натрием регулирует водный баланс в организме и нормализует ритм сердца, поддерживает концентрацию и физиологические функций магния.

Калий участвует в процессе проведения нервных импульсов и передачи их на иннервируемые органы. Способствует лучшей деятельности головного мозга, улучшая снабжение его кислородом. Оказывает положительное влияние при многих аллергических состояниях. Калий необходим для осуществления сокращений скелетных мышц. Калий регулирует содержание в организме солей, щелочей и кислот, чем способствует уменьшению отёков.

Калий содержится во всех внутриклеточных жидкостях, он необходим для нормальной жизнедеятельности мягких тканей (мышц, сосудов и капилляров, желез внутренней секреции и т.д.)

Усвояемость калия

Калий всасывается в организм из кишечника, куда поступает с пищей, выводится с мочой обычно в таком же количестве. Излишний калий выводится из организма тем же путём, не задерживается и не накапливается. Препятствиями для нормального всасывания калия могут послужить чрезмерное употребление кофе, сахара, алкоголя.

Взаимодействие с другими

Калий работает в тесном контакте с натрием и магнием, при росте концентрации калия из организма стремительно выводится натрий, а уменьшение количества магния может нарушить усвоение калия.

Признаки нехватки калия

Нехватка калия в организме характеризуется мышечной слабостью, быстрой утомляемостью, снижением иммунитета, сбоями в работе миокарда, нарушениями показателей артериального давления, учащённым и затруднённым дыханием. Кожные покровы могут шелушиться, повреждения плохо заживают, волосы становятся очень сухими и ломкими. Происходят сбои в работе желудочно-кишечного тракта – тошнота, рвота, расстройства желудка вплоть до гастрита и язвы.

Признаки избытка калия

Переизбыток калия наступает при передозировке препаратов, содержащих калий и характеризуется нервно-мышечными расстройствами, повышенной потливостью, возбудимостью, раздражительностью и плаксивостью. Человек постоянно испытывает чувство жажды, которое приводит к частым мочеиспусканиям. Желудочно-кишечный тракт реагирует кишечными коликами, чередованием запоров и поносов.

Применение калия в жизни

Калий в виде основных соединений находит широкое применение в медицине, сельском хозяйстве и промышленности. Калийные удобрения необходимы для нормального роста и вызревания растений, а всем известная марганцовка, это не что иное, как перманганат калия, испытанный временем антисептик.

Автор: Виктория Н. (специально для Calorizator.ru)
Копирование данной статьи целиком или частично запрещено.

(от араб. аль-кали — поташ; лат. Kalium) К, хим. элемент I гр. периодич. системы; относится к щелочным металлам, ат. н. 19, ат. м. 39,0983. Состоит из двух стабильных изотопов ³⁹ К (93,259%) и ⁴¹ К (6,729%), а также радиоактивного изотопа ⁴⁰ К (Т _1/2 1,32.10⁹ лет). Поперечное сечение захвата тепловых нейтронов для прир. смеси изотопов 1,97.10^—28 м ². Конфигурация внеш. электронной оболочки 4s¹; степень окисления +1; энергия ионизации К ⁰ : К ⁺ : К ²⁺ соотв. 4,34070 эВ и 31,8196 эВ; сродство к электрону 0,47 эВ; электроотрицательность по Полингу 0,8; атомный радиус 0,2313 нм, ионный радиус (в скобках указано координац. число) К ⁺ 0,151 нм (4), 0,152 нм (6), 0,160 нм (7), 0,165 нм (8), 0,178 нм (12). Содержание в земной коре 2,41% по массе. Осн. минералы: сильвин КСl, карналлит KCl.MgCl₂.6H₂O, калиевый полевой шпат (ортоклаз) K[AlSi₃O₈], мусковит калиевая слюда) KAl₂[AlSi₃O₁₀](OH, F)₂, каинит КСl.MgSO₄.3H₂O, полигалит K₂SO₄.MgSO₄.2CaSO₄.2Н ₂ О, алунит KAl₃(SO₄)₂(OH)₆.
Свойства. К. — мягкий серебристо-белый металл с кубич. решеткой, а =0,5344 нм, z =2, пространств. группа Im3m. Т. пл. 63,51 °С, т. кип. 761 °С; плотн. 0,8629 г/см ³; С ⁰_p 29,60 Дж/(моль. К); DH⁰ _пл 2,33 кДж/моль, DH⁰ _возг 89,0 кДж/моль; S⁰₂₉₈64,68 Дж/(моль. К); ур-ние температурной зависимости давления пара (в мм рт. ст.): lgp =7,34 — 4507/T(373-474 К); r 6,23.10^-8 Ом. м (0°С), 8,71.10^—8 Ом. м (25°C) и 13,38.10^-8 Ом. м (77°С); g 0,114 Н/м (334 К), h 5,096.10^—4 Н. с/м ² (350 К); теплопроводность 99,3 Вт/(м. К) при 273 К и 44,9 Вт/(м. К) при 473 К; при 273-323 К температурные коэф. линейного и объемного расширения составляют соотв. 8,33.10^—5 К ^—1 и 2,498.10^—4 К ^—1. К. может обрабатываться прессованием и прокаткой. К. химически очень активен. Легко взаимод. с О ₂ воздуха, образуя калия оксид К ₂ О, пероксид К ₂ О ₂ и надпероксид КО ₂; при нагр. на воздухе загорается. С водой и разб. к-тами взаимод. со взрывом и воспламенением, причем H₂SO₄ восстанавливается до S^2—, S⁰ и SO₂, а НNО ₃ — до NO, N₂O и N₂. При нагр. до 200-350 °С реагирует с Н ₂ с образованием гидрида КН. Воспламеняется в атмосфере F₂, слабо взаимод. с жидким Сl₂, но взрывается при соприкосновении с Вr₂ и растирании с I₂; при контакте с межгалогенными соед. воспламеняется или взрывается. С S, Se и Те при слабом нагревании образует соотв. K₂S, K₂Se и К ₂ Те, при нагр. с Р в атмосфере азота — К ₃ Р и К ₂ Р ₅, с графитом при 250-500 °С — слоистые соед. состава С ₈ К-С ₆₀ К. С СО ₂ не реагирует заметно при 10-30°С; стекло и платину разрушает выше 350-400 °С. К. раств. в жидком NH₃ (35,9 г в 100 мл при — 70 °С), анилине, этилендиамине, ТГФ и диглиме с образованием р-ров с металлич. проводимостью. Р-р в NH₃ имеет темно-синий цвет, в присут. Pt и следов воды разлагается, давая KNH₂ и Н ₂. С азотом К. не взаимод. даже под давлением при высоких т-рах. При взаимод. с NH₄N₃ в жидком NH₃ образуется азид KN₃. К. не раств. в жидких Li, Mg, Cd, Zn, Al и Ga и не реагирует с ними. С натрием образует интерметаллид KNa₂ (плавится инконгруэнтно при 7°С), с рубидием и цезием -твердые р-ры, для к-рых миним. т-ры плавления составляют соотв. 32,8 °С (81,4% по массе Rb) и — 37,5 °С (77,3% Cs). С ртутью дает амальгаму, содержащую два меркурида -KHg₂ и KHg (т. пл. соотв. 270 °С и 180°С). С таллием образует КТl (т. пл. 335 °С), с оловом — K₂Sn, KSn, KSn₂ и KSn₄, со свинцом — КРb (т. пл. 568 °С) и фазы состава К ₂ Рb₃, КРb₂ и КРb₄, с сурьмой — К ₃Sb и KSb (т. пл. соотв. 812 и 605 °С), с висмутом — K₃Bi, K₃Bi₂ и KBi₂ (т. пл. соотв. 671, 420 и 553 °C). К. энергично взаимод. с оксидами азота, а при высоких т-рах — с СО и СО ₂. Восстанавливает В ₂ О ₃ и SiO₂ соотв. до В и Si, оксиды Al, Hg, Ag, Ni и др. — до своб. металлов, сульфаты, сульфиты, нитраты, нитриты, карбонаты и фосфаты металлов — до оксидов или своб. металлов. Со спиртами К. образует алкоголяты, с галогеналкилами и галогенарилами — соотв. калийалкилы и калийарилы. Важнейшим соед. К. посвящены отд. статьи (см., напр., Калия гидроксид, Калия иодид, Калия карбонат, Калия хлорид). Ниже приводятся сведения о др. важных соединениях. Пероксид К ₂ О ₂ — бесцв. кристаллы с ромбич. решеткой ( а =0,6736 нм, b= 0,7001 нм, с =0,6479 нм, z = 4, пространств. группа Рппп); т. пл. 545 °С; плотн. 2,40 г/см ³; C⁰_p 90,8 Дж/(моль. К); DH⁰ _пл 20,5 кДж/моль, DH⁰ _обр -443,0 кДж/моль; S⁰₂₉₈ 117 Дж/(моль. К). На воздухе К ₂ О ₂ мгновенно окисляется до КО ₂; энергично взаимод. с водой с образованием КОН и О ₂, с СО ₂ дает К ₂ СО ₃ и О ₂. Получают пероксид пропусканием дозированного кол-ва О ₂ через р-р К. в жидком NH₃, разложением КО ₂ в вакууме при 340-350 °С. Надпероксид КО ₂ — желтые кристаллы с тетрагон. решеткой ( а Ч0,5704 нм, с Ч0,6699 нм, z = 4, пространств. группа I4/mmm; плотн. 2,158 г/см ³); выше 149°С переходит в кубич. модификацию (а= 0,609 нм); т. пл. 535 °С; С ⁰_p 77,5 Дж/(моль. К);

DH⁰ _обр — 283,2 кДж/моль, DH⁰ _пл 20,6 кДж/моль; S⁰₂₉₈ 125,4 Дж/(моль. К); парамагнетик. Сильный окислитель. Взаимод. с водой с образованием КОН и О ₂. Диссоциирует, напр., в бензоле, давая анион-радикал О ₂^—. Сера при нагр. с КО ₂ воспламеняется и дает K₂SO₄. С влажными СО ₂ и СО надпероксид образует К ₂ СО ₃ и О ₂, с NO₂ — KNO₃ и О ₂, с SO₂ — K₂SO₄ и О ₂. При действии конц. H₂SO₄ выделяется О ₃, при р-ции с NH₃ образуются N₂, H₂O и КОН. Смесь КО ₂ с графитом взрывает. Получают надпероксид сжиганием К. в воздухе, обогащенном влажным О ₂ и нагретом до 75-80 °С. Озонид КО ₃ — красные кристаллы с тетрагон. решеткой ( а =0,8597 нм, с= 0,7080 нм, пространств. группа I4/mcm); плотн. 1,99 г/см ³; устойчив только при хранении в герметически закрытых сосудах ниже 0°С, при более высокой т-ре распадается на КО ₂ и О ₂; С ⁰_p >75 Дж/(моль. К); S⁰₂₉₈ 105 Дж/(моль. К); парамагнетик. Р-римость в жидком NH₃ (г в 100 г); 14,82 (-35°С), 12,00 (-63,5°С), эвтектика NH₃ — KO₃ (5 г в 100 г NH₃) имеет т. пл. — 80 °С; при длит. хранении аммиачные р-ры разлагаются на NH₄O₃ и KNH₂. Раств. в фреонах. КО ₃ — сильный окислитель; мгновенно реагирует с водой уже при 0 °С, давая КОН и О ₂; с влажным СО ₂ образует К ₂ СО ₃ и КНСО ₃. Получают: взаимод. смеси О ₃ и О ₂ с КОН или КО ₂ при т-ре ниже 0°С с послед. экстракцией жидким NH₃; озонированием суспензии КО ₂ или КОН во фреоне 12. Надпероксид, пероксид и озонид К. — компоненты составов для регенерации воздуха в замкнутых системах (шахты, подводные лодки, космич. корабли). Гидрид КН — бесцв. кристаллы с кубич. решеткой ( а =0,570 нм, z = 4, пространств. группа Fm3m); т. пл. 619 °С при давлении водорода 6,86 МПа; плотн. 1,52 г/см ³; С ⁰_p38,1 Дж/(моль. К); DH⁰ _обр -57,8 кДж/моль, DH⁰ _пл 21,3 кДж/моль; S⁰₂₉₈ 50,2 Дж/(моль. К). Разлагается при нагр. на К. и Н ₂. Сильный восстановитель. Воспламеняется во влажном воздухе, в среде F₂ и Сl₂. Энергично взаимод. с водой, давая КОН и Н ₂. При нагр. с N₂ или NH₃ образует KNH₂, с H₂S — K₂S и Н ₂, с расплавл. серой — K₂S и H₂S, с влажным СO₂ — НСООК, с СО — НСООК и С. Получают КН взаимод. К. с избытком Н ₂ при 300-400 °С. Гидрид восстановитель в неорг. и орг. синтезах. Азид КN₃ — бесцв. кристаллы с тетрагон. решеткой (а= 0,6094 нм, с= 0,7056 нм, z = 4, пространств. группа I4/mcm); С ⁰ _р >76,9 Дж/(моль. К); DH⁰ _обр -1,7 кДж/моль; S⁰₂₉₈ 104,0 Дж/(моль. К); т. пл. 354 °С, выше разлагается на К. и азот; плотн. 2,056 г/см ³. Р-римость (г в 100 г р-рителя): вода -41,4 (0°С), 105,7 (100°С), этанол — 0,137 (16°С). В эфире и бензоле раств. плохо. Водой гидролизуется; эвтектика Н ₂ О -KN₃ (35,5 г в 100 г) имеет т. пл. — 12°С. Получают: взаимод. N₂O с расплавл. KNH₂ при 280 °С; действием КОН на р-р HN₃.
Получение. К. производят взаимод. Na с КОН при 380-450 °С или КСl при 760-890 °С. Р-ции проводят в атмосфере N₂. Взаимод. КОН с жидким Na осуществляют противотоком в тарельчатой колонне из Ni. При р-ции с КСl пары Na пропускают через расплав КСl. Продукт р-ций -сплав К — Na. Его подвергают ректификации и получают К. с содержанием примесей (в % по массе): 1.10^—3 Na и 1.10^—4 С1. Кроме того, К. получают: вакуум-термич. восстановлением КСl карбидом СаС ₂, сплавами Si-Fe или Si-Al при 850-950 °С и остаточном давлении 1,33 13,3 Па; нагреванием К ₂ Сr₂ О ₇ с Zr при 400 °С; электролизом КОН с железным катодом; электролизом КСl (или его смеси К ₂ СО ₃) с жидким свинцовым катодом при 680-720 °С с послед. разделением К и Рb вакуумной дистилляцией; электролизом 50%-ного р-ра KNH₂ в жидком NH₃ при 25 °С и давлении 0,75 МПа с амальгамой К. в качестве анода и катодом из нержавеющей стали, при этом образуется 30%-ный р-р К. в жидком NH₃, к-рый выводится из аппарата для отделения К.
Определение. Качественно К. обнаруживают по розово-фиолетовому окрашиванию пламени и по характерным лииниям спектра: 404,41, 404,72, 766,49, 769,90 нм. Наиб. распространенные количеств. методы, особенно в присут. др. щелочных металлов, — эмиссионная пламенная фотометрия (чувствительность 1.10^—4 мкг/мл) и атомно-абсорбц. спектрометрия (чувствительность 0,01 мкг/мл). В меньшей степени используются химико-спектральный и спектрофотометрич. методы с применением дипикриламина (чувствительность 0,2-1,0 мкг/мл). При большом содержании К. в пробе применяют гравиметрич. метод с осаждением К. в виде тетрафенилбората, K₂[PtCl₆] или КСlO₄ в среде бутанола.
Применение. К. — материал электродов в хим. источниках тока; компонент катодов-эмиттеров фотоэлементов и термоэмиссионных преобразователей, а также фотоэлектронных умножителей; геттер в вакуумных радиолампах; активатор катодов газоразрядных устройств. Сплав К. с Na -теплоноситель в ядерных реакторах. Радиоактивный изотоп ⁴⁰ К служит для определения возраста горных пород (калийаргоновый метод). Искусств. изотоп ⁴² К (T_1/2 12,52 года) -радиоактивный индикатор в медицине и биологии. К. — важнейший элемент (в виде соед.) для питания растений (см. Калийные удобрения). К. вызывает сильные ожоги кожи, при попадании мельчайших его крошек в глаза тяжело их поражает, возможна потеря зрения. Загоревшийся К. заливают минер. маслом или засыпают смесью талька и NaCl. Хранят К. в герметически закрытых железных коробках под слоем обезвоженного керосина или минер. масла. Отходы К. утилизируют обработкой их сухим этанолом или пропанолом с послед. разложением образовавшихся алкоголятов водой. К. открыл Г. Дэви в 1807. Лит.: Натрий и калий, Л., 1959; Коренман И. М., Аналитическая химия калия, М., 1964; Теплофизическиe свойства щелочных металлов, М., 1970, (Тр. МЭИ, в. 75); Вольнов И. И., Перекисные соединения щелочных металлов, М., 1980. Б. Д. Степин.

Глоссарий. Химия