Lithium

Educational institution

Summary

Subject: Lithium

He performed the: .

Checked: .

City: 20 г.

Content

1 Introduction

2 Characteristics

2.1 Physical

2.2 Chemical

2.3 Lithium compounds

2.4 Isotopes

3 History and etymology

4 Occurrence

5 Production

6 Applications

6.1 Medical use

6.2 Other uses

7 Precautions

7.1 Regulation

8 Conclusion

9 References

1. Introduction

Lithium (pronounced /?l?иi?m/, LITH-ee-?m) is a soft, silver-white metal that belongs to the alkali metal group of chemical elements. It is represented by the symbol Li, and it has the atomic number three. 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 petroleum. When cut open, lithium exhibits a metallic luster, but contact with oxygen quickly turns it back to a dull silvery gray color. Lithium in its elemental state is highly flammable.

According to theory, lithium was one of the few elements synthesized in the Big Bang. Since its current estimated abundance in the universe is vastly less than that predicted by physical theories, the processes by which new lithium is created and destroyed, and the true value of its abundance,[1] continue to be active matters of study in astronomy.[2][3][4] The nuclei of lithium are relatively fragile: the two stable lithium isotopes found in nature have lower binding energies per nucleon than any other stable compound nuclides, save deuterium, and helium-3 (3He).[5] Though very light in atomic weight, lithium is less common in the solar system than 25 of the first 32 chemical elements.[6]

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 lithium ion Li+ administered as any of several lithium salts has proved to be useful as a mood stabilizing drug due to neurological effects of the ion in the human body.[7] 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 transmutation of lithium atoms to tritium was the first man-made form of a nuclear fusion reaction, and lithium deuteride serves as a fusion fuel in staged thermonuclear weapons.

Figure. 0. Silvery white (seen here in oil)

2. Characteristics

2.1 Physical

Like the other alkali metals, lithium has a single valence electron that is easily given up to form a cation.[8] Because of this, it is a good conductor of both heat and electricity and highly reactive, though it is the least reactive of the alkali metals due to the proximity of its valence electron to its nucleus.[8]

Lithium is soft enough to be cut with a knife, and it is the lightest of the metals of the periodic table. When cut, it possesses a silvery-white color that quickly changes to gray due to oxidation.[8] It also has a low density (approximately 0.534 g/cm3) and thus will float on water, with which it reacts easily. This reaction is energetic, forming hydrogen gas and lithium hydroxide in aqueous solution.[8] Due to its reactivity with water, lithium is usually stored in mineral oil or kerosene.[8]

Lithium possesses a low coefficient of thermal expansion and the highest specific heat capacity of any solid element. Lithium is superconductive below 400 мK at standard pressure[9] and at higher temperatures (more than 9 kelvin) at very high pressures (over 200,000 atmospheres)[10] At cryogenic temperatures, lithium, like sodium, undergoes diffusionless phase change transformations. At 4.2K it has a rhombohedral crystal system (with a nine-layer repeat spacing)[11]; 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.

Figure. 1. Lithium pellets (covered in white lithium hydroxide)

2.2 Chemical

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.[13]

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

2.3 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]

2.4 Isotopes

Naturally occurring lithium is composed of two stable isotopes, 6Li and 7Li, the latter being the more abundant (92.5% natural abundance).[8][14] 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.[15] Additional small amounts of lithium of both 6Li and 7Li may be generated from solar wind, cosmic rays, and early solar system 7Be and 10Be radioactive decay.[16] 7Li can also be generated in carbon stars.[17]

Lithium isotopes fractionate substantially during a wide variety of natural processes,[18] 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.

3. History and etymology

Petalite (LiAlSi4O10, which is lithium aluminium silicate) was first discovered in 1800 by the Brazilian chemist Josй Bonifбcio de Andrade e Silva, who discovered this mineral in a mine on the island of Utц, Sweden.[19][20][21] However, it was not until 1817 that Johan August Arfwedson, then working in the laboratory of the chemist Jцns Jakob Berzelius, detected the presence of a new element while analyzing petalite ore.[22][23][24] This element formed compounds similar to those of sodium and potassium, though its carbonate and hydroxide were less soluble in water and more alkaline.[25] Berzelius gave the alkaline material the name "lithos", from the Greek word лйиoт (transliterated as lithos, meaning "stone"), to reflect its discovery in a solid mineral, as opposed to sodium and potassium, which had been discovered in plant tissues. The name of this element was later standardized as "lithium".[8][20][24] Arfwedson later showed that this same element was present in the minerals spodumene and lepidolite.[20] In 1818, Christian Gmelin was the first man to observe that lithium salts give a bright red color in flame.[20] However, both Arfwedson and Gmelin tried and failed to isolate the element from its salts.[20][24][26] This element, lithium, was not isolated until 1821, when William Thomas Brande isolated the element by performing electrolysis on lithium oxide, a process that had previously employed by the chemist Sir Humphry Davy to isolate the alkali metals potassium and sodium.[26][27][28] Brande also described some pure salts of lithium, such as the chloride, and he performed an estimate of its atomic weight. In 1855, larger quantities of lithium were produced through the electrolysis of lithium chloride by Robert Bunsen and Augustus Matthiessen.[20] The discovery of this procedure henceforth led to commercial production of lithium metal, beginning in 1923 by the German company Metallgesellschaft AG, which performed an electrolysis of a liquid mixture of lithium chloride and potassium chloride.[20][29]

The production and use of lithium underwent several drastic changes in history. The first major application of lithium became high temperature grease for aircraft engines or similar applications in World War II and shortly after. This small market was supported by several small mining operations mostly in the United States. The demand for lithium increased dramatically when in the beginning of the cold war the need for the production of nuclear fusion weapons arose and the dominant fusion material tritium had to be made by irradiating lithium-6. The United States became the prime producer of lithium in the period between the late 1950s and the mid 1980s. At the end the stockpile of lithium was roughly 42.000 tons of lithium hydroxide. The stockpiled lithium was depleted in lithium-6 by 75% .[30]

Lithium was used to decrease the melting temperature of glass and to improve the melting behavior of aluminium chloride when using the Hall-Hйroult process.[31][31] These two uses dominated the market until the middle of the 1990's. After the end of the nuclear arms race the demand for lithium decreased and the sale of Department of Energy stockpiles on the open market further reduced prices.[30] Then, in the mid 1990's several companies started to extract lithium from brine; this method proved to be less expensive than underground or even open pit mining. Most of the mines closed or shifted their focus to other materials as only the ore from zoned pegmatites could be mined for a competitive price. For example, the US mines near Kings Mountain, North Carolina closed before the turn of the century. The use in lithium ion batteries increased the demand for lithium and became the dominant use in 2007.[30] New companies have expanded brine extraction efforts to meet the rising demand.[32]

4. Occurrence

According to theory, the stable isotopes 6Li and 7Li were created in the Big Bang, but the amounts are unclear. Lithium is a fusion fuel in main sequence stars. 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 between carbon (element 6) and 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. In turn the destruction of lithium isotopes is due to their very low energy of binding per nucleon with regard to all other nuclides save deuterium (also destroyed in stars) and helium-3.[5] This low energy of binding encourages breakup of lithium in favor of more tightly-bound nuclides under thermonuclear reaction conditions.

Lithium is widely distributed on Earth but does not naturally occur in elemental form due to its high reactivity.[8] 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 the Western Lithium Corporation in the United States.[34]

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."[35] At 20 mg lithium per kg of Earth's crust [36], lithium is the 25th most abundant element. Nickel and lead have the about the same abundance.

Страницы: 1, 2