Thallium

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 θαλλός, thallus, 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.
Thallium tends to form the +3 and +1 oxidation states. The +3 state resembles that of the other elements in group 13. However, the +1 state, which is far more prominent in thallium than the elements above it, recalls the chemistry of alkali metals and thallium ions are found geologically mostly in potassium-based ores and 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 is used in small amounts as an agent in a nuclear medicine scan, during one type of nuclear cardiac stress test.
Soluble thallium salts 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".

Characteristics

A thallium atom has 81 electrons, arranged in the electron configuration 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.
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:
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. Accordingly, thallium oxide and hydroxide are more basic and thallium oxide and hydroxide are more acidic, showing that thallium conforms to the general rule of elements being more electropositive in their lower oxidation states.
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 and nitric acids dissolve thallium rapidly to make the sulfate and nitrate salts, while hydrochloric acid forms an insoluble thallium chloride layer.

Isotopes

Thallium has 41 known isotopes with atomic masses from 176 to 216. ²⁰³Tl and ²⁰⁵Tl are the only stable isotopes and make up all natural thallium. The five short-lived isotopes ²⁰⁶Tl through ²¹⁰Tl inclusive occur in nature, but only as part of the natural decay chains of heavier elements. ²⁰⁴Tl is the most stable radioisotope, with a half-life of 3.78 years; the next most stable are ²⁰²Tl and ²⁰¹Tl. It is made by the neutron activation of stable thallium in a nuclear reactor.
The isotope ²⁰¹Tl is useful in nuclear medicine; it decays by electron capture, emitting X-rays, 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 compounds resemble the corresponding aluminium compounds. They are moderately strong oxidizing agents and are usually unstable, as illustrated by the positive reduction potential for the Tl³⁺/Tl couple. Some mixed-valence compounds are also known, such as Tl₄O₃ and TlCl₂, which contain both thallium and thallium. Thallium oxide, Tl₂O₃, is a black solid which decomposes above 800 °C, forming the thallium oxide and oxygen.
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 fluoride, TlF₃, has the β-BiF₃ structure rather than that of the lighter group 13 trifluorides, and does not form the complex anion in aqueous solution. The trichloride and tribromide disproportionate just above room temperature to give the monohalides, and thallium triiodide contains the linear triiodide anion and is actually a thallium compound. Thallium sesquichalcogenides do not exist.

Thallium(I)

The thallium halides are stable. In keeping with the large size of the Tl⁺ 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. The stability of thallium compounds demonstrates its differences from the rest of the group: a stable oxide, hydroxide, and carbonate are known, as are many chalcogenides.
The double salt has been shown to have hydroxyl-centred triangles of thallium,, as a recurring motif throughout its solid structure.
The metalorganic compound thallium ethoxide is a heavy liquid, often used as a basic and soluble thallium source in organic and organometallic chemistry.

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 R₂TlX. Thallium forms the stable ⁺ ion in aqueous solution; like the isoelectronic Hg₂ and ²⁺, 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, in contrast to gallium.

History

Thallium was discovered by William Crookes and Claude Auguste Lamy, working independently, both using flame spectroscopy. The name comes from thallium's bright green spectral emission lines derived from the Greek 'thallos', meaning a green twig.
After the publication of the improved method of flame spectroscopy by Robert Bunsen and Gustav Kirchhoff and the discovery of caesium and rubidium in the years 1859 to 1860, flame spectroscopy became an approved method to determine the composition of minerals and chemical products. Crookes and Lamy both started to use the new method. Crookes used it to make spectroscopic determinations for tellurium on selenium compounds deposited in the lead chamber of a sulfuric acid production plant near Tilkerode in the Harz mountains. He had obtained the samples for his research on selenium cyanide from August Hofmann years earlier. By 1862, Crookes was able to isolate small quantities of the new element and determine the properties of a few compounds. Claude-Auguste Lamy used a spectrometer that was similar to Crookes' to determine the composition of a selenium-containing substance which was deposited during the production of sulfuric acid from pyrite. He also noticed the new green line in the spectra and concluded that a new element was present. Lamy had received this material from the sulfuric acid plant of his friend Frédéric Kuhlmann and this by-product was available in large quantities. Lamy started to isolate the new element from that source. The fact that Lamy was able to work ample quantities of thallium enabled him to determine the properties of several compounds and in addition he prepared a small ingot of metallic thallium which he prepared by remelting thallium he had obtained by electrolysis of thallium salts.
As both scientists discovered thallium independently and a large part of the work, especially the isolation of the metallic thallium was done by Lamy, Crookes tried to secure his own priority on the work. Lamy was awarded a medal at the International Exhibition in London 1862: For the discovery of a new and abundant source of thallium and after heavy protest Crookes also received a medal: thallium, for the discovery of the new element. The controversy between both scientists continued through 1862 and 1863. Most of the discussion ended after Crookes was elected Fellow of the Royal Society in June 1863.
The dominant use of thallium was the use as poison for rodents. After several accidents the use as poison was banned in the United States by Presidential Executive Order 11643 in February 1972. In subsequent years several other countries also banned its use.