Math Is Fun Forum

Potassium

Gist

Potassium, with the chemical symbol K and atomic number 19, is a soft, silvery-white metal that is a member of the alkali metal group. It is highly reactive and not found naturally in its pure form. Potassium is crucial for plant growth, being a key component in fertilizers, and also plays an essential role in human and animal nutrition.

Potassium is a vital mineral and electrolyte that plays a crucial role in various bodily functions. It helps regulate fluid balance, muscle contractions, nerve signals, and the heartbeat. Most people can obtain sufficient potassium through their diet, with fruits and vegetables being excellent sources.

Summary

Potassium is a chemical element; it has symbol K (from Neo-Latin kalium) and atomic number 19. It is a silvery white metal that is soft enough to easily cut with a knife. 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), 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 40K is radioactive. Traces of 40K 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.

Details

Potassium (K), chemical element of Group 1 (Ia) of the periodic table, the alkali metal group, indispensable for both plant and animal life. Potassium was the first metal to be isolated by electrolysis, by the English chemist Sir Humphry Davy, when he obtained the element (1807) by decomposing molten potassium hydroxide (KOH) with a voltaic battery.

Element Properties

atomic number  :  19
atomic weight  :  39.098
melting point  :  63.28 °C (145.90 °F)
boiling point  :  760 °C (1,400 °F)
specific gravity  :  0.862 (at 20 °C, or 68 °F)
oxidation states  :  +1, −1 (rare)

Properties, occurrence, and uses

Potassium metal is soft and white with a silvery lustre, has a low melting point, and is a good conductor of heat and electricity. Potassium imparts a lavender colour to a flame, and its vapour is green. It is the seventh most abundant element in Earth’s crust, constituting 2.6 percent of its mass.

The potassium content of the Dead Sea is estimated at approximately 1.7 percent potassium chloride, and many other salty bodies of water are rich in potassium. The waste liquors from certain saltworks may contain up to 40 grams per litre of potassium chloride and are used as a source of potassium.

Most potassium is present in igneous rocks, shale, and sediment in minerals such as muscovite and orthoclase feldspar that are insoluble in water; this makes potassium difficult to obtain. As a result, most commercial potassium compounds (often loosely called potash) are obtained via electrolysis from soluble potassium compounds, such as carnallite (KMgCl3∙6H2O), sylvite (potassium chloride, KCl), polyhalite (K2Ca2Mg[SO4]4∙2H2O), and langbeinite (K2Mg2[SO4]3), which are found in ancient lake beds and seabeds.

Potassium is produced by sodium reduction of molten potassium chloride, KCl, at 870 °C (1,600 °F). Molten KCl is continuously fed into a packed distillation column while sodium vapour is passed up through the column. By condensation of the more volatile potassium at the top of the distillation tower, the reaction Na + KCl → K + NaCl is forced to the right. Efforts to devise a scheme for commercial electrolytic production of potassium have been unsuccessful because there are few salt additives that can reduce the melting point of potassium chloride to temperatures where electrolysis is efficient.

There is little commercial demand for potassium metal itself, and most of it is converted by direct combustion in dry air to potassium superoxide, KO2, which is used in respiratory equipment because it liberates oxygen and removes carbon dioxide and water vapour. (The superoxide of potassium is a yellow solid consisting of K+and O2− ions. It also can be formed by oxidation of potassium amalgam with dry air or oxygen.) The metal is also used as an alloy with sodium as a liquid metallic heat-transfer medium. Potassium reacts very vigorously with water, liberating hydrogen (which ignites) and forming a solution of potassium hydroxide, KOH.

Sodium-potassium alloy (NaK) is used to a limited extent as a heat-transfer coolant in some fast-breeder nuclear reactors and experimentally in gas-turbine power plants. The alloy is also used as a catalyst or reducing agent in organic synthesis.

In addition to the alloys of potassium with lithium and sodium, alloys with other alkali metals are known. Complete miscibility exists in the potassium-rubidium and potassium-cesium binary systems. The latter system forms an alloy melting at approximately −38 °C (−36 °F). Modification of the system by the addition of sodium results in a ternary eutectic melting at approximately −78 °C (−108 °F). The composition of this alloy is 3 percent sodium, 24 percent potassium, and 73 percent cesium. Potassium is essentially immiscible with all the alkaline-earth metals, as well as with zinc, aluminum, and cadmium.

Potassium (as K+) is required by all plants and animals. Plants need it for photosynthesis, regulation of osmosis and growth, and enzyme activation. Every animal has a closely maintained potassium level and a relatively fixed potassium-sodium ratio. Potassium is the primary inorganic cation within the living cell, and sodium is the most abundant cation in extracellular fluids. In higher animals, selective complexants for Na+ and K+ act at cell membranes to provide “active transport.” This active transport transmits electrochemical impulses in nerve and muscle fibres and in balancing the activity of nutrient intake and waste removal from cells. Too little or too much potassium in the body is fatal; however, potassium in the soil ensures the presence of this indispensable element in food.

The potassium content of plants varies considerably, though it is ordinarily in the range of 0.5–2 percent of the dry weight. In humans the ratio of potassium between the cell and plasma is approximately 27:1. The potassium content of muscle tissue is approximately 0.3 percent, whereas that of blood serum is about 0.01–0.02 percent. The dietary requirement for normal growth is approximately 3.3 grams (0.12 ounce) of potassium per day, but the ingestion of more than 20 grams (0.7 ounce) of potassium results in distinct physiological effects. Excess potassium is excreted in the urine, and a significant quantity may be lost during sweating.

Natural potassium consists of three isotopes: potassium-39 (93.26 percent), potassium-41 (6.73 percent), and radioactive potassium-40 (about 0.01 percent); several artificial isotopes have also been prepared. Potassium-39 is normally about 13.5 times more plentiful than potassium-41. The natural radioactivity of potassium is due to beta radiation from the potassium-40 isotope ({10}^{9} years half-life). The disintegration of potassium-40 is used in geological age calculations (see potassium-argon dating). Potassium easily loses the single 4s electron, so it normally has an oxidation state of +1 in its compounds, although compounds that contain the anion, K−, can also be made.

Additional Information

Potassium is a vital mineral and electrolyte that plays a crucial role in various bodily functions. It helps regulate fluid balance, muscle contractions, nerve signals, and the heartbeat. Most people can obtain sufficient potassium through their diet, with fruits and vegetables being excellent sources.

Lithium

Gist

Lithium is a chemical element; it has symbol Li and atomic number 3. It is a soft, silvery-white alkali metal. Under standard conditions, it is the least dense metal and the least dense solid element.

Lithium is a type of medicine known as a mood stabiliser. It's used to treat mood disorders such as: mania (feeling highly excited, overactive or distracted) hypo-mania (similar to mania, but less severe)

Summary

Lithium is a chemical element; it has symbol Li and atomic number 3. It is a soft, silvery-white alkali metal. Under standard conditions, it is the least dense metal and the least dense solid element. Like all alkali metals, lithium is highly reactive and flammable, and must be stored in vacuum, inert atmosphere, or inert liquid such as purified kerosene or mineral oil. It exhibits a metallic luster. It corrodes quickly in air to a dull silvery gray, then black tarnish. It does not occur freely in nature, but occurs mainly as pegmatitic minerals, which were once the main source of lithium. Due to its solubility as an ion, it is present in ocean water and is commonly obtained from brines. Lithium metal is isolated electrolytically from a mixture of lithium chloride and potassium chloride.

The nucleus of the lithium atom verges on instability, since the two stable lithium isotopes found in nature have among the lowest binding energies per nucleon of all stable nuclides. Because of its relative nuclear instability, lithium is less common in the Solar System than 25 of the first 32 chemical elements even though its nuclei are very light: it is an exception to the trend that heavier nuclei are less common. For related reasons, lithium has important uses in nuclear physics. The transmutation of lithium atoms to helium in 1932 was the first fully human-made nuclear reaction, and lithium deuteride serves as a fusion fuel in staged thermonuclear weapons.

Lithium and its compounds have several industrial applications, including heat-resistant glass and ceramics, lithium grease lubricants, flux additives for iron, steel and aluminium production, lithium metal batteries, and lithium-ion batteries. These uses consume more than three-quarters of lithium production.

Lithium is present in biological systems in trace amounts. It has no established metabolic function in humans. Lithium-based drugs are useful as a mood stabilizer and antidepressant in the treatment of mental illness such as bipolar disorder.

Details

Lithium (Li) is a chemical element of Group 1 (Ia) in the periodic table, the alkali metal group, lightest of the solid elements. The metal itself—which is soft, white, and lustrous—and several of its alloys and compounds are produced on an industrial scale.

Element Properties

atomic number  :  3
atomic weight  :  6.941
melting point  :  180.5 °C (356.9 °F)
boiling point  :  1,342 °C (2,448 °F)
specific gravity  :  0.534 at 20 °C (68 °F)
oxidation state  :  +1

Occurrence and Production

Discovered in 1817 by Swedish chemist Johan August Arfwedson in the mineral petalite, lithium is also found in brine deposits and as salts in mineral springs; its concentration in seawater is 0.1 part per million (ppm). Lithium is also found in pegmatite ores, such as spodumene (LiAlSi2O6) and lepidolite (of varying structure), or in amblygonite (LiAlFPO4) ores, with Li2O contents ranging between 4 and 8.5 percent. It constitutes about 0.002 percent of Earth’s crust.

Until the 1990s the lithium chemical and metal market was dominated by American production from mineral deposits, but by the turn of the 21st century most production was derived from non-U.S. sources; Australia, Chile, and Portugal were the world’s largest suppliers. (Bolivia has half the world’s lithium deposits but is not a major producer of lithium.) The major commercial form is lithium carbonate, Li2CO3, produced from ores or brines by a number of different processes. Addition of hydrochloric acid (HCl) produces lithium chloride, which is the compound used to produce lithium metal by electrolysis. Lithium metal is produced by electrolysis of a fused mixture of lithium and potassium chlorides. The lower melting point of the mixture (400–420 °C, or 750–790 °F) compared with that of pure lithium chloride (610 °C, or 1,130 °F) permits lower-temperature operation of the electrolysis. Since the voltage at which decomposition of lithium chloride takes place is lower than that of potassium chloride, lithium is deposited at a purity level greater than 97 percent. Graphite anodes are used in the electrolytic production of lithium, while the cathodes are made of steel. The pure lithium formed at the cathode coalesces at the surface of the electrolyte to form a molten pool, which is protected from reaction with air by a thin film of the electrolyte. The lithium is ladled from the cell and cast by pouring it into a mold at a temperature only slightly above the melting point, leaving the solidified electrolyte behind. The solidified lithium is then remelted, and materials insoluble in the melt either float to the surface or sink to the bottom of the melt pot. The remelting step reduces the potassium content to less than 100 parts per million. Lithium metal, which can be drawn into wire and rolled into sheets, is softer than lead but harder than the other alkali metals and has the body-centred cubic crystal structure.

Many lithium alloys are produced directly by the electrolysis of molten salts, containing lithium chloride in the presence of a second chloride, or by the use of cathode materials that interact with the deposited lithium, introducing other elements into the melt.

Significant uses

The principal industrial applications for lithium metal are in metallurgy, where the active element is used as a scavenger (remover of impurities) in the refining of such metals as iron, nickel, copper, and zinc and their alloys. A large variety of nonmetallic elements are scavenged by lithium, including oxygen, hydrogen, nitrogen, carbon, sulfur, and the halogens. Lithium is utilized to a considerable extent in organic synthesis, both in laboratory reactions and industrially. A key reagent that is produced commercially on a large scale is n-butyllithium, C4H9Li. Its principal commercial use is as an initiator of polymerization, for example, in the production of synthetic rubber. It is also extensively used in the production of other organic chemicals, especially pharmaceuticals. Because of its light weight and large negative electrochemical potential, lithium metal, either pure or in the presence of other elements, serves as the anode (negative electrode) in many nonrechargeable lithium primary batteries. Since the early 1990s much work has been done on high-power rechargeable lithium storage batteries for electric vehicles and for power storage. The most successful of these provides for separation of the anode and a cathode such as LiCoO2 by a solvent-free conducting polymer that permits migration of the lithium cation, Li+. Smaller rechargeable lithium batteries are extensively used for cell phones, cameras, and other electronic devices.

Lightweight lithium-magnesium alloys and tough lithium-aluminum alloys, harder than aluminum alone, have structural applications in the aerospace and other industries. Metallic lithium is used in the preparation of compounds such as lithium hydride.

Additional Information

Lithium was discovered from a mineral, while other common alkali metals were discovered from plant material. This is thought to explain the origin of the element’s name; from ‘lithos’ (Greek for ‘stone’).

Appearance

A soft, silvery metal. It has the lowest density of all metals. It reacts vigorously with water.

Uses

The most important use of lithium is in rechargeable batteries for mobile phones, laptops, digital cameras and electric vehicles. Lithium is also used in some non-rechargeable batteries for things like heart pacemakers, toys and clocks.

Lithium metal is made into alloys with aluminium and magnesium, improving their strength and making them lighter. A magnesium-lithium alloy is used for armour plating. Aluminium-lithium alloys are used in aircraft, bicycle frames and high-speed trains.

Lithium oxide is used in special glasses and glass ceramics. Lithium chloride is one of the most hygroscopic materials known, and is used in air conditioning and industrial drying systems (as is lithium bromide). Lithium stearate is used as an all-purpose and high-temperature lubricant. Lithium carbonate is used in drugs to treat manic depression, although its action on the brain is still not fully understood. Lithium hydride is used as a means of storing hydrogen for use as a fuel.

Biological role

Lithium has no known biological role. It is toxic, except in very small doses.

Natural abundance

Lithium does not occur as the metal in nature, but is found combined in small amounts in nearly all igneous rocks and in the waters of many mineral springs. Spodumene, petalite, lepidolite, and amblygonite are the more important minerals containing lithium.

Most lithium is currently produced in Chile, from brines that yield lithium carbonate when treated with sodium carbonate. The metal is produced by the electrolysis of molten lithium chloride and potassium chloride