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lithium

 
Dictionary: lith·i·um   (lĭth'ē-əm) pronunciation
 
n. (Symbol Li)
  1. A soft, silvery, highly reactive metallic element that is used as a heat transfer medium, in thermonuclear weapons, and in various alloys, ceramics, and optical forms of glass. Atomic number 3; atomic weight 6.941; melting point 179°C; boiling point 1,317°C; specific gravity 0.534; valence 1.
  2. Any of several salts of lithium, especially lithium carbonate.

[From LITHIA.]


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A chemical element, Li, atomic number 3, and atomic weight 6.939. Lithium heads the alkali metal family in the periodic table. In nature it is a mixture of the isotopes 6Li and 7Li. Lithium, the lightest solid element, is a soft, low-melting, reactive metal. In many physical and chemical properties it resembles the alkaline-earth metals as much as, or more than, it does the alkali metals. See also Alkaline-earth metals; Periodic table.

The major industrial use of lithium is in the form of lithium stearate as a thickener for lubricating greases. Other important uses of lithium compounds are in ceramics, specifically in porcelain enamel formulation; as an additive to give longer life and higher output in alkaline storage batteries; and in welding and brazing fluxes.

Lithium is a moderately abundant element and is present in the Earth's crust to the extent of 65 parts per million (ppm). This places lithium a little below nickel, copper, and tungsten, and a little above cerium and tin in abundance.

Noteworthy among lithium's physical properties are the high specific heat (heat capacity), large temperature range of the liquid phase, high thermal conductivity, low viscosity, and very low density. Lithium metal is soluble in liquid ammonia and is slightly soluble in the lower aliphatic amines, such as ethyl-amine. It is insoluble in hydrocarbons.

Lithium undergoes a large number of reactions with both organic, and inorganic, reagents. It reacts with oxygen to form the monoxide, Li2O, and the peroxide, Li2O2. Lithium is the only alkali metal that reacts with nitrogen at room temperature to form a nitride, Li3N, which is black. Lithium reacts readily with hydrogen at about 930°F (500°C) to form lithium hydride, LiH. The reaction of lithium metal with water is exceedingly vigorous. Lithium reacts directly with carbon to form the carbide, Li2C2. Lithium combines readily with the halogens, forming halides with the emission of light. While lithium does not react with paraffin hydrocarbons, it does undergo addition reactions with arylated alkenes and with dienes. Lithium also reacts with acetylenic compounds, forming lithium acetylides, which are important in the synthesis of vitamin A.

The most important lithium compound is lithium hydroxide. It is a white powder, and the material of commerce is actually lithium hydroxide monohydrate, LiOH · H2O. Lithium carbonate, LiCO3, finds application in the ceramic industries and in medicine as an antidepressant. Both lithium halides, lithium chloride and lithium bromide, form concentrated brines with ability to absorb moisture over a wide temperature range; these brines are used in commercial air conditioning systems.


 
Food and Nutrition: lithium
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Metal not known to have any physiological function, although it occurs in food and water; lithium salts are used in the treatment of manic-depressive illness.

 
Drug Info: Lithium
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Brand names: Eskalith CR®Eskalith®Lithobid®Lithonate®

Chemical formula:



Lithium Carbonate Oral capsule

What is this medicine?

LITHIUM is used to prevent and treat the manic episodes caused by manic-depressive illness.

This medicine may be used for other purposes; ask your health care provider or pharmacist if you have questions.

What should I tell my health care provider before I take this medicine?

They need to know if you have any of these conditions:
•dehydration (diarrhea or sweating)
•heart or blood vessel disease
•kidney disease
•low level of salt in the blood, or on a low salt diet
•an unusual or allergic reaction to lithium, other medicines, foods, dyes, or preservatives
•pregnant or trying to get pregnant
•breast-feeding

How should I use this medicine?

Take this medicine by mouth with a glass of water. Follow the directions on the prescription label. Take after a meal or snack to avoid stomach upset. Take your doses at regular intervals. Do not take your medicine more often than directed. The amount of this medicine you take is very important. Taking more than the prescribed dose can cause serious side effects. Do not stop taking except on the advice of your doctor or health care professional.

Talk to your pediatrician regarding the use of this medicine in children. Special care may be needed. While this drug may be prescribed for children as young as 12 years for selected conditions, precautions do apply.

Overdosage: If you think you have taken too much of this medicine contact a poison control center or emergency room at once.
NOTE: This medicine is only for you. Do not share this medicine with others.

What if I miss a dose?

If you miss a dose, take it as soon as you can. If it is almost time for your next dose, take only that dose. Do not take double or extra doses. You must leave a suitable interval between doses. If you are taking one dose a day and have to take a missed dose, make sure there is at least 10 to 12 hours between doses. If you are taking two doses a day and have to take a missed dose, make sure there is at least 5 to 6 hours between doses.

What may interact with this medicine?

Do not take this medicine with any of the following medications:
•stimulant medicines used to treat ADHD or narcolepsy

This medicine may also interact with the following medications:
•caffeine
•calcium iodide
•carbamazepine
•diuretics
•medicines for high blood pressure
•medicines for mental problems and psychotic disturbances
•metronidazole
•NSAIDs, medicines for pain and inflammation, like ibuprofen or naproxen
•phenytoin
•potassium iodide, KI
•sodium bicarbonate
•sodium chloride
•urea

This list may not describe all possible interactions. Give your health care provider a list of all the medicines, herbs, non-prescription drugs, or dietary supplements you use. Also tell them if you smoke, drink alcohol, or use illegal drugs. Some items may interact with your medicine.

What should I watch for while using this medicine?

Visit your doctor or health care professional for regular checks on your progress. It can take several weeks of treatment before you start to get better.

The amount of salt (sodium) in your body influences the effects of this medicine, and this medicine can increase salt loss from the body. Eat a normal diet that includes salt. Do not change to salt substitutes. Avoid changes involving diet, or medications that include large amounts of sodium like sodium bicarbonate. Ask your doctor or health care professional for advice if you are not sure.

Drink plenty of fluids while you are taking this medicine. Avoid drinks that contain caffeine, such as coffee, tea and colas. You will need extra fluids if you have diarrhea or sweat a lot. This will help prevent toxic effects from this medicine. Be careful not to get overheated during exercise, saunas, hot baths, and hot weather. Consult your doctor or health care professional if you have a high fever or persistent diarrhea.

You may get drowsy or dizzy. Do not drive, use machinery, or do anything that needs mental alertness until you know how this medicine affects you. Do not stand or sit up quickly, especially if you are an older patient. This reduces the risk of dizzy or fainting spells.

What side effects may I notice from receiving this medicine?

Side effects that you should report to your doctor or health care professional as soon as possible:
•blurred vision
•clumsiness or loss of balance
•confusion
•difficulty speaking or swallowing
•dizziness
•loss of appetite
•muscle weakness
•nausea, vomiting
•pain, coldness, or blue coloration of fingers or toes
•sensitivity to cold
•seizures
•slow, fast, or irregular heartbeat (palpitations)
•slurred speech
•swelling in the neck
•unusually weak or tired
•unusual weight gain

Side effects that usually do not require medical attention (report to your doctor or health care professional if they continue or are bothersome):
•increased thirst
•increased frequency and urgency to pass urine
•muscle twitches
•skin rash
•stomach bloating, full feeling
•trembling of the hands

This list may not describe all possible side effects. Call your doctor for medical advice about side effects. You may report side effects to FDA at 1-800-FDA-1088.

Where should I keep my medicine?

Keep out of the reach of children.

Store at room temperature between 15 and 30 degrees C (59 and 86 degrees F). Throw away any unused medicine after the expiration date.

Last updated: 7/1/2002

Important Disclaimer: The drug information provided here is for educational purposes only. It is intended to supplement, not substitute for, the diagnosis, treatment and advice of a medical professional. This drug information does not cover all possible uses, precautions, side effects and interactions. It should not be construed to indicate that this or any drug is safe for you. Consult your medical professional for guidance before using any prescription or over the counter drugs.

 

Chemical element, lightest alkali metal, chemical symbol Li, atomic number 3. It is soft, white, lustrous, and very reactive, forming compounds in which it has valence 1. The metal is used in certain alloys, as a coolant in nuclear reactors, and (because of its reactivity) as a reagent, scavenger, and rocket fuel. Lithium hydride is used as a source of hydrogen; lithium hydroxide is used as an additive in storage batteries and to absorb carbon dioxide. Halides (see halogen) of lithium are used as moisture absorbents, and lithium soaps are used as thickeners in lubricating greases. Lithium carbonate is an important drug for treating depression and bipolar disorder.

For more information on lithium, visit Britannica.com.

 
lithium (lĭth'ēəm) [Gr.,=stone], metallic chemical element; symbol Li; at. no. 3; at. wt. 6.941; m.p. about 180.54°C; b.p. about 1,342°C; sp. gr. .534 at 20°C; valence +1. Lithium is a soft, silver-white metal. It is one of the alkali metals in Group 1 of the periodic table. It is the least dense metal. Because it has high specific heat, it has found some use in cooling systems for nuclear reactors; such use is limited because lithium is very corrosive. Lithium metal is prepared by electrolysis of fused lithium chloride. Lithium reacts with water less readily than sodium. It burns in air with a brilliant white flame. Lithium forms many inorganic compounds, among them a hydride (LiH), a nitride (Li3N), an oxide (lithia, Li2O), a hydroxide (LiOH), a carbide (Li2C2), a carbonate (Li2CO3), and a phosphate (Li3PO4). When heated it reacts directly with the halogens to form halides. Lithium aluminum hydride (LiAlH4) is an important reagent in organic chemistry. Lithium also forms numerous organic compounds. One compound of major importance is lithium stearate, produced by cooking tallow (or other animal fat) with lithium hydroxide; lithium stearate is used to transform oil into lithium-base lubricating greases, which have found extensive use in the automotive industry. Lithium carbonate is used in special glasses and ceramic glazes. Lithium chloride and bromide are used as brazing and welding fluxes; they are also used in air conditioning systems because they are very hygroscopic, i.e., they absorb moisture. Lithium hydroxide is used to increase the capacity of alkaline storage cells. Lithium compounds are used in the nuclear energy industry, in the preparation of plastics and synthetic rubber, and in the synthesis of vitamin A. Lithium is added in small amounts to magnesium, aluminum, or lead-base alloys; it is also used as a degasifier in iron, steel, and copper refining. In addition, lithium is used to scavenge small amounts of oxygen and nitrogen in electronic vacuum tubes. Trace amounts of lithium and its compounds color a flame bright red; they are used in pyrotechnics. Lithium in the salt form has recently come into use as a medical treatment for bipolar disorder. Lithium is widely distributed in nature; it is found in the soil, in plants, in animals, and in the human body. It is also found in the sun. Lithium may be profitably extracted from ores containing as little as 1% lithium (measured as lithium oxide). Some commercially important minerals are lepidolite, petalite, spodumene, and amblygonite. Lithium is also produced from brines such as those in Searles Lake, Calif., and in the Great Salt Lake, Utah. Lithium was discovered in 1817 by J. A. Arfvedson.


 

A chemical element, atomic number 3, atomic weight 6.939, symbol Li.

  • l. carbonate — used in the treatment of canine cyclic hematopoiesis to stabilize numbers of neutrophils.
 
Wikipedia: Lithium
Top
heliumlithiumberyllium
H

Li

Na
Appearance
silvery white (seen here in oil)
General
Name, symbol, number lithium, Li, 3
Element category alkali metal
Group, period, block 12, s
Standard atomic weight 6.941(2)g·mol−1
Electron configuration 1s2 2s1
Electrons per shell 2, 1 (Image)
Physical properties
Phase solid
Density (near r.t.) 0.534 g·cm−3
Liquid density at m.p. 0.512 g·cm−3
Melting point 453.69 K
(180.54 °C, 356.97 °F)
Boiling point 1615 K
(1342 °C, 2456.6 °F)
Critical point (extrapolated)
3223 K, 67 MPa
Heat of fusion 3.00 kJ·mol−1
Heat of vaporization 147.1 kJ·mol−1
Specific heat capacity (25 °C) 24.860 J·mol−1·K−1
Vapor pressure
P/Pa 1 10 100 1 k 10 k 100 k
at T/K 797 885 995 1144 1337 1610
Atomic properties
Oxidation states +1, -1
(strongly basic oxide)
Electronegativity 0.98 (Pauling scale)
Ionization energies 1st: 520.2 kJ·mol−1
2nd: 7298.1 kJ·mol−1
3rd: 11815.0 kJ·mol−1
Atomic radius 152 pm
Covalent radius 128±7 pm
Van der Waals radius 182 pm
Miscellaneous
Crystal structure body-centered cubic
Magnetic ordering paramagnetic
Electrical resistivity (20 °C) 92.8 nΩ·m
Thermal conductivity (300 K) 84.8 W·m−1·K−1
Thermal expansion (25 °C) 46 µm·m−1·K−1
Speed of sound (thin rod) (20 °C) 6000 m/s
Young's modulus 4.9 GPa
Shear modulus 4.2 GPa
Bulk modulus 11 GPa
Mohs hardness 0.6
CAS registry number 7439-93-2
Most stable isotopes
Main article: Isotopes of lithium
iso NA half-life DM DE (MeV) DP
6Li 7.5% 6Li is stable with 3 neutrons
7Li 92.5% 7Li is stable with 4 neutrons
6Li content may be as low as 3.75% in
natural samples. 7Li would therefore
have a content of up to 96.25%.
References
Lithium ingots with a thin layer of black oxide tarnish

Lithium (pronounced /ˈlɪθiəm/) is the chemical element with atomic number 3, and is represented by the symbol Li. It is a soft alkali metal with a silver-white color. Under standard conditions it is the lightest metal and the least dense solid element. Like all alkali metals lithium is highly reactive, corroding quickly in moist air to form a black tarnish. For this reason lithium metal is typically stored under the cover of oil.[1] When cut open lithium exhibits a metallic lustre, but contact with oxygen quickly turns it back to a dull silvery grey color. Lithium is highly flammable.

According to theory lithium was one of the few elements synthesized in the Big Bang; its abundance is now vastly less than that predicted by theory[2]; the processes by which new lithium is created and destroyed, and the true value of its abundance,[3] continue to be active matters of study in astronomy.[4][5][6] Though very light in atomic weight, lithium is less common in the universe than any of the first 20 elements due to its low nuclear binding energy.

Due to its high reactivity it only appears naturally in the form of compounds. Lithium occurs in a number of pegmatitic minerals, but is also commonly obtained from brines and clays; on a commercial scale, lithium metal is isolated electrolytically from a mixture of lithium chloride and potassium chloride.

Trace amounts of lithium are present in the oceans and in some organisms, though the element serves no apparent vital biological function in humans, though the neurological effect of the lithium ion Li+ makes some lithium salts useful as a mood stabilizing drug. Lithium and its compounds have several industrial applications, including heat-resistant glass and ceramics, high strength-to-weight alloys used in aircraft, and lithium batteries. Lithium also has important links to nuclear physics. the splitting of lithium atoms was the first man-made form of a nuclear reaction, and lithium deuteride serves as a fusion fuel in staged thermonuclear weapons.

Contents

History and etymology

Petalite (lithium aluminium silicate) was first discovered in 1800 by the Luso-Brazilian scientist José Bonifácio de Andrade e Silva, who discovered the mineral in a Swedish iron mine on the island of Utö. However, it was not until 1817 that Johan August Arfwedson, then a trainee in the laboratory of Jöns Jakob Berzelius, discovered the presence of a new element while analyzing petalite ore. The element formed compounds similar to those of sodium and potassium, though its carbonate and hydroxide were less water soluble and had a larger capacity to neutralize acid. Berzelius gave the alkaline material the name "lithos", from the Greek λιθoς (lithos, "stone"), to reflect its discovery in a mineral, as opposed to sodium and potassium which had been discovered in plant tissue; its name was later standardized as "lithium". Arfwedson later showed that this same element was present in the mineral ores spodumene and lepidolite. In 1818 Christian Gmelin was the first to observe that lithium salts give a bright red color in flame. However, both Arfwedson and Gmelin tried and failed to isolate the element from its salts.[7][8][9] The element was not isolated until 1821, when William Thomas Brande performed electrolysis on lithium oxide, a process previously employed by Sir Humphry Davy to isolate potassium and sodium.[8][10] Brande also described pure salts of lithium, such as the chloride, and performed an estimate of its atomic weight. In 1855 Robert Bunsen and Augustus Matthiessen produced large quantities of the metal by electrolysis of lithium chloride. Commercial production of lithium metal began in 1923 by the German company Metallgesellschaft AG through the electrolysis of a molten mixture of lithium chloride and potassium chloride.[7][11]

Properties

Lithium pellets (covered in white lithium hydroxide)

Like other alkali metals, lithium has a single weakly held valence electron which it will readily lose to form a cation. As a result lithium is easily deformed, highly reactive, and has a low melting and boiling point. The properties and chemistry of Lithium are modified further due to its small atomic radius or ionic radius. Lithium is the least electropositive of the alkali group.

Lithium is soft enough to be cut with a knife with some difficulty; it is the hardest of the alkali metals. The fresh metal has a silvery-white color which remains untarnished only in dry air.[12] Lithium has about half the density of water, similar to pine wood, and lithium sticks have a heft similar to wooden dowels. Lithium floats in hydrocarbons, so laboratory stock lithium sticks are typically held under the protective liquid by the container lid.

Lithium possesses a low coefficient of thermal expansion and the highest specific heat capacity of any solid element. Lithium is superconductive below 400 μK at standard pressure[13] and at higher temperatures (more than 9 Kelvin) at very high pressures (over 200,000 atmospheres)[14]

At cryogenic temperatures, lithium, like sodium, undergoes diffusionless phase change transformations. At 4.2K it has a Rhombohedral (with a nine-layer repeat spacing)[15], at higher temperatures it transforms to Face-centered cubic and then Body-centered cubic. At liquid-helium temperatures (4 K) the rhombohedral structure is the most prevalent.

Chemistry

In moist air, lithium metal rapidly tarnishes to form a black coating of lithium hydroxide (LiOH and LiOH·H2O), lithium nitride (Li3N) and lithium carbonate (Li2CO3, the result of a secondary reaction between LiOH and CO2).[12]

When placed over a flame lithium gives off a striking crimson color, but when it burns strongly the flame becomes a brilliant white. Lithium will ignite and burn in oxygen when exposed to water or water vapours. It is the only metal that reacts with nitrogen at room temperature.[clarification needed]

Lithium metal is flammable and potentially explosive when exposed to air and especially water, though less so than other alkali metals. The lithium-water reaction at normal temperatures is brisk but not violent, though the hydrogen produced can ignite. Like all alkali metals, lithium fires are difficult to extinguish, requiring dry powder fire extinguishers, specifically Class D type (see Types of extinguishing agents).

Lithium compounds

Lithium has a diagonal relationship with magnesium, an element of similar atomic and ionic radius. Chemical resemblances between the two metals include the formation of a nitride by reaction with N2, the formation of an oxide when burnt in O2, salts with similar solubilities, and thermal instability of the carbonates and nitrides.[12]

Isotopes

Naturally occurring lithium is composed of two stable isotopes, 6Li and 7Li, the latter being the more abundant (92.5 percent natural abundance).[16] Both natural isotopes have anomalously low nuclear binding energy per nucleon compared to the next lighter and heavier elements, Helium and Beryllium, which means that alone among stable light elements, Lithium can produce net energy through nuclear fission. Seven radioisotopes have been characterized, the most stable being 8Li with a half-life of 838 ms and 9Li with a half-life of 178.3 ms. All of the remaining radioactive isotopes have half-lives that are shorter than 8.6 ms. The shortest-lived isotope of lithium is 4Li, which decays through proton emission and has a half-life of 7.58043x10-23 s.

7Li is one of the primordial elements (or, more properly, primordial isotopes) produced in Big Bang nucleosynthesis. A small amount of both 6Li and 7Li are produced in stars, but are thought to be burned as fast as it is produced.[17] Lithium isotopes fractionate substantially during a wide variety of natural processes, including mineral formation (chemical precipitation), metabolism, and ion exchange. Lithium ions substitute for magnesium and iron in octahedral sites in clay minerals, where 6Li is preferred to 7Li, resulting in enrichment of the light isotope in processes of hyperfiltration and rock alteration. The exotic 11Li is known to exhibit a nuclear halo.

Natural occurrence

Lithium is about as common as chlorine in the Earth's upper continental crust, on a per-atom basis.

According to theory, the stable isotopes lithium-6 and lithium-7 were created in the Big Bang, but the amounts are unclear. There is general agreement that they were larger than the cosmos contains today. Because of the method by which elements are built up by fusion in stars, there is a general trend in the cosmos that the lighter elements are more common. However, lithium (element number 3) is tied with krypton as the 32nd/33rd most abundant element in the cosmos (see Cosmochemical Periodic Table of the Elements in the Solar System), being less common than any element before scandium (element 21). It is not until atomic number 36 (krypton) and beyond that chemical elements are found to be universally less common in the cosmos than lithium. The reasons have to do with the failure of any good mechanisms to synthesize lithium in the fusion reactions between nuclides in supernovae. Due to the absence of any quasi-stable nuclide with five nucleons, nuclei of lithium-5 produced from helium and a proton has no time to fuse with a second proton or neutron to form a six nucleon isotope which might decay to lithium-6, even under extreme conditions of bombardment. Also, the product of helium-helium fusion (berylium-8) is immediately unstable toward disintegration to helium again, and is thus not available for formation of lithium. Some lithium-7 is formed in the pp III branch of the proton-proton chain in main sequence and red giant stars, but it is normally consumed by lithium burning as fast as it is formed. This leaves new formation of the stable isotopes lithium 6 and 7 to rare cosmic ray spallation on carbon or other elements in cosmic dust. Meanwhile, existing Li-6 and Li-7 is destroyed in many nuclear reactions in supernovae and by lithium burning in main sequence stars, resulting in net removal of lithium from the cosmos.

Lithium is widely distributed on Earth[18] but does not naturally occur in elemental form due to its high reactivity. Estimates for crustal content range from 20 to 70 ppm by weight.[12] In keeping with its name, lithium forms a minor part of igneous rocks, with the largest concentrations in granites. Granitic pegmatites also provide the greatest abundance of lithium-containing minerals, with spodumene and petalite being the most commercially viable sources.[12] A newer source for lithium is hectorite clay, the only active development of which is through Western Lithium Corp. in the USA. [19]

According to the Handbook of Lithium and Natural Calcium, "Lithium is a comparatively rare element, although it is found in many rocks and some brines, but always in very low concentrations. There are a fairly large number of both lithium mineral and brine deposits but only comparatively a few of them are of actual or potential commercial value. Many are very small, others are too low in grade."[20] The most important deposit of lithium is in the Salar de Uyuni area of Bolivia, which holds half of the world's known reserves. According to the US Geological Survey the reserves of lithium in Bolivia are estimated at 5.4 million tons, compared with 3 million tons in Chile, 1.1 million tons in China and just 410,000 tons in the United States.[21][22] Worldwide lithium reserves are estimated at 30 million tonnes in 2015.[23]

Seawater contains an estimated 230 billion tons of lithium, though at a low concentration of 0.1 to 0.2 ppm.[24]

Major applications of the metal

Because of its specific heat capacity, the highest of all solids, lithium is often used in heat transfer applications.

In the latter years of the 20th century lithium became important as an anode material. Used in lithium-ion batteries because of its high electrochemical potential, a typical cell can generate approximately 3 volts, compared with 1.5 volts for lead/acid or zinc cells. Because of its low atomic mass, it also has a high charge- and power-to-weight ratio.

Lithium is also used in the pharmaceutical and fine-chemical industry in the manufacture of organolithium reagents, which are used both as strong bases and as reagents for the formation of carbon carbon bonds. Organolithiums are also used in polymer synthesis as catalysts/initiators[25] in anionic polymerisation of unfunctionalised olefins.[26][27][28]

Medical use

Lithium salts were used during the 19th century to treat gout. Lithium salts such as lithium carbonate (Li2CO3), lithium citrate, and lithium orotate are mood stabilizers. They are used in the treatment of bipolar disorder since, unlike most other mood altering drugs, they counteract both mania and depression. Lithium can also be used to augment antidepressants. Because of Lithium's nephrogenic diabetes insipidus effects, it can be used to help treat the syndrome of inappropriate diuretic hormone (SIADH). It was also sometimes prescribed as a preventive treatment for migraine disease and cluster headaches.[29]

The active principle in these salts is the lithium ion Li+. Although this ion has a smaller diameter than either Na+ or K+, in a watery environment like the cytoplasmic fluid, Li+ binds to the hydrogen atoms of water, making it effectively larger than either Na+ or K+ ions. How Li+ works in the CNS is still a matter of debate. Li+ elevates brain levels of tryptophan, 5-HT (serotonin), and 5-HIAA (a serotonin metabolite). Serotonin is related to mood stability. Li+ also reduces catecholamine activity in the brain (associated with brain activation and mania), by enhancing reuptake and reducing release. Therapeutically useful amounts of lithium (~ 0.6 to 1.2 mmol/l) are only slightly lower than toxic amounts (>1.5 mmol/l), so the blood levels of lithium must be carefully monitored during treatment to avoid toxicity.

Common side effects of lithium treatment include muscle tremors, twitching, ataxia[30] and hypothyroidism. Long term use is linked to hyperparathyroidism[31], hypercalcemia (bone loss), hypertension, kidney damage, nephrogenic diabetes insipidus (polyuria and polydipsia), seizures[32] and weight gain.[33] Some of the side-effects are a result of the increased elimination of potassium.

There appears to be an increased risk of Ebstein (cardiac) Anomaly in infants born to women taking lithium during the first trimester of pregnancy.

According to a study in 2009 at Oita University in Japan and published in the British Journal of Psychiatry, communities whose water contained larger amounts of lithium had significantly lower suicide rates[34][35][36][37] but did not address whether lithium in drinking water causes the negative side effects associated with higher doses of the element.[38]

Other uses

The red lithium flame leads to Lithium's use in flares and pyrotechnics
  • Electrical and electronic uses:
  • Chemical uses:
  • General engineering:
  • Alloys of the metal with aluminium, cadmium, copper and manganese are used to make high-performance aircraft parts (see also Lithium-aluminium alloys).
  • Optics:
  • Rocketry:
  • Nuclear applications:
    • Lithium deuteride was the fusion fuel of choice in early versions of the hydrogen bomb. When bombarded by neutrons, both 6Li and 7Li produce tritium—this reaction, which was not fully understood when hydrogen bombs were first tested, was responsible for the runaway yield of the Castle Bravo nuclear test. Tritium fuses with deuterium in a fusion reaction that is relatively easy to achieve. Although details remain secret, lithium-6 deuteride still apparently plays a role in modern nuclear weapons, as a fusion material.
    • Lithium fluoride (highly enriched in the common isotope lithium-7) forms the basic constituent of the preferred fluoride salt mixture (LiF-BeF2) used in liquid-fluoride nuclear reactors. Lithium fluoride is exceptionally chemically stable and LiF/BeF2 mixtures have low melting points and the best neutronic properties of fluoride salt combinations appropriate for reactor use.[clarification needed]
    • In conceptualized nuclear fusion power plants, Lithium will be used to produce tritium in magnetically confined reactors using deuterium and tritium as the fuel. Tritium does not occur naturally and will be produced by surrounding the reacting plasma with a 'blanket' containing lithium where neutrons from the deuterium-tritium reaction in the plasma will react with the lithium to produce more tritium. 6Li + n → 4He + 3H. Various means of doing this will be tested at the ITER reactor being built at Cadarache, France.
    • Lithium is used as a source for alpha particles, or helium nuclei. When 7Li is bombarded by accelerated protons 8Be is formed, which undergoes spontaneous fission to form two alpha particles. This was the first man-made nuclear reaction, produced by Cockroft and Walton in 1929.
  • Other uses:
    • Lithium hydroxide (LiOH) is an important compound of lithium obtained from lithium carbonate (Li2CO3). It is a strong base, and when heated with a fat it produces a lithium soap. Lithium soap has the ability to thicken oils and is used to manufacture lubricating greases.
    • Lithium hydroxide and lithium peroxide are used in confined areas, such as aboard spacecraft and submarines, for air purification. Lithium hydroxide absorbs carbon dioxide from the air by reacting with it to form lithium carbonate, and is preferred over other alkaline hydroxides for its low weight. Lithium peroxide (Li2O2) in presence of moisture not only absorbs carbon dioxide to form lithium carbonate, but also releases oxygen. E.g. 2 Li2O2 + 2 CO2 → 2 Li2CO3 + O2.
    • Lithium compounds are used in red fireworks and flares.
    • The Mark 50 Torpedo Stored Chemical Energy Propulsion System (SCEPS) uses a small tank of sulfur hexafluoride gas which is sprayed over a block of solid lithium. The reaction generates enormous heat which is used to generate steam from seawater. The steam propels the torpedo in a closed Rankine cycle.[41]

Production and world supply

Lithium mine, Salar del Hombre Muerto, Argentina. The brine in this salar is rich in lithium, and the mine concentrates the brine by pumping it into solar evaporation ponds. 2009 image from NASA’s EO-1 satellite.

Since the end of World War II lithium metal production has greatly increased. The metal is separated from other elements in igneous minerals such as those above. Lithium salts are extracted from the water of mineral springs, brine pools and brine deposits.

The metal is produced electrolytically from a mixture of fused lithium and potassium chloride. In 1998 it was about US$ 43 per pound ($95 per kg).[42]

Deposits of lithium are found in South America throughout the Andes mountain chain. Chile is the leading lithium metal producer, followed by Argentina. Both countries recover the lithium from brine pools. In the United States lithium is recovered from brine pools in Nevada.[43]

Nearly half the world's known reserves are located in Bolivia, a nation sitting along the central eastern slope of the Andes. In 2009 Bolivia is negotiating with Japanese and French firms to begin extraction.[44] According to the US Geological Survey, Bolivia's Uyuni Desert has 5.4 million tons of lithium, which can be used to make batteries for hybrid and electric vehicles.[44] This is the largest amount of lithium in any country, compared to Chile's 3 million tons and the United States's 760,000 tons.[44][45]

China may emerge as a significant producer of brine-source lithium carbonate around 2010. There is potential production of up to 55,000 tons per year if projects in Qinghai province and Tibet proceed.[46]

The total amount of lithium recoverable from global reserves has been estimated at 35 million tonnes, which includes 15 million tons of the known global lithium reserve base.[47]

In 1976 a National Research Council Panel estimated lithium resources at 10.6 million tons for the Western World.[48] With the inclusion of Russian and Chinese resources as well as new discoveries in Australia, Serbia, Argentina and the United States, the total had nearly tripled by 2008.[49][50]

Precautions

Due to its alkaline tarnish, lithium metal is corrosive and requires special handling to avoid skin contact. Breathing lithium dust or lithium compounds (which are often alkaline) initially irritate the nose and throat, while higher exposure can cause a buildup of fluid in the lungs, leading to pulmonary edema. The metal itself is a handling hazard because of the caustic hydroxide produced when it is in contact with moisture. Lithium is safely stored in non-reactive compounds such as naphtha.[51]

Regulation

Some jurisdictions limit the sale of lithium batteries, which are the most readily available source of lithium metal for ordinary consumers. Lithium can be used to reduce pseudoephedrine and ephedrine to methamphetamine in the Birch reduction method, which employs solutions of alkali metals dissolved in anhydrous ammonia.

Carriage and shipment of some kinds of lithium batteries may be prohibited aboard certain types of transportation (particularly aircraft) because of the ability of most types of lithium batteries to fully discharge very rapidly when short-circuited, leading to overheating and possible explosion in a process called thermal runaway. Most consumer lithium batteries have thermal overload protection built-in to prevent this type of incident, or their design inherently limits short-circuit currents. Internal shorts have been known to develop due to manufacturing defects or damage to batteries that can lead to spontaneous thermal runaway.[52]

See also

References

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External links


 
Translations: Lithium
Top

Dansk (Danish)
n. - lithium

Nederlands (Dutch)
lithium (chemisch element)

Français (French)
n. - lithium

Deutsch (German)
n. - (Chem.) Lithium

Ελληνική (Greek)
n. - (χημ.) λίθιο(ν)

Italiano (Italian)
litio

Português (Portuguese)
n. - lítio (m) (Quím.)

Русский (Russian)
литий

Español (Spanish)
n. - litio

Svenska (Swedish)
n. - (kem.) litium

中文(简体)(Chinese (Simplified))

中文(繁體)(Chinese (Traditional))
n. - 鋰

한국어 (Korean)
n. - 리튬

日本語 (Japanese)
n. - リチウム, リチウム塩

العربيه (Arabic)
‏(الاسم) الليثيوم, عنصر فلزي فضي البياض‏

עברית (Hebrew)
n. - ‮היסוד המתכתי הקל ביותר, משמש כחומר לסוללות וברפואה, ליתיום, אבנן‬


 
 

 

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