Potassium

Potassium is a chemical element; it has symbol K and atomic number19. It is a silvery white metal that is soft enough to easily cut with a knife. Potassium metal reacts rapidly with atmospheric oxygen to form flaky white potassium peroxide in only seconds of exposure. It was first isolated from potash, the ashes of plants, from which its name derives. In the periodic table, potassium is one of the alkali 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. 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, and occurs in many minerals such as orthoclase, a common constituent of granites and other igneous rocks.
Potassium is chemically very similar to sodium, the previous element in group 1 of the periodic table. They have a similar first ionization energy, which allows for each atom to give up its sole outer electron. It was first suggested in 1702 that they were distinct elements that combine with the same anions to make similar salts, which was demonstrated in 1807 when elemental potassium was first isolated via electrolysis. Naturally occurring potassium is composed of three isotopes, of which potassium-40| is radioactive. Traces of are found in natural sources of 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 by increasing potassium excretion by the kidneys and sequestering potassium in the liver and muscles to avoid changes in serum potassium levels.
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.

Etymology

The English name for the element potassium comes from the word potash, which refers to an early method of extracting various potassium salts: placing in a pot the ash of burnt wood or tree leaves, adding water, heating, and evaporating the solution. Humphry Davy named the element potassium after isolating the metal itself.
The symbol K stems from kali, itself from the root word alkali, which in turn comes from 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. In 1807, Humphry Davy produced the element via electrolysis: in 1809, Ludwig Wilhelm Gilbert proposed the name Kalium for Davy's "potassium". In 1814, the Swedish chemist Berzelius advocated the name kalium for potassium, with the chemical symbol K.
The English and French-speaking countries adopted the name Potassium, which was favored by Davy and French chemists Joseph Louis Gay-Lussac and Louis Jacques Thénard, whereas the other Germanic countries adopted Gilbert and Klaproth's name Kalium. The "Gold Book" of the International Union of Pure and Applied Chemistry has designated the official chemical symbol as K.

Discovery

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

Properties

Potassium is a soft silvery solid that easily cut with a knife. 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. It readily dissolves in liquid ammonia, up to 480 g per 1000 g of ammonia at 0°C to form the electride, which features an electron as an anion.

Compounds

Reflecting its low first ionization energy of 418.8kJ/mol, potassium is a strong reducing agent, i.e., it readily releases an electron upon contact with other materials. With graphite, potassium metal forms graphite intercalation compounds. One such compound has the formula KC₈, a gold colored solid that is described as a K⁺ salt of negatively charged graphite. Potassium can reduce many salts to the metal as illustrated by the Rieke method for making magnesium powder from magnesium chloride:
Most potassium compounds are ionic. Owing to the high hydration energy of the ion, these salts often exhibit excellent water solubility. The main species in water solution are the aquo complexes where n = 6 and 7. Although typically insoluble in organic solvents, potassium salts dissolve in polar organic solvents in the presence of crown ethers and cryptand. These organic ligands envelop K⁺ ions, giving lipophilic coordination complexes. Similar complexation phenomena are found for some ion-binding antibiotics.

Binary compounds

Potassium forms many binary compounds, i.e., compounds of potassium and one other element. The inventory is so extensive that one gap merits mention: no nitride of potassium is known. Potassium hydride forms directly from the elements:
It is a white, pyrophoric solid that finds some use as a base.
All of the halides salts are well known: potassium fluoride, potassium chloride, and potassium bromide, potassium iodide. Four oxides of potassium are well studied: potassium oxide, potassium peroxide, potassium superoxide and potassium ozonide. These species all hydrolyze to give potassium hydroxide. Similarly an extensive array of sulfides, selenides, and tellurides are well characterized. Although such simple salts are typically white and diamagnetic, is something of an exception, being deep yellow and paramagnetic.

Ternary and more complex compounds

Although rarely encountered in anhydrous form, KOH is one of the dominant compounds of potassium from the commercial perspective. It is a strong base and highly corrosive. Illustrative of its hydrophilic character, as much as 1.21 kg of KOH can dissolve in a liter of water. KOH reacts readily with carbon dioxide to produce potassium carbonate, 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. Potassium-based soaps are used in soap dispensers because they more soluble in water than sodium soaps.
Nitrate, nitrite, sulfate, and various phosphates also form potassium salts, all white solids, that are widely used. Illustrating the thermal stability typical for these materials, potassium nitrate, sodium nitrate, and sodium nitrite form a eutectic, which remains liquid from 142 to 600 °C.
Sodium and potassium salts display virtually identical properties in aqueous solution, but their differing solubilities are of practical value The distinctive solubility of potassium heptafluorotantalate allows the purification of tantalum from the otherwise persistent contaminant of niobium. The solubility of the K⁺ compound differs strikingly from that for the Na⁺ compound in the pairs sodium tetraphenylborate/potassium tetraphenylborate, sodium cobaltinitrite/potassium cobaltinitrite, and sodium hexachloroplatinate/ potassium hexachloroplatinate. These differences are the bases for gravimetric analysis for K⁺.
Several potassium-containing reagents, so-called primary standards, have the advantage of being non-hygroscopic, in contrast the corresponding sodium salts. Thus, the oxidant potassium dichromate, the acid potassium hydrogen phthalate, and the reductant potassium ferrocyanide can be handled in air without gaining weight by hydration. Potassium salts are often produced from the sodium salts e.g., sodium chromate and sodium permanganate, which are more directly obtained from ores.

Organopotassium compounds

s are mainly of academic interest. They feature highly polar covalent K–C bonds. An example is potassium diphenylmethyl.

Isotopes

There are 25 known isotopes of potassium, three of which occur naturally: , , and . Naturally occurring potassium-40| has a half-life of years. It decays to stable Argon| by electron capture or positron emission or to stable Calcium| by beta decay. This decay results in a relatively higher concentration of Argon in the atmosphere. The decay of to is the basis of a common method for dating rocks. The conventional potassium–argon dating method depends on the assumption that the rocks contained no argon at the time of formation and that all the subsequent radiogenic argon was quantitatively retained. Minerals are dated by measurement of the concentration of potassium and the amount of radiogenic 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. 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 on Earth.
occurs in natural potassium in sufficient quantity that large bags of those substitutes can be used as a radioactive source for classroom demonstrations. is the radioisotope with the largest abundance in the human body. In healthy animals and people, represents the largest source of radioactivity, greater even than Carbon-14|. In a human body of 70 kg, about 4,400 nuclei of decay per second. The activity of natural potassium is 31 Bq/g.