Lithium CAS 7439-93-2
Introduction：Basic information about Lithium CAS 7439-93-2, including its chemical name, molecular formula, synonyms, physicochemical properties, and safety information, etc.
Lithium Basic informationIntroduction Physical Properties Uses Production Reactions
| Product Name: | Lithium |
| Synonyms: | elementallithium;Li;lithium,elemental;LITHIUM AA STANDARD;LITHIUM, AAS STANDARD SOLUTION;LITHIUM AA SINGLE ELEMENT STANDARD;LITHIUM ATOMIC ABSORPTION SINGLE ELEMENT STANDARD;LITHIUM ATOMIC SPECTROSCOPY STANDARD |
| CAS: | 7439-93-2 |
| MF: | Li |
| MW: | 6.94 |
| EINECS: | 231-102-5 |
| Product Categories: | Chemical Synthesis;Electrode Materials;Lithium-Ion Batteries;Materials Science;Metal and Ceramic Science;Synthetic Reagents;metal or element;Alkali Metals;Alternative Energy;Anode Materials;Battery Electrode Materials;Inorganic Chemicals;Alkali MetalsMetal and Ceramic Science;Lithium;Metals;Reduction;Synthetic Reagents;Alkali MetalsMaterials Science;Alternative Energy;Electrode MaterialsMetal and Ceramic Science;LSpectroscopy;AA Standard SolutionsSpectroscopy;AAS;Alphabetic;ChlorideAnalytical Standards;Matrix Selection;Reference/Calibration Standards;Single Solution;Standard Solutions;Inorganics |
| Mol File: | 7439-93-2.mol |
Lithium Chemical Properties
| Melting point | 180 °C (lit.) |
| Boiling point | 1342 °C (lit.) |
| density | 0.534 g/mL at 25 °C (lit.) |
| vapor pressure | 1 hPa (723 °C) |
| storage temp. | water-free area |
| solubility | reacts with H2O |
| form | wire |
| color | Silvery |
| Specific Gravity | 0.534 |
| Flame Color | Pink-red or magenta |
| Odor | Odorless |
| PH | 5.0 (20°C in H2O) |
| PH Range | >12 |
| resistivity | 9.446 μΩ-cm, 20°C |
| Water Solubility | REACTS |
| Sensitive | air sensitive, moisture sensitive |
| Merck | 13,5542 |
| Exposure limits | ACGIH: TWA 2 ppm; STEL 4 ppm OSHA: TWA 2 ppm(5 mg/m3) NIOSH: IDLH 25 ppm; TWA 2 ppm(5 mg/m3); STEL 4 ppm(10 mg/m3) |
| Stability: | Stable, but reacts violently with water. |
| InChI | 1S/Li |
| InChIKey | WHXSMMKQMYFTQS-UHFFFAOYSA-N |
| SMILES | [Li] |
| LogP | -0.77 at 25℃ |
| Hardness, Vickers | <= 5.0 |
| CAS DataBase Reference | 7439-93-2(CAS DataBase Reference) |
| NIST Chemistry Reference | Lithium(7439-93-2) |
| EPA Substance Registry System | Lithium (7439-93-2) |
Safety Information
| Hazard Codes | Xi,C,F |
| Risk Statements | 36/38-34-14/15-23 |
| Safety Statements | 8-43-45-43C-36/37/39-26 |
| RIDADR | UN 3264 8/PG 3 |
| WGK Germany | 2 |
| RTECS | OJ5540000 |
| F | 10 |
| Autoignition Temperature | 179oC |
| TSCA | TSCA listed |
| HS Code | 2805 19 90 |
| HazardClass | 4.3 |
| PackingGroup | I |
| Storage Class | 13 - Non Combustible Solids |
| Hazardous Substances Data | 7439-93-2(Hazardous Substances Data) |
| Toxicity | An element used clinicallyas one of its salts. It is effective against both mania anddepression. Despite its effectiveness, there are no clearmechanisms that have been directly related to its therapeuticeffectiveness although its inhibition of the formation of inositolfrom inositol phosphate is thought to be important. At therapeuticconcentrations, lithium causes almost no discerniblepsychotropic effects in healthy humans. The major complaintswhen the serum concentrations of the drug are carefully monitoredinclude slight muscular weakness, thirst, and excessiveurination. The major difficulty with lithium is that a fairly highconcentration of the ion is needed in the blood(0.5_x0002_1.0 mmol/L) for maintenance, higher for acute mania.Toxic symptoms (which can involve many physiologicalsymptoms) may occur, however, at doses of 1.5 mmol/L orhigher. This low therapeutic index is indicative of the need forregular monitoring of lithium concentrations in the serum. |
| Introduction | Lithium, the lightest of the alkali metals, has an atomic number of 3 and an atomic weight of 6.94. Lithium exhibits oxidation states of zero and plus one and is clearly an alkali metal. However, lithium and its compounds are not always typical of the other alkali metals. The high ionic charge density and the strong tendency for lithium to form a monopositive ion strongly influence the stability of lithium compounds and the type of bond which lithium forms with other atoms, ions and radicals. The unusually high charge density of the lithium ion is a crucial factor in setting lithium and its compounds apart from the other alkali metals and their compounds. Lithium has numerous industrial applications. It is used to make highenergy lithium batteries. Lithium and its aluminum alloys are used as anodes in non-aqueous solid-state batteries. Also, many of its salts are used as electrolytes in these batteries. Another major application is in metallurgy. Lithium is alloyed with lead, magnesium, aluminum and other metals. Its alloy Bahnmetall is used for wheel bearings in railroad cars, and its magnesium alloy is used in aerospace vehicles. Probably the most important applications of lithium are in preparative chemistry. It is the starting material to prepare lithium hydride, amide, nitride, alkyls and aryls. Lithium hydrides are effective reducing agents. The alkyls are used in organic syntheses. |
| Physical Properties | Soft silvery-white metal; body-centered cubic structure; density 0.531 g/cm3; burns with a carmine-red flame, evolving dense white fumes; melts at 180.54°C; vaporizes at 1,342°C; vapor pressure 1 torr at 745°C and 10 torr at 890°C; electrical resistivity 8.55 microhm-cm at 0°C and 12.7 microhm-cm at 100°C; viscosity 0.562 centipoise at 200°C and 0.402 centipoise at 400°C; reacts with water; soluble in liquid ammonia forming a blue solution. |
| Uses | Lithium is used as a starting material or chemical intermediate in a number of reactions. Both lithium hydride and lithium nitride, which are prepared by the direct combination of their elements. Other inorganic compounds of lithium may be prepared by the direct combination of the elements or by the reaction of lithium with an acid gas if an unusually high purity product is required. For example, extremely dry lithium sulfide can be prepared by the reaction of lithium metal with hydrogen sulfide. Such processes are hazardous and expensive and are generally avoided if possible. Lithium may be used in the preparation of both alkyl- and aryl-lithium compounds and lithium alkoxides. |
| Production | Lithium is obtained primarily from its ore, spodumene. Another important source is natural brine found in many surface and ground waters, from which the metal also is produced commercially. |
| Reactions | Lithium metal is highly reactive but less so than other alkali metals. Its chemical properties, however, are more like those of the alkaline earth metals. At ordinary temperatures, lithium does not react with dry oxygen. However, it reacts above 100°C, forming lithium oxide, Li2O: The metal ignites in air near its melting point, burning with intense white flame, forming Li2O. Lithium reacts with water forming lithium hydroxide with evolution of hydrogen: 2Li + 2H2O → 2LiOH + H2 The reaction is violent when lithium metal is in finely divided state. Lithium reacts violently with dilute acids, liberating hydrogen: Li + 2HCl → LiCl + H2 Reaction with cold concentrated sulfuric acid is slow. The metal dissolves in liquid ammonia, forming a blue solution, lithium amide, LiNH2: 2Li + 2NH3 → 2LiNH2 + H2 The same product also is obtained from ammonia gas. Unlike other alkali metals, lithium reacts with nitrogen in the presence of moisture at ordinary temperatures, forming the black lithium nitride, Li3N: 6Li + N2 → 2Li3N The above reaction is exothermic. Lithium reacts with hydrogen at red heat forming lithium hydride: Reactions with sulfur and selenium in liquid ammonia yield lithium sulfide and selenide, respectively: The metal combines with chlorine and other halogens, forming their halides: Li + Cl2 → 2LiCl When heated with carbon at 800°C, the product is lithium carbide: 2Li + 2C→Li2C2 2Li + S→Li2S 2Li + H2→2LiH Li + O2→2Li2O The metal reacts with carbon dioxide at elevated temperatures, forming lithium carbonate, Li2CO3. Lithium forms alloys with several metals including aluminum, calcium, copper, magnesium, mercury, sodium, potassium, silver, tin and zinc. It combines with phosphorus, arsenic and antimony on heating, forming their binary salts: The metal behaves as a reducing agent at high temperatures. It reduces aluminum chloride to aluminum and boron oxide to boron: Lithium liberates hydrogen from ethanol, forming lithium ethoxide: 2Li + 2C2H5OH → 2C2H5OLi + H2 Several organolithium compounds have important applications in organic syntheses. These may be readily synthesized by reactions of lithium with organics. The metal reacts with alkyl or aryl halides or mercury alkyls or aryls to produce alkyl or aryl lithium. |
| Description | Lithium is an alkali metal with physiologic actions similar topotassium and sodium. Lithium was discovered as a salt in1817 by Johan August Arfwedson. It does not occur in nature asa free metal but is found in minerals such as spodumene,petalite, and eucryptite. It is the 27th most abundant elementin Earth’s crust. Lithium was used to treat gout, as a saltsubstitute, and as a major constituent of the soft drink 7-Upbefore 1950. |
| Chemical Properties | Lithium is a silvery to grayish-white metal thatturns yellow on exposure to air and/or moisture. |
| Physical properties | In the metallic state, lithium is a very soft metal with a density of 0.534 g/cm3. When asmall piece is placed on water, it will float as it reacts with the water, releasing hydrogen gas.Lithium’s melting point is 179°C, and it has about the same heat capacity as water, with aboiling point of 1,342°C. It is electropositive with an oxidation state of +1, and it is an excellent conductor of heat and electricity. Its atom is the smallest of the alkali earth metals andthus is the least reactive because its valence electron is in the K shell, which is held closest toits nuclei. |
| Isotopes | There are two stable lithium isotopes: Li-6.015, which makes up 7.5% of all lithium atoms, and Li-7.016, which makes up 92.5% of lithium atoms found in the Earth’scrust. Less prevalent isotopes of lithium are Li-4, Li-5, Li-8, Li-9, Li-10, and Li-11. Theyare unstable with short half-lives and make up only a very small fraction of Lithium’stotal averaged atomic weight. |
| Origin of Name | The name lithium comes from the Greek word lithos, meaning “stone”because it was found in rocks on Earth. |
| Occurrence | Lithium ranks 33rd among the most abundant elements found on Earth. It does not existin pure metallic form in nature because it reacts with water and air. It is always combinedwith other elements in compound forms. These lithium mineral ores make up only about0.0007%, or about 65 ppm, of the Earth’s crust.Lithium is contained in minute amounts in the mineral ores of spodumene, lepidolite, andamblygonite, which are found in the United States and several countries in Europe, Africa, andSouth America. High temperatures are required to extract lithium from its compounds and byelectrolysis of lithium chloride. It is also concentrated by solar evaporation of salt brine in lakes.Metallic lithium is produced on a commercial scale by electrolysis of molten lithium chloride (LiCl) that is heated as a mixture with potassium chloride (KCl). Both have a rather highmelting point, but when mixed, the temperature required to melt them (400°C) is severalhundred degrees lower than their individual melting points. This liquid mixture of LiCl andKCl becomes the electrolyte. The anode is graphite (carbon) and the cathode is steel. Themolten liquid positive lithium cations collect at the cathode while negative chlorine anionscollect at the anode, and the potassium chloride remains in the electrolyte. Each positive ion oflithium that collects at the cathode gains an electron, thus producing neutral atoms of moltenlithium metal, which is then further purified. |
| History | Discovered by Arfvedson in 1817. Lithium is the lightest of all metals, with a density only about half that of water. It does not occur free in nature; combined it is found in small amounts in nearly all igneous rocks and in the waters of many mineral springs. Lepidolite, spodumene, petalite, and amblygonite are the more important minerals containing it. Lithium is presently being recovered from brines of Searles Lake, in California, and from Nevada, Chile, and Argentina. Large deposits of spodumene are found in North Carolina. The metal is produced electrolytically from the fused chloride. Lithium is silvery in appearance, much like Na and K, other members of the alkali metal series. It reacts with water, but not as vigorously as sodium. Lithium imparts a beautiful crimson color to a flame, but when the metal burns strongly the flame is a dazzling white. Since World War II, the production of lithium metal and its compounds has increased greatly. Because the metal has the highest specific heat of any solid element, it has found use in heat transfer applications; however, it is corrosive and requires special handling. The metal has been used as an alloying agent, is of interest in synthesis of organic compounds, and has nuclear applications. It ranks as a leading contender as a battery anode material because it has a high electrochemical potential. Lithium is used in special glasses and ceramics. The glass for the 200-inch telescope at Mt. Palomar contains lithium as a minor ingredient. Lithium chloride is one of the most hygroscopic materials known, and it, as well as lithium bromide, is used in air conditioning and industrial drying systems. Lithium stearate is used as an all-purpose and hightemperature lubricant. Other lithium compounds are used in dry cells and storage batteries. Seven isotopes of lithium are recognized. Natural lithium contains two isotopes. The metal is priced at about $1.50/g (99.9%). |
| Characteristics | While classified as an alkali metal, lithium also exhibits some properties of the alkali earthmetals found in group 2 (IIA). Lithium is the lightest in weight and softest of all the metalsand is the third lightest of all substances listed on the periodic table, with an average atomicweight of about 7. (The other two are hydrogen and helium.) Although it will float on water,it reacts with water, liberating explosive hydrogen gas and lithium hydroxide (2Li + 2H2O →2LiOH + H2?). It will also ignite when exposed to oxygen in moist air (4Li + O2 → 2Li2O).It is electropositive and thus an excellent reducing agent because it readily gives up electrons inchemical reactions. Lithium is the only metal that reacts with nitrogen at room temperature.When a small piece of the metal, which is usually stored in oil or kerosene, is cut, the newsurface has a bright, shiny, silvery surface that soon turns gray from oxidation. |
| Uses | Lithium has many uses in today’s industrial society. It is used as a flux to promote the fusing of metals during welding and soldering. It also eliminates the formation of oxides duringwelding by absorbing impurities. This fusing quality is also important as a flux for producingceramics, enamels, and glass.A major use is as lithium stearate for lubricating greases. It makes a solid grease that canwithstand hard use and high temperatures.Lithium is used to manufacture electric storage cells (batteries) that have a long shelf lifefor use in heart pacemakers, cameras, and so forth.Some lithium compounds are used as rocket propellants, nuclear reactor coolants, alloyhardeners, and deoxidizers and to make special ceramics.Lithium is a source of alpha particles when bombarded in a nuclear accelerator. The following occurs: lithium nuclei (3 protons + 4 neutrons) are targeted by high-speed protons(hydrogen nuclei), resulting in lithium absorbing a proton to form 4 protons + 4 neutrons,which become two alpha particles (helium nuclei). For example, see the following equation:1H + 3Li = 2He + 2He. This is an example of the first man-made nuclear reaction produced bySir John Cockcroft (1897–1967) and Ernest Walton (1903–1995) in 1929.Several compounds of lithium are used as pharmaceuticals to treat severe psychoticdepression (as antidepressant agents). And lithium carbonate is also used as a sedative ormild tranquilizer to treat less severe anxiety, which is a general feeling of uneasiness or distressabout present condition or future uncertainties. Lithium is also used in the production ofvitamin A. |
| Uses | In production of organometallic alkyl and aryl lithium compounds; in production of high-strength, low-density aluminum alloys for the aircraft industry; extremely tough, low-density alloys with aluminum and magnesium used for armour plate and aerospace components. In polymerization catalysts for the polyolefin plastics industry; manufacture of high-strength glass and glass-ceramics. As anode in electrochemical cells and batteries; as chemical intermediate in organic syntheses. Lithium stearate as thickener and gelling agent to transform oils into lubricating greases. |
| Uses | The most widely known use of lithium is that of lithiumcarbonate in treating manic-depressive affective disorders; themechanism by which it alleviates depression in some people isnot known. Industrially, lithium and its compounds are usedin metal alloys, lubricants, nuclear reactor coolant, ceramics,alkaline storage batteries, and electronic tubes. It is also used asa catalyst and as a reducing agent. |
| Definition | A light silvery moderately reactive metal; the first member of the alkali metals (group 1 of the periodic table). It occurs in a number of complex silicates, such as spodumene, lepidolite, and petalite, and a mixed phosphate, tryphilite. It is a rare element accounting for 0.0065% of the Earth’s crust. Lithium ores are treated with concentrated sulfuric acid and lithium sulfate subsequently separated by crystallization. The element can be obtained by conversion to the chloride and electrolysis of the fused chloride or of solutions of LiCl in pyridine. Lithium has the electronic configuration 1s22s1 and is entirely monovalent in its chemistry. The lithium ion is, however, much smaller than the ions of the other alkali metals; consequently it is polarizing and a certain degree of covalence occurs in its bonds. Lithium also has the highest ionization potential of the alkali metals. The element reacts with hydrogen to form lithium hydride, LiH, a colorless high-melting solid, which releases hydrogen at the anode during electrolysis (confirming the ionic nature Li+H–). The compound reacts with water to release hydrogen and is also frequently used as a reducing agent in organometallic synthesis. Lithium reacts with oxygen to give Li2O (sodium gives the peroxide) and with nitrogen to form Li3N on fairly gentle warming. The metal itself reacts only slowly with water, giving the hydroxide (LiOH) but lithium oxide reacts much more vigorously to give again the hydroxide; the nitride is hydrolyzed to ammonia. The metal reacts with halogens to form halides (LiX). Apart from the fluoride the halides are readily soluble both in water and in oxygen- containing organic solvents. In this property lithium partly resembles magnesium which has a similar charge/size ratio. Compared to the other carbonates of group 1, lithium carbonate is thermally unstable decomposing to Li2O and CO2. This is because the small Li+ ion leads to particularly high lattice energies favouring the formation of Li2O. Lithium also forms a wide range of alkyl- and aryl-compounds with organic compounds, which are particularly useful in organic synthesis. Lithium compounds impart a characteristic purple color to flames. Symbol: Li; m.p. 180.54°C; b.p. 1347°C; r.d. 0.534 (20°C); p.n. 3; r.a.m. 6.941. |
| Biological Functions | For more than 40 years, Li+ has been used to treat mania.While it is relatively inert in individuals without amood disorder, lithium carbonate is effective in 60 to80% of all acute manic episodes within 5 to 21 days ofbeginning treatment. Because of its delayed onset of actionin the manic patient, Li+ is often used in conjunctionwith low doses of high-potency anxiolytics (e.g., lorazepam)and antipsychotics (e.g. haloperidol) tostabilize the behavior of the patient. Over time, increasedtherapeutic responses to Li+ allow for a gradualreduction in the amount of anxiolytic or neuroleptic required,so that eventually Li+ is the sole agent used tomaintain control of the mood disturbance. In addition to its acute actions, Li+ can reduce thefrequency of manic or depressive episodes in the bipolarpatient and therefore is considered a mood-stabilizingagent. Accordingly, patients with bipolar disorder areoften maintained on low stabilizing doses of Li+ indefinitelyas a prophylaxis to future mood disturbances.Antidepressant medications are required in addition toLi+ for the treatment of breakthrough depression. |
| General Description | A soft silvery metal that is normally grayish white due to oxide formation. Spontaneous ignition is likely if heated to melting point. |
| Air & Water Reactions | Highly flammable. Is readily ignited by and reacts with most extinguishing agents such as water, carbon dioxide, and carbon tetrachloride [Mellor 2, Supp 2:71. 1961]. Reacts with water to form caustic Litium hydroxide and hydrogen gas (H2). Litium is spontaneously flammable in air if heated to 180°C if the surface of the metal is clean. |
| Reactivity Profile | Burns in air, oxygen, nitrogen, hydrogen, and carbon dioxide. The reactions can become extremely violent at higher temperatures. The disposition to ignite of surfaces of molten Litium exposed to any of these gases is increased by the presence of Litium oxides and nitrides. Litium reacts avidly with water to generate gaseous hydrogen and a solution of Litium hydroxide (a caustic). Contact with halogenated hydrocarbons can produce extremely violent reactions, especially on impact [Haz. Chem. Data 1966]. Boron trifluoride reacts with incandescence when heated with Litium [Merck 11th ed. 1989]. Maleic anhydride decomposes explosively in the presence of Litium [Chemical Safety Data Sheet SD-88. 1962, Chem. Haz. Info. Series C-71. 1960]. Chlorine vapors and Litium react producing a luminous flame [Mellor 2, Supp. 1:380. 1956]. The product of the reaction between Litium and carbon monoxide, Litium carbonyl, detonates violently with water, igniting the gaseous products [Mellor 2, Supp. 2:84. 1961]. The reaction of Litium and ferrous sulfide starts around 260°C with subsequent rise in temperature to 950° C [Mellor 2, Supp. 2:80. 1961]. A truck, which was carrying Litium batteries, sodium dithionite and derivatives of cyanide, caught fire; multiple explosions occurred as the cargo was exposed to the air. |
| Hazard | Lithium metal is highly flammable, explosive, and toxic. It will ignite when exposed towater, acids, and even damp air. Metallic lithium is a reducing agent that readily gives up anelectron to active oxidizing agents that require an electron to complete their outer valenceshell—thus the violent chemical reaction that follows. Lithium will even burn in nitrogengas, which is relatively stable. In addition, many of its compounds also react violently whenexposed to water. As an element (metal), it must be stored in oil or in some type of air and moisture-free container,given that many of its compounds will also burn when exposed to air or water. Lithiumfires are difficult to extinguish. If water is poured on the fire, lithium will just burn faster orexplode. A supply of special chemicals or even dry sand is required to extinguish such fires. Solutions and powders of several lithium salts are very toxic to the human nervous system,thus requiring close observation by a physician when used as antidepressant drugs. |
| Health Hazard | Lithium can react with moisture on the skinto produce corrosive hydroxide. Thus, contactof this metal with the skin or eyes can causeburn. The fumes are irritating to the skin,eyes, and mucous membranes. Ingestion of lithium can cause kidney injury, especiallywhen sodium intake is limited (Merck 1996). |
| Fire Hazard | Lithium is less active than sodium or potassium. Finely divided metal ignites spontaneously in air. The ignition of the bulk metal occurs when heated to its melting point. It burns with a carmine-red flame. Burning evolves dense white and opaque fumes. Vigorous reaction occurs when the metal is mixed with water. The heat of reaction, if not dissipated, can ignite or explode hydrogen that is liberated. Violent explosive reactions occur with carbon tetrachloride; carbon tetrabromide; chloroform, bromoform, or iodoform (on heating); carbon monoxide in the presence of water; phosphorus (on heating); arsenic (on heating); and sulfur (molten). Among the substances that constitute high explosion hazards, the halogenated hydrocarbons are most significant. A number of compounds of this class in addition to those mentioned above form impact-sensitive products that can detonate on heating or impact. Heating with nitric acid can cause fires. Lithium reacts with nitrogen at elevated temperatures to form lithium nitride, which can ignite on heating. |
| Flammability and Explosibility | Contact with water liberates highly flammable gases |
| reaction suitability | reagent type: reductant |
| Pharmaceutical Applications | Lithium inhibits GSK-3 and InsP, and both pathways have therefore been suggested to be involved in thetreatment of BD and schizophrenia. The theory behind this hypothesis is that overactive InsP signalling in thebrain of these patients potentially causes BD and this may be reduced by the inhibitory effect of lithium onsuch signalling. It is believed that lithium potentially can protect against disease-induced cell death. GSK-3 has been implicated in the origins of schizophrenia, butwith the availability of many antipsychotic drugs on the market, lithium ions are not in common use for thetreatment of schizophrenia.There are also several direct roles of lithium in the treatment of Alzheimer’s disease. Alzheimer’s diseaseis a neurodegenerative brain disorder causing neuronal dysfunction and ultimately cell death. Onset occurs with the accumulation of extracellular senile plaques composed of amyloid-β peptides and withthe accumulation of intercellular neurofibrillary tangles. |
| Industrial uses | This lightest of all metals is found in more than 40 minerals, but is obtained chiefly from lepidolite, spodumene, and salt brines. It is unstable chemically and burns in the air with a dazzling white flame when heated to just above its melting point. The metal is silvery white but tarnishes quickly in the air. The metal is kept submerged in kerosene. Lithium resembles sodium, barium, and potassium, but has a wider reactive power than the other alkali metals. It combines easily with oxygen, nitrogen, and sulfur to form low melting-point compounds that pass off as gases, and is thus useful as a deoxidizer and degasifier of metals. |
| Mechanism of action | Lithium is a monovalent cation that can replace Na+ insome biological processes. It can be argued that competitionby Li+ for active Na+ sites may lead to alteredneuronal functions that may account for its antimanicand mood-stabilizing actions. In this regard, the failureof Li+ to maintain a normal membrane potential becauseof its lower affinity for the Na+ pump has beendemonstrated. However, this action of Li+ would notexplain its relatively selective effects on the CNS, sparingcomparable excitable tissues (e.g. cardiac muscle) inthe periphery.Moreover, an action on membrane polaritywould be so general that the entire pool of brainneurons would be affected by Li+. It seems more reasonablethat Li+ produces its psychotropic actions byperturbation of molecular events common to a few CNSsynapses that might have been disturbed during thecourse of the manic-depressive illness. Recently, attention has focused on the actions of Li+on receptor-mediated second-messenger signaling systemsof the brain. In this regard, interactions betweenLi+ and guanine nucleotide (GTP) binding proteins (Gproteins) have been the target of many studies, since Gproteins play a pivotal role in the function of many second-messenger signaling systems. Lithium is capable ofaltering G-protein function. It can diminish the couplingbetween the receptor recognition site and the Gprotein.The molecular mechanism involves the competitionfor Mg++ sites on the G protein, which are essentialfor GTP binding. Guanine nucleotide activates theG protein. Accordingly, in the presence of Li+, receptormediatedactivation of these G proteins is attenuated.This action of Li+ has been selectively demonstrated forG proteins associated with β-adrenoceptors and M1muscarinic receptors of the CNS. While it is not possible at present to assign a therapeuticrole to this action of Li+, it is a step toward explainingthe stabilizing actions of this drug. Since severalneurotransmitter receptors share common Gprotein–regulated second-messenger signaling systems,Li+ could simultaneously correct the alterations at individualsynapses associated with depression and maniaby a single action on the function of specific G proteins.An   |