Thallium is a chemical element; it has symbol Tl and atomic number 81. It is a silvery-white post-transition metal that is not found free in nature. When isolated, thallium resembles tin, but discolors when exposed to air. Chemists William Crookes and Claude-Auguste Lamy discovered thallium independently in 1861, in residues of sulfuric acid production. Both used the newly developed method of flame spectroscopy, in which thallium produces a notable green spectral line. Thallium, from Greek , , meaning "green shoot" or "twig", was named by Crookes. It was isolated by both Lamy and Crookes in 1862, Lamy by electrolysis and Crookes by precipitation and melting of the resultant powder. Crookes exhibited it as a powder precipitated by zinc at the International Exhibition, which opened on 1 May that year.The Mining and Smelting Magazine . Ed. Henry Curwen Salmon. Vol. iv, July–Dec 1963, p. 87.
Thallium tends to form the +3 and +1 oxidation states. The +3 state resembles that of the other elements in group 13 (boron, aluminium, gallium, indium). However, the +1 state, which is far more prominent in thallium than the elements above it in the periodic table, recalls the chemistry of alkali metals. Thallium(I) ions are found geologically mostly in potassium-based ores and (when ingested) are handled in many ways like potassium ions () by ion pumps in living cells.
Commercially, thallium is produced not from potassium ores, but as a byproduct from refining of heavy-metal sulfide ores. Approximately 65% of thallium production is used in the electronics industry and the remainder is used in the pharmaceutical industry and in glass manufacturing. It is also used in infrared detectors. The radioisotope thallium-201 (as the soluble chloride TlCl) is used in small amounts as an agent in a nuclear medicine scan, during one type of nuclear cardiac stress test.
Soluble thallium salts (many of which are nearly tasteless) are highly toxic and they were historically used in rat poisons and insecticides. Because of their nonselective toxicity, use of these compounds has been restricted or banned in many countries. Thallium poisoning usually results in hair loss. Because of its historic popularity as a murder weapon, thallium has gained notoriety as "the poisoner's poison" and "inheritance powder" (alongside arsenic).
Characteristics A thallium atom has 81 electrons, arranged in the electron configuration [Xe]4f145d106s26p1; of these, the three outermost electrons in the sixth shell are valence electrons. Due to the inert pair effect, the 6s electron pair is relativistically stabilised and it is more difficult to get these involved in chemical bonding than it is for the heavier elements. Thus, very few electrons are available for metallic bonding, similar to the neighboring elements mercury and lead. Thallium, then, like its congeners, is a soft, highly electrically conducting metal with a low melting point, of 304 °C.Greenwood and Earnshaw, pp. 222–224
Thallium is malleable and sectile enough to be cut with a knife at room temperature. It has a metallic luster that, when exposed to air, quickly tarnishes to a bluish-gray tinge, resembling lead. It may be preserved by immersion in oil. A heavy layer of oxide builds up on thallium if left in air. The metal also reacts with water, forming thallium hydroxide. Sulfuric, nitric, and hydrofluoric acid dissolve thallium rapidly to make the corresponding thallium(I) salts, but the other hydrohalic acids form a passivating layer of thallium(I) halide.
A number of standard electrode potentials, depending on the reaction under study, are reported for thallium, reflecting the greatly decreased stability of the +3 oxidation state: +0.73 V -0.336 V Thallium is the first element in group 13 where the reduction of the +3 oxidation state to the +1 oxidation state is spontaneous under standard conditions. Since bond energies decrease down the group, with thallium, the energy released in forming two additional bonds and attaining the +3 state is not always enough to outweigh the energy needed to involve the 6s-electrons.Greenwood and Earnshaw, pp. 224–7
Thallium(I) oxide and hydroxide are more basic and thallium(III) oxide and hydroxide are more acidic, showing that thallium conforms to the general rule of elements being more electropositive in their lower oxidation states.
Isotopes
Thallium has 41 known isotopes with atomic masses from 176 to 216. Tl and Tl are the only stable isotopes and make up all natural thallium. The five short-lived isotopes Tl through Tl inclusive occur in nature, but only as part of the natural decay chains of heavier elements. Tl is the most stable radioisotope, with a half-life of 3.78 years; the next most stable are Tl (half-life 12.31 days) and Tl (half-life 3.0421 days). It is made by the neutron activation of stable thallium in a nuclear reactor.
The isotope Tl is useful in nuclear medicine; it decays by electron capture, emitting X-rays (~70–80 keV), and gamma rays of 135 and 167 keV; therefore, it has good imaging characteristics without an excessive patient-radiation dose. It is the most popular isotope used for thallium nuclear cardiac stress tests.
Compounds
Thallium(III) Thallium(III) compounds resemble the corresponding aluminium(III) compounds. They are moderately strong oxidizing agents and are usually unstable, as illustrated by the positive reduction potential for the /Tl couple. Some mixed-valence compounds are also known, such as and , which contain both thallium(I) and thallium(III).
The simplest possible thallium compound, thallane (), is too unstable to exist in bulk, both due to the instability of the +3 oxidation state as well as poor overlap of the valence 6s and 6p orbitals of thallium with the 1s orbital of hydrogen. The trihalides are more stable, although they are chemically distinct from those of the lighter group 13 elements and are still the least stable in the whole group. For instance, thallium(III) fluoride, , but not the lighter group 13 trifluorides, has the structure and does not form the complex anion in aqueous solution. The trichloride and tribromide disproportionate just above room temperature to give the monohalides. Thallium triiodide is actually a thallium(I) compound with the linear triiodide anion ()Greenwood and Earnshaw, p. 239 in neutral solution, but forms a stable [TlI4]− anion; similar behavior occurs for the common pseudohalides.
Thallium(III) oxide, , is a black solid which decomposes above 800 °C, forming the thallium(I) oxide and oxygen. The material is amphoteric, and forms a number of ternary (complex) oxides. Heavier thallium(III) sesquichalcogenides do not exist.Greenwood and Earnshaw, p. 254 As thallic ion is relatively oxidizing, oxyacid salts are generally with the most oxidized form of the central atom: thus a phosphate, a nitrate, a sulfate, and many carboxylates. All are moisture-sensitive. The carboxylates are much stronger electrophiles than the trihalides.
Thallium(III) complexes are generally tetrahedral or octacoordinate pseudotetrahedral.
Certain electropositive and coordinating anions prefer the thallic state. Thus thallous tetracarbonylcobaltate can only be made through carbonylating a mixture of the metals; Tl+ ions react with NaCo(CO)4 to give Tl[Co(CO)4]3 instead. Similar behavior occurs with Mn(CO), GaH, and BH; and is the rule in organothallium chemistry.
Thallium(I) Thallium(I) chemistry is quite different than the rest of the boron group and somewhat resembles that of the alkali metals. It forms salts with virtually all the classical anions of inorganic chemistry, both monatoamic and polyatomic. Thallium halide and pseudohalide salts generally have little covalent character, with thallous cyanamide being a notable exception. However, bonds between thallium and the rest of the main group are considerably more covalent than their alkali analogues.
The simplest salts are the thallium monohalides. In keeping with the large size of the cation, the chloride and bromide have the caesium chloride structure, while the fluoride and iodide have distorted sodium chloride structures. Like the analogous silver compounds, TlCl, TlBr, and TlI are photosensitive and display poor solubility in water.Greenwood and Earnshaw, p. 241
A stable oxide and hydroxide are known, as are many chalcogenides.Greenwood and Earnshaw, pp. 246–247 The metalorganic compound thallium ethoxide (TlOEt, ) is a heavy liquid (ρ , m.p. −3 °C), often used as a basic and soluble thallium source in organic and organometallic chemistry. It and the other alkoxides are however somewhat unstable at room temperature.
Thallium(I) forms ternary oxides and sulfides with most elements. Thallous sulfate and nitrate are known, as are salts of virtually all the lower nitrogenous and sulfurous anions. Complex salts with transition metal ions sometimes exhibit electrical conductivity or spectral shifts from perturbative oxidation to Tl2+.
Thallous nitride is formed by metathesis between thallium nitrate and potassamide: TlNO3 + KNH2 → Tl3N + KNO3 + NH3 The product reacts explosively with water. In dry conditions, it can be used to substitute Tl for H on e.g. amides and imides. Absent an additional Lewis base functionality, these salts are generally subsceptible to hydrolysis.
Within the adamantogens, thallium(I) forms a carbonate and a bicarbonate.Greenwood and Earnshaw, pp. 246–247 The double salt has been shown to have hydroxyl-centred triangles of thallium, , as a recurring motif throughout its solid structure. Carboxylic acids react with thallous hydroxide to form thallium salts (thallium soaps), but reducing sugars generally deposit thallium metal.
Organothallium compounds
Organothallium compounds tend to be thermally unstable, in concordance with the trend of decreasing thermal stability down group 13. The chemical reactivity of the Tl–C bond is also the lowest in the group, especially for ionic compounds of the type . Thallium forms the stable ion in aqueous solution; like the isoelectronic and , it is linear. Trimethylthallium and triethylthallium are, like the corresponding gallium and indium compounds, flammable liquids with low melting points. Like indium, thallium cyclopentadienyl compounds contain thallium(I), in contrast to gallium(III).Greenwood and Earnshaw, pp. 262–264
Exotica Tl(+2) has been observed in tetrakis(hypersilyl)dithallium ([(Me3Si)Si]2Tl—Tl[Si(SiMe3)]2), produced from hypersilylrubidium and thallium hexamethyldisilazane in toluene at low temperature. Using thallous halides instead of the amide gives disproportionated products. The analogous supersilyl compound is believed to dissociate to Tl(II) radicals at room temperature.
History Thallium (Greek , , meaning "a green shoot or twig")Liddell, Henry George and Scott, Robert (eds.) "θαλλος ", in A Greek–English Lexicon, Oxford University Press. was discovered by William Crookes and Claude Auguste Lamy, working independently, both using flame spectroscopy (Crookes was first to publish his findings, on March 30, 1861).* Crookes, William (March 30, 1861) "On the existence of a new element, probably of the sulphur group," Chemical News, vol. 3, pp. 193–194; reprinted in: Crookes, William (May 18, 1861) "Further remarks on the supposed new metalloid," Chemical News, vol. 3, p. 303. Crookes, William (June 19, 1862) "Preliminary researches on thallium," Proceedings of the Royal Society of London, vol. 12, pp. 150–159. Lamy, A. (May 16, 1862) "De l'existencè d'un nouveau métal, le thallium," Comptes Rendus, vol. 54, pp. 1255–1262. . The name comes from thallium's bright green spectral emission lines.