Lithium and its Compounds

Lithium and its Compounds

1. Discovery and naming of Lithium

Lithium was discovered in 1817 by J.A. Arfvedson at Stockholm Sweden. Isolated in 1821 by W.T. Brande

Source: The Elements(John Emsley) Page 116; (WSU Reference QD 466.e48 1999)

The lithium containing materials petalite and spodumene, were discovered by Jose de Andrada between 1790 and 1800 on Utö island in Sweden. In 1817 J.A. Arfvedson discovered the new alkali metal lithium in petalite. He noted that lithium carbonate is sparingly soluble , that the hydroxide is much less soluble than the hydrogen oxides of other alkali metals and that lithium compounds are similar to those of other alkaline earth metals.

Source: Comprehensive Inorganic Chemistry (J.C. Bailar et al.) Page 331 (WSU Library QD 451.2 C64 V.1)

The Name Lithium was given to the new element in recognition of the fact that lithium was recovered first from a mineral whereas both sodium and potassium were first derived from plant matter.

Source: Comprehensive Inorganic Chemistry (J.C. Bailar et al.) Page 331 (WSU Library QD 451.2 C64 V.1)

[from the journal of J.A. Arfvedson] It’s solution could not be precipitated either by tartaric acid in excess or by platinum chloride. Consequently it could not be potassium. I mixed another portion of the solution with a few drops of pure potash but without it’s becoming cloudy. Therefore it contained no magnesia. … At last having studied more closely the sulfate in question I soon found that it contained a definite fixed alkali whose nature had not been previously known.

Source: The chemical elements (H.M. Davis and G. Seaborg) page 22 (WSU Library QD 466.d35 1961)

2. Occurrence and extraction of Lithium

The abundance of lithium in the earth’s crust is estimated to be about 0.005 percent. That places it among the top 15 elements found in the earth. The most common ores of lithium are spodumme, petalite, and lepidolite. Lithium is also obtained from salt water. … The worlds largest producer of lithium is the United States

Source: Chemical Elements (David e.Newton) Pages 309-310 (WSU Reference QD 466.N464 1999V.2)

Melt [lithium chloride] and then insert the electrodes so that they dip well into the solution. Shortly after you have immersed the electrodes a silver substance will form on the cathode. This is the metal lithium.

Source: Chemistry magic (Sweezey) page 72 (WSU Library QD 43 S8)

3. Uses of Lithium

The first commercial use of lithium occurred towards the end of World War I when small amounts were used in an aluminum zinc allow, Scleron. After the war lithium was used as a hardener in a lead bearing alloy-bearing material. Between world Wars I and II there was little production of lithium materials. Most lithium trade was as ores sold as additives of frit and glass formulators. ... LiH was used in military sea rescue equipment as a source of hydrogen upon reaction with water. the hydrogen inflated rescue balloons to carry radio antennas needed for the SOS signal broadcast. During this period all temperature greases using lithium stearate were produced for military applications.

Source: Encyclopedia of chemical technology fifth Edition volume 15 pages 120 to 121 (WSU Reference TP9.K54 2004 V.15)

4. General Properties of Lithium

The group has been much studied as it best illustrated the effects of increasing atomic

size and mass on chemical and physical properties.

Source: Encyclopedia of inorganic chemistry (R. Bruce King) Pages 35-36 (WSU Library QD 148.E53 1994 V.1)

Lithium is a very soft silvery metal. It has a melting point of 180.54°C and a boiling point of about 1335°C. Its density is 0.534 grams per cubic centimeter. … Lithium’s hardness is 0.6 on the Mohs scale. Lithium is an active element but not as active as the other alkali metals. It reacts slowly with water at room temperature. Lithium does not react with oxygen at room temperature. Under proper conditions lithium also combines with sulfur, hydrogen, nitrogen, and the halogens.

Source: Chemical Elements (David E. Newton) Page 308 (WSU Reference QD 466.N464 1999V.2)

One of the most interesting properties [of lithium] is the extremely low density 0.534 gm/cm3. At normal temperatures lithium has the lowest density of a non-gaseous element. The metal is harder and has a higher melting point than the other alkali metals.

Above -117 °C lithium metal has a body centered cubic crystal structure typical of the alkali metals. On cooling to -201°C the metal begins to transform to the face centered cubic structure.

Source: Comprehensive Inorganic Chemistry (J.C. Bailar et al.) Page 337(WSU Library QD 451.2 C64 V.1)

Just as the properties of lithium set it apart [physically] from the other alkali metals these same properties cause the compounds of lithium to differ from the other alkali metal compounds. ... Several general statements may be useful in the understanding of lithium compounds. The lattice energies of ionic lithium compounds are greater than the lattice energies of the corresponding salts of the other alkali metals

Source: The Elements(John Emsley) Page 116; (WSU Reference QD 466.e48 1999)

5. Aqueous chemistry of Lithium

There are numerous crystal structures of the Li+ ion with crown ethers. .. In fact most of these crystal structures are misleading in that the Li+ ion forms no complexes, or at best very weak complexes, with the familiar crown ethers such as 12-crown-4 in any known solvent

Source: Metal Complexes in Aqueous Solutions (A E Martell, R D Hancock) introduction page v (WSU library QD474.M375 1996)

6. Lithium-oxygen compounds

Lithium superoxide is a major product of the co-condensation reaction of an atomic beam of lithium with a jet of molecular oxygen at 4.2-15K.

Source: Inorganic Chemistry of the main Group Elements-volume 1(CC Addison et al) page 14 (WSU Library QD 146.I54 V.1)

The compound Li4XeO6·2H2O is Amorphous to X-rays, unlike the analogous sodium and potassium compounds, and is less stable but shows the absorption bands in the Ir as 640-720 and 2900-3600 cm-1 which are due to the XeO6- ion and H2O respectively.

Source: Inorganic Chemistry of the main Group Elements-volume 1(CC Addison et al) page 16 (WSU Library QD 146.I54 V.1)

7. Lithium-hydrogen compounds

Lithium hydride is prepared both industrially and in the laboratory by the reaction of molten lithium metal with hydrogen gas. Industrially lithium metal is heated to initiate its reaction with hydrogen gas.

Lithium hydride is also used commercially in the preparation of lithium aluminum hydride LiAlH₄ and to prepare an intermediate used in the synthesis of vitamin A.

LiH is a very strong base. It is insoluble in solvents except those where it reacts to form a soluble compound.

Source: Comprehensive Inorganic Chemistry (J.C. Bailar et al.) Page 331 (WSU Library QD 451.2 C64 V.1)

8. Lithium in the synthesis of organic compounds

The discovery by F.W. Stavely that metallic lithium and many organo-lithium compounds initiate the polymerization of isoprene to for an elastomer of nearly identical properties to those of natural rubber began a new era in rubber technology.

There is a broad range of polymerizations which can be initiated by lithium or it’s compounds. The most commonly used catalyst is butyl-lithium. The quantity of Li in the final product is of course minute.

Source: Specialty inorganic chemicals (R. Thompson) pages 112-113 (WSU library QD 151.2.S664)

Organolithium compounds are central to many aspects of synthetic organic chemistry and are primarily used as carbanions to construct carbon skeletons of a wide variety of organic compounds. Tremendous efforts have been devoted to the development of convenient methods of for generation of tailor made organolithium compounds.

Source: Main group metals in organic synthesis (H. Yamamoto, K Oshima) page 1 (WSU Library QD 411.7.S94 M33 2004 v.1)

Although [organo-lithium compounds] do not contain free carbanions they act as if they did. They are sources of R:^-. The very polar carbon-metal bond attacks all manner of lewis and bronstead acids. Water is more than sufficiently stron enough to protonate a Grignard or organo-lithium reagent.

The end product is a hydrocarbon. The sequence constitutes a new synthesis of hydrocarbons from alkyl, alkenyl, or aryl halides. … This reaction can be put to good use in the construction of isotypically labeled reagents because D₂O can be used in place of water to produce specifically deuteriolabled hydrocarbons.

CH₃CH₂Br + Li CH₃CH₂Li+D₂OCH₃CH₂D

Source: Organic Chemistry second edition (M. Jones) pages 734-735 (Stoker Library)

9. Lithium electrochemical cells

Lithium is the most electropositive of the metals with a standard electrode potential of 3.045v… Consequently Li metal cells have the potential to store the largest amount of power/unit weight or volume. … The greatest potential of course lies in the development of the electric vehicle and off-peak power storage.

Source: Specialty inorganic chemicals (R. Thompson) page 117 (WSU library QD 151.2.S664)

The lithium rechargeable battery used in portable computers and phones uses LiCoO₂ as the anode with a graphite cathode. Lithium ions are produced at the cathode and to maintain charge balance Co(III) is oxidized to Co(IV) at the anode.

LiCoO₂Li_1-xCoO₂+xLi⁺(solvent)+xe⁺

The lithium ions are intercalated into the graphite and return to the LiCoO₂ when the cell is discharged. The battery is rechargeable because for the cathode and the anode can act as hosts for the Li⁺ ions which can move back and forth between them.

Source: Inorganic Chemistry (Shriver & Atkins) page 261 (Stoker Library)

10. Toxicity of Lithium

Lithium is moderately toxic by ingestion but there are wide variations of tolerance. Even Lithium Carbonate which is used in psychiatry is prescribed at doses near to the toxic level. Some Lithium compounds are carcinogenic and teratogenic.

Source: The Elements (John Emsley) Page 116; (WSU Reference QD 466.e48 1999)

Lithium oxides, hydroxide, and carbonate are strong bases and their water solutions are very caustic. The toxicity of lithium compounds is a function of their solubility in water. Lithium ion is toxic to the central nervous system. Lithium is commonly ingested at dosages of 0.5g/d of lithium carbonate for treatment of bipolar disorders. Therapeutic vs. Toxic levels are small: 2mmol/L is toxic 4mmol/L is fatal. LiOH either directly or formed by the hydrolysis of other salts can cause caustic burns. Skin contact with lithium halides can result in skin dehydration. Organo-lithium compounds are often pyrophoric and require special handling.

Source: Encyclopedia of chemical technology fifth Edition volume 15 pages 134 to 135 (WSU Reference TP9.K54 2004 V.15)