Thallium (symbol Tl, atomic number 81) is a heavy, soft post-transition (poor) metal in group 13 of the periodic table. In its pure form it is a silvery-gray metal that tarnishes quickly in air. At standard conditions it is a solid (melting point ≈ 303 °C, boiling point ≈ 1457 °C) with a density around 11.8 g/cm³ (about 11.85 g/cm³). Thallium most commonly exhibits oxidation state +1 (thallium(I)), with +3 (thallium(III)) also known but much less stable. Like its lighter congeners (aluminium, gallium, indium), thallium’s chemistry is dominated by its valence ℓ = 6s²6p¹ electron arrangement Xe]4f¹⁴5d¹⁰6s²6p¹), but the so-called “inert pair effect” makes the +1 state unusually favored. In compounds Tl(I) behaves somewhat like alkali or coinage metals: thallium(I) ions are large (ionic radius ≈ 150 pm) and form basic oxides and hydroxides. In contrast, Tl(III) compounds (where thallium loses the 6s electrons) are strong oxidizers and resemble aluminium(III) or gallium(III) chemistry. Naturally occurring thallium is never found free; it is always combined, typically in sulfide or selenide minerals, because metallic Tl is easily oxidized.

Key Facts:.

Symbol: Tl; Atomic Number: 81
Group/Period: 13 (IIIA), Period 6, p-block element (post-transition metal)
Electron Configuration: [Xe] 4f¹⁴ 5d¹⁰ 6s² 6p¹ (valence shell 6s² 6p¹)
Common Oxidation States: +1 (dominant), +3 (stable under strong oxidizing conditions)
Phase at STP: Solid (soft, malleable metal)
Appearance: Silvery-gray metal; rapidly forms a bluish–gray tarnish (oxide layer) in air.

Atomic Structure and Trends

Thallium
Atomic number	81
Symbol	Tl
Group	13 (boron group)
Boiling point	1473 °C
Electronegativity	1.62 (Pauling)
Electron configuration	[Xe] 4f14 5d10 6s2 6p1
Melting point	304 °C
Period	6
Phase STP	Solid
Block	p
Oxidation states	+1, +3
Wikidata	Q932

Thallium’s atomic number (81) puts it below indium (49) and above lead (82) in the periodic table. It has 81 electrons; the two 6s electrons and one 6p electron form the valence shell. The filled 5d¹⁰ and 4f¹⁴ subshells lie below these and are relatively inert (core-like). As a heavy element, thallium experiences relativistic effects that influence its chemistry. One manifestation is the so-called inert pair effect: the 6s electrons are held more tightly and participate less in bonding, making the +1 state more stable than +3.

In periodic trends, thallium is a very large atom. Depending on the measure, its atomic radius is on the order of ~170–190 pm (empirical “atomic radius”), with a single-bond covalent radius ~148 pm. For comparison, indium’s atomic radius is about 167 pm, so thallium is slightly larger. Thallium’s electronegativity (Pauling scale ≈ 1.62) is modest – similar to aluminium – and lower than indium (1.78) or gallium (1.81). Its first ionization energy (~6.08 eV) is also relatively low (lower than indium’s ~5.78 eV), reflecting a large atomic size and electron shielding by the filled inner shells. These trends (larger radius, lower electronegativity and ionization energy than lighter congeners) are typical moving down a group, although thallium’s IE does not drop as much as expected due to the inert pair stabilization of the 6s electrons.

Thallium’s metallic character is pronounced: it is a poor metal (post-transition metal). It is very soft (Mohs hardness ~1.2, comparable to lead or tin) and can be cut with a knife at room temperature. Its mechanical strength is very low (yield strength only a few megapascals), but it is ductile and malleable. In bulk, thallium metal crystallizes in a hexagonal close-packed (hcp) lattice (space group P6₃/mmc) with lattice parameters a = 345.66 pm, c = 552.48 pm.

Isotopes and Nuclear Properties

Naturally occurring thallium consists of two stable isotopes: ^203Tl and ^205Tl. These two isotopes have nearly 30% and 70% natural abundance, respectively, and both have nuclear spin ½. They form the basis of thallium’s atomic mass (204.38 amu). ^203Tl and ^205Tl are the endpoints of a long-lived decay chain from stellar nucleosynthesis. A noteworthy derived isotope is ^201Tl (half-life ≈ 73 hours), which is produced in cyclotrons and used in medical imaging (see Applications). Thallium-201 decays by electron capture to ^201Hg, emitting characteristic x-rays and low-energy gamma rays useful for cardiac perfusion scans. Another isotope, ^204Tl (half-life ≈ 3.8 years), decays by beta emission and is important in environmental tracing but generates relatively little radiation dose. Heavier radioisotopes like ^206Tl (t½ ≈ 4.2 min) and ^207Tl (t½ ≈ 4.8 min) decay very rapidly to lead isotopes and have only niche scientific uses.

Because both stable isotopes have spin ½, thallium nuclei are NMR-active (like ^203Tl and ^205Tl NMR), though such NMR is specialized and not routine. ^205Tl NMR is sometimes used in research (thallium’s nuclear magnetic moment is moderately large), but outside nuclear medicine ^201Tl is the main isotope of technological importance (for SPECT imaging). Thallium has no stable isotope with an appreciable natural radioactivity (both its “long-lived” isotopes are effectively stable for practical purposes), unlike e.g. uranium or thorium.

Allotropes and Typical Compounds

Thallium has no distinct allotropes similar to carbon or phosphorus. Its only stable solid form is the metallic hcp phase; there are no widely known crystalline allotropes or molecular forms of elemental thallium. The metal is easily melted and resolidified, but this yields the same structure (no allotropic transformations). When molten, thallium is a silvery liquid metal (its melting point ~577 K).

Thallium’s chemistry is dominated by the monovalent Tl⁺ state; Tl³⁺ is analogous to Al³⁺ but is strongly oxidizing and relatively rare. Common Tl(I) compounds include:

Oxides and hydroxides: Thallium(I) oxide (Tl₂O) is a yellow solid formed when thallium metal is heated in oxygen. Tl₂O dissolves in water to give thallium(I) hydroxide, TlOH, a colorless, strongly basic solution. Indeed, TlOH is unusual among heavy-metal hydroxides: it readily dissolves CO₂ to form thallium(I) carbonate (Tl₂CO₃), the only soluble carbonate of any heavy metal. Thallium(III) oxide (Tl₂O₃) is a black oxide formed by oxidizing Tl₂O in strong oxidizers; it decomposes above ~800 °C back to Tl₂O. Tl(III) oxide is analogous to Al₂O₃.

Halides: Thallium(I) halides (TlCl, TlBr, TlI) are all fairly stable ionic compounds. For example, TlCl is a white salt (used historically as a reagent) and TlI is a yellow salt that is soluble in strong acids but photosensitive (yellow → red under light). Thallium polyhalides such as TlCl₂⁻ or TlCl₄²⁻ can occur in solution, reflecting Tl(I)·Cl⁻ interactions. Thallium(III) trihalides also exist but are powerful oxidizers – for example, TlF₃, TlCl₃ and TlBr₃ can be made by reacting Tl metal with excess halogen. These compounds (particularly TlF₃ and TlCl₃) are only stable at room temperature and in the absence of moisture. In water, TlCl₃ hydrolyzes to TlCl and chlorine gas. TlBr₃ and TlI₃ likewise are not stable in water.

Other salts: Thallium forms soluble salts with common anions in the +1 state: e.g. thallium nitrate (TlNO₃ – white, very soluble), Tl₂SO₄ (thallium sulfate), and TlAc (acetate). Thallium(I) thiocyanate (TlSCN) is soluble and used in certain analyses. Thallium(I) sulfide (Tl₂S) is a black solid analogous to Ag₂S.

Complexes: In chemistry, Tl⁺ behaves like a large alkali or coinage metal cation. A classic example is that Tl₂CO₃ dissolves in strong base to form the thallate complex [Tl(OH)₄]⁻, analogous to aluminate [Al(OH)₄]⁻. Thallium(I) also forms ammonia complexes (e.g. [Tl(NH₃)ₙ]⁺) and halide complexes TlCl₂]⁻, [TlCl₃]²⁻, etc.) under various conditions. Tl(III) forms coordination compounds similarly to other group-13 trications (for example, TlCl₃·6H₂O and Tl(NO₃)₃·3H₂O exist).

Special compounds:

* Scintillator crystals: Many inorganic scintillators incorporate thallium. The classic example is sodium iodide doped with thallium, NaI(Tl), a scintillating crystal that emits visible light when hit by X-rays or gamma rays. Cesium iodide can similarly be doped with thallium (CsI(Tl)). These compounds leverage the efficient energy transfer afforded by Tl⁺ emitters at ~550 nm. (These are not exactly stoichiometric salt compounds but solid solutions.)

Luminous and optical materials: Thallium’s heavy atomic weight and electron structure make certain thallium compounds useful in optics. For example, TlBr and TlI can form high-refractive-index semiconducting crystals used in infrared optics. Mixed halide glasses containing thallium (e.g. mixed TlBr–TlI glasses) are transparent in the IR and are used for special lenses and windows. Thallium oxide (Tl₂O) has also been studied in specialty glass formulations to expand transmission into the infrared.

* Sulfur and Selenium minerals: In nature, thallium often appears in rare minerals (see Occurrence). These are technically compounds: e.g. crooksite (CuTlSe₂) and lorandite (TlAsS₂) contain thallium bound to selenium or sulfur. These crystalline compounds have semiconducting properties and have been investigated for detectors (lorandite was famously studied as a possible solar neutrino detector because of its nuclear reactions).

In summary, thallium forms typical +1 antiseptic salts and a few +3 compounds. The +1 compounds tend to look and behave somewhat like those of potassium or silver (weak acid anhydrides, basic oxides, alkaline hydroxide, etc.), while +3 compounds are rarer, somewhat alkaline-earth-like or aluminum-like, and strongly oxidizing. Thallium does not form many complex organic or organometallic compounds; it is generally encountered as the simple salts described above.

Physical Properties

Thallium is a heavy, metallic element with characteristic properties: For bulk metal, density ≈ 11.85 g/cm³ (comparable to lead’s 11.34 g/cm³). It is very soft (Mohs hardness ~1.2) and highly malleable; it has a low tensile strength (a few MPa) similar to other poor metals. Thallium has a relatively low melting point of 302.9 °C (575.9 K) and boils at about 1457 °C (1730 K). Its specific heat capacity is low (~0.13 J/g·K), typical for heavy metals.

The thermal conductivity of thallium is moderate (~46 W/m·K at room temperature), and its electrical resistivity at 20 °C is about 180 nΩ·m (nanohm·meters). This corresponds to an electrical conductivity of roughly 5.6×10^6 S/m (nearly an order of magnitude lower conductivity than copper, which is ~58×10^6 S/m). In other words, thallium conducts electricity well as a metal, but not as well as the best conductors. Magnetically, thallium is strongly diamagnetic (all electrons paired). It also exhibits high electron reflectivity (shiny metal surface) until it tarnishes.

Crystallographically, solid thallium has hexagonal close-packed (hcp) symmetry (space group P6₃/mmc) at and below room temperature. The unit cell lattice constants are a = b = 345.66 pm and c = 552.48 pm, with the ideal close-packed c/a ratio. Under pressure or potentially at very low temperature no alternate allotropes are reported; the hcp structure persists. (Unlike tin, thallium does not have a low-temperature β–γ transition or multiple metallic phases.)

Spectroscopically, atomic thallium is known for bright optical emission lines, which historically signaled its discovery. The most prominent line is a green "resonance" transition around 535 nm (in vacuum). Indeed, a green emission near 535.04 nm is often cited as the signature line of thallium in flames or discharges. Other notable thallium lines occur in the visible and near-infrared: for example, ionized thallium (Tl II) has lines at 696.3 nm, 817.2 nm, 1162.6 nm, 1307.5 nm, and 1561.6 nm (observed in laboratory spectra), but neutral thallium’s lines around 535 nm and in the near-UV (at 290.4 nm and 293.1 nm) are most historically important. These spectral lines are used in atomic absorption/emission spectroscopy as calibration standards or as sensitive flame test indicators for thallium.

Thallium’s optical properties (for a metal) are similar to lead or tin. The pure metal is opaque and reflective at visible wavelengths, but it also has interband transitions in the blue–UV that give tarnished surfaces a dull gray-blue appearance. There are no naturally luminescent allotropes or unusual color phases of the metal itself.

Chemical Reactivity and Trends

Thallium’s chemical reactivity reflects its low first ionization energy and its heavy-metal character. The Tl⁺ ion, being large and singly charged, is relatively inert towards water removal reactions – in fact, freshly cut thallium metal does not react with dry air or moisture at room temperature. It tarnishes slowly in moist air, forming a thin, protective layer of Tl₂O (and perhaps Tl₂CO₃ if CO₂ is present). A shiny thallium sample can stay metallic for some time, unlike iron which rusts quickly. However, if heated strongly in air, thallium burns to Tl₂O and eventually Tl₂O₃ (the latter at very high temperature or in strong oxygen). The oxide layer passivates the surface, so bulk thallium is not extremely susceptible to corrosion, but it can slowly dissolve in water, producing thallium hydroxide and hydrogen gas: 2 Tl(s) + 2 H₂O(l) → 2 TlOH(aq) + H₂(g). This reaction proceeds slowly unless the water is acidified.

With halogens and other oxidizers, thallium reacts more vigorously. At room temperature, chlorine and bromine attack thallium metal to form thallium(III) halides (TlF₃, TlCl₃, TlBr₃) as noted above. Fluorine reacts explosively even with solid Tl. Iodine reacts similarly to give TlI (with some TlI₃ in excess iodine). Thallium thus forms +3 halides readily, unlike many of the lighter post-transition metals.

Acids attack thallium only slowly. For example, concentrated hydrochloric or sulfuric acid will dissolve thallium metal only sluggishly because the TlCl and Tl₂SO₄ produced are quite insoluble and form a coating. Similarly, nitric acid oxidizes thallium to Tl(III) and back-reduces to Tl(I) rapidly, making nitric acid somewhat ineffective at dissolving large amounts of thallium. In contrast, strong oxidizing acids (like hot perchloric acid) will convert thallium entirely to Tl(III) solutions.

Because the +1 state is so stable, most soluble thallium salts contain Tl⁺, which behaves nearly inertly as a cation. Thallium(I) salts are essentially ionic and often colorless or white (unless anion is colored). Tl⁺ complexes are typically 2-coordinate or linear (similar to Ag⁺) when complexed to organic ligands or halides, reflecting a 6s² “lone pair” effect. Tl(III) salts, when isolated, often hydrolyze. For example, Tl(III) chloride in water yields TlCl and Cl₂ spontaneously because Tl(III) is a strong oxidizer (reducing its own fraction to +1). Thus Tl(III) compounds behave chemically like strong oxidizing M(III) species (akin to permanganate or dichromate in reactivity).

In acid-base terms, thallium hydroxides are basic: TlOH dissolves in water to give strongly alkaline solutions. Thallium does not have an amphoteric hydroxide known in the +1 state; rather, TlOH simply dissolves and can convert to Tl₂O upon heating. Tl(III) hydroxide (Tl(OH)₃) is known in solution and is amphoteric (analogous to Al(OH)₃), but it quickly disproportionates unless strongly oxidizing conditions are maintained. For example, precipitates of Tl(OH)₃ in strong base can redissolve as [Tl(OH)₆]³⁻ complexes, analogous to [Al(OH)₄]⁻.

In the electrochemical series, thallium lies below copper and silver, meaning it is somewhat nobler than these, and it does not aggressively reduce water (i.e. it does not react vigorously with non-oxidizing acids). Thallium metal will slowly liberate hydrogen from dilute acids, but much more slowly than sodium or calcium; this is because Tl₂O is not very soluble. Compared with its group neighbors, thallium is less reactive: for instance, indium will dissolve more readily in HCl than thallium does. The trend is that thallium(I) salts are fairly stable and unreactive (it is the Tl(III) chemistry that tends to be reactive/oxidizing).

Corrosion/Passivation: Thallium metal, like many heavy post-transition metals, tarnishes and forms a passive oxide layer in air, which slows further corrosion. In practice, thallium metal can be stored under inert atmosphere or oil to prevent tarnishing. In aqueous environments, Tl⁺ stays dissolved or precipitates as insoluble Tl₂CO₃ or Tl₂S. Thallium is not known to form a strongly protective passivation like aluminum; rather, its slow oxidation and insoluble salts offer some protection.

Complexation: In solution, Tl⁺ can form complexes with soft ligands (Cl⁻, I⁻, CN⁻) due to its large size and Argentinelike affinity. For example, thallium(I) chloride is sparingly soluble, but in excess chloride it forms [TlCl₂]⁻ complexes. Thallium is often compared to copper(I) or silver(I) in this regard. In strongly alkaline solutions, thallium(I) forms soluble thallate complexes as noted. Thallium does not have a rich redox chemistry like transition metals, but Tl(III) salts are strong oxidizers (e.g. TlCl₃ in solution oxidizes iodide to iodine, etc.).

Occurrence and Production

Thallium is quite rare in the Earth’s crust — on the order of a few tenths of a part per million by weight. It is not mined as a primary ore; nearly all thallium is obtained as a byproduct of processing other ores. Thallium tends to occur in trace amounts in sulfide minerals of lead, zinc and copper. For example, thallium is found in minor quantities in the ores of zinc sulfide (sphalerite) and lead sulfide (galena). Some rare minerals are inherently thallium-rich: crooksite (CuTlSe₂), hutchinsonite ((Tl,Pb)₂As₅S₉) and lorandite (TlAsS₂) are thallium-bearing sulfide or selenide minerals, but these are very uncommon. Thallium also substitutes into potassium minerals like sylvite (KCl) and carnallite, and into cesium-bearing minerals like pollucite (CsAlSi₂O₆), though again at low levels.

Worldwide, the principal source of thallium is scrubbed flue dust from copper, zinc, or lead smelting. During the smelting of sulfide ores, volatile thallium collects in the acid plant gases; contacting these with diluted acid or chloride solutions extracts thallium into water or scrubber solutions. From such waste solutions, thallium is then recovered by precipitation and reduction. Historically, zinc smelters (e.g. in Belgium’s Vieille-Montagne refinery) and lead smelters (e.g. in Kazakhstan) have been large thallium producers. For example, the Ust-Kamenogorsk Metallurgical Complex in Kazakhstan processes lead-copper dust to yield on the order of 5–10 tonnes of metallic thallium per year. Thallium production is very small by tonnage (a few to a few dozen tonnes globally per year) compared with major metals.

Production process: A typical recovery route is to leach smelter dust with dilute sulfuric acid, precipitate thallium as Tl₂CO₃ or Tl₂S or basic Tl salts, then chemically reduce with iron or hydrogen to metal. In the Kazakh process, for instance, thallium is recovered by forming thallium(I) chromate (Tl₂CrO₄) which is then decomposed and reduced to yield Tl metal, which is cast into ingots or briquettes. (One method: scrap iron is added to a thallium salt solution to cement out Tl metal.) The metal is usually handled under inert atmosphere or stored under kerosene to prevent oxidation and vaporization (thallium has a high vapor pressure at melting point).

Abundance: In the broader environment, thallium occurs only at trace levels. Abundance by weight is roughly 0.000053% (0.53 ppm) in the continental crust. It is much lower in seawater (<0.001 ppb, since Tl strongly adsorbs to sediments) and in the human body (millions of times lower than major elements). Thallium is slightly concentrated in some plants (especially tobacco) and in soils near mining or smelting sites. Coal and fly ash can contain elevated thallium (as thallium follows potassium and sulfur when coal is burned).

Major Producers: Today the leading sources of thallium metal are in Kazakhstan and China. Kazzinc (Kazakhstan) recovers ~5–7 t/yr. China historically has produced thallium as well (often from zinc mines). The United States used to produce small quantities but stopped domestic production by the 1980s, instead importing all needed thallium.

Applications and Technology

Thallium’s toxicity severely limits its use, so its applications are confined to niche areas where other elements cannot substitute. About two-thirds of the thallium refined is used in electronics and specialty materials; the rest goes into glass, optics, and some medical uses. Key applications include:

Radiation detectors and scintillators: Thallium is essential to many scintillation detectors. The classic NaI(Tl) crystal (sodium iodide doped with thallium chloride at ~0.1% by weight) is the most widely used scintillator for gamma-ray spectroscopy. Thallium activator ions serve as luminescence centers, converting high-energy radiation into blue-green light (peak ~415 nm). Similarly, thallium-doped cesium iodide (CsI(Tl)) is a very bright, efficient scintillator used in X-ray imaging and high-energy physics detectors (e.g. cosmic-ray calorimeters). Thallium is also used as an additive in other scintillators (e.g. in some thallium-activated aluminum oxide fibers and crystals). These scintillators are found in medical imaging (gamma cameras, CT scanners), mineral exploration instruments, and radiation detection equipment.

Memory and photomultiplier devices: Thallium can be used in photomultiplier tube photocathodes (e.g. bialkali antimonide photocathodes often contain Tl, as “Sb-K-Cs” or “Sb-Na-K” with some Tl). The precise role is to adjust the work function and enhance sensitivity. Specialized photo-sensitive surfaces (e.g. for some night-vision or UV sensors) sometimes incorporate thallium compounds.

Semiconductor and infrared technology: Certain thallium compounds serve in optoelectronics. Thallium bromide (TlBr) and thallium iodide (TlI) are wide-bandgap semiconductors (bandgaps ~2.6–2.7 eV) with high atomic numbers (Tl=81) and density (~7.6 g/cm³ for TlBr). These materials are used experimentally as room-temperature X-ray and gamma-ray detectors because their high Z and density give high photon absorption, and they have decent charge transport properties. Research continues into TlBr detector crystals for medical imaging and astrophysics (e.g. the Brookhaven National Lab TlBr arrays). TlI has also been used in electroluminescent and cathodoluminescent phosphors. In telecommunications, thallium-doped fluoride glasses can transmit further into the infrared (>5 μm) than conventional glasses, for fiber optics and thermal imaging lenses.

Optical glasses and lenses: Thallium oxide is used to prepare specialized high-refractive-index glass. Thallium-doped barium crown glasses and thallium-containing lead-free optical glasses have been studied for infrared transparency (important for lenses in thermal imaging cameras). The Kazzinc product sheet notes *“optical glass, prisms” as applications of thallium metal. In practice, mill kilogram of Tl are added to exotic glass formulations to achieve particular refractive indices or infrared cutoff properties, replacing parts of lead oxide or other heavy-metal components.

High-temperature superconductors: Thallium plays a role in a class of cuprate superconductors discovered in the late 1980s. Compounds such as Tl₂Ba₂Ca₂Cu₃O₁₀ (often abbreviated Tl-2223) and TlBa₂Ca₂Cu₃O₉ (“Tl-1223”) can superconduct above 100 K (around 120 K was achieved). In these layered oxide ceramics, thallium acts as a charge reservoir layer between CuO₂ planes. While these materials are mainly of scientific interest, thallium-gallium barium copper oxides (Tl-Ba-Ca-Cu-O) demonstrated that thallium inclusion can raise the critical temperature (Tc) of superconductivity. Research into thallium-based cuprates helped advance understanding of high-Tc materials.

Catalysts and chemical research: There are a few catalytic or research uses for Tl compounds. Thallium(I) salts can act as homogeneous catalysts in organic chemistry, for example in certain organic synthesis steps (though any such use is rare due to toxicity). Thallium doping has been investigated in ferroelectric and piezoelectric ceramics (e.g. Tl–doped lead titanate to adjust properties), again mostly at a laboratory scale. Thallium(I) and Tl(III) salts have been used as reference chemicals in electrochemistry due to their well-known oxidation potentials.

Other historical/limited uses: Historically, thallium were used in rodenticides and insecticides (“thallium sulfate” was once a common rat poison until banned). It was also used for a short time in homoeopathy (under the name much-needed remedy?) and was studied for glaucoma treatment in the 1900s (thallium nitrate can lower eye pressure) – though such medical uses are obsolete today. Thallium salts were used in small amounts in anticoagulant rodenticides for laboratory use until alternatives took over. Additionally, some street-lamp and mercury-vapor lamp alloys include thallium compounds to modify the emission spectrum (the thallium line at 535 nm adds green to sodium lamps, for example, or thallium halides are dopants in metal-halide lamps to improve color rendering).

Overall, the key industries for thallium are specialized optics, high-energy detectors, and electronic materials. Many recording devices (NaI detectors, IR cameras) quietly rely on a few milligrams of thallium dopant. In all these cases, any benefit of thallium must outweigh its toxicity, so the applications tend to be ones where unique electronic or optical properties are essential and no safer substitute is available.

Biology, Environment, and Safety

Thallium is acutely and chronically highly toxic to living organisms. It has no known essential biological role — instead, it disrupts biological processes. Thallium(I) ions mimic potassium (K⁺) in biological systems (similar ionic radius and charge), so the body tends to accumulate thallium where potassium would go. This interference hampers enzymes and ion channels critical for nerve function and cell metabolism.

Toxic effects: Exposure to thallium (especially soluble thallium salts like TlNO₃ or Tl₂SO₄) can cause severe poisoning. Symptoms of thallium poisoning include gastrointestinal distress (nausea, vomiting, diarrhea), followed by neurological effects: numbness, paralysis, hallucinations, and in severe cases coma. A characteristic sign is alopecia (rapid hair loss) occurring days after exposure. Thallium also damages kidneys and blood cells. The lethal dose (LD₅₀) of thallium salts in humans is on the order of a few milligrams per kilogram of body weight; as little as 1–3 g total intake can be fatal to an adult. Because of these effects, thallium compounds are considered among the most dangerous heavy-metal poisons (often compared with arsenic and mercury). Chronic low-level exposure can cause peripheral neuropathy and other subtle harms.

Environmental cycling: Thallium released into the environment is persistent. It can enter soils from mining/smelter waste or coal combustion fly ash. Plants readily take up thallium through roots (similar to potassium uptake), so food crops can accumulate trace thallium if grown in contaminated soil. Tobacco plants are particularly notorious for accumulating thallium from soil, so cigarette smoke can be a source of exposure (smokers have roughly double the thallium body burden of non-smokers). Thallium in water can bioaccumulate in aquatic food chains to some extent, though information is sparse. In soil, thallium strongly adsorbs to clays and organic matter, so it tends to stick rather than leach away, forming a long-lasting hazard in polluted sites.

Regulations and exposure limits: Because of its toxicity, thallium is tightly regulated. Occupational exposure limits (e.g., OSHA PEL or ACGIH TLV in the United States) are typically around 0.1 mg/m³ for soluble thallium compounds (workplace air concentrations). The IDLH (immediately dangerous to life or health) for thallium compounds is about 15–20 mg/m³. For drinking water, agencies often set maximum contaminant levels in the low parts-per-billion range: for example, the U.S. EPA’s criterion is ~2 ppb (0.002 mg/L) for thallium in drinking water, and some guidelines allow up to 10–13 ppb as a limit for lifetime use. Soil action levels are similarly low; even tens of ppm in soil can be a concern because of ingestion risk. While elemental thallium metal is less bioavailable than soluble salts, any dust or fumes (e.g. from machining) would rapidly oxidize to soluble Tl⁺ salts.

Safety handling: In practice, thallium metal and compounds are handled in glove boxes or fume hoods. Personnel use gloves and protective clothing; inhalation masks or ventilated enclosures are mandatory when powdered or when producing Tl salts. A major safety incident is the lethal poisoning of workers who accidentally ingested thallium sulfate (it has no odor or taste). Today, thallium is classified as an extremely hazardous substance under many chemical safety laws. Industrial processes that produce thallium must have effluent controls; thallium is considered a priority pollutant.

Biology: No life-form uses thallium beneficially; it is purely an environmental contaminant. At very low levels, however, the human body can excrete small amounts of thallium via urine. There is an FDA-approved antidote: “Prussian blue” (ferric hexacyanoferrate) binds thallium ions in the gut, preventing reabsorption and promoting excretion. Chelating agents like DTPA or EDTA have limited effect on Tl, so the main treatment is Prussian blue plus supportive care.

In summary, thallium is one of the most dangerous non-radioactive poisons known. Strict regulatory controls and safer alternatives have greatly reduced its prevalence; for example, it is no longer used in consumer products or medicines, removed from rodenticides, and emissions from industry are monitored. Any release or use today requires careful management to avoid severe health and ecological impacts.

History and Etymology

Thallium was discovered in 1861 by British chemist Sir William Crookes. Using the new technique of flame spectroscopy, Crookes heated a sample of pyrites (iron sulfide) residues from a sulfuric acid plant and observed an intense green spectral line that did not correspond to any known element. This “brilliant green” emission prompted him to identify a new element. Crookes named it thallium from the Greek thallos (θαλός, “green shoot” or “twig”), referring to the green line in the spectrum.

At nearly the same time, in 1862 French chemist Claude-Auguste Lamy independently discovered the same element while analyzing residues from a pyrite smelter and noted the green spectral line. There was a dispute over priority, but Crookes published first in the Philosophical Transactions of the Royal Society (1863) and is generally credited as discoverer. Both scientists isolated minuscule amounts of thallium by precipitation from residues, and determined some of its properties. Crookes received the first Davy Medal (1860) in recognition of his contributions to the isolation of thallium.

The 1860s also saw early characterization: it was found that thallium metal is soft and quickly tarnishes in air. In 1869 Dmitri Mendeleev used thallium’s properties (atomic weight ~204, valence ~+1) to place it in his periodic table under aluminium and gallium, predicting some properties correctly. Early chemists discovered thallium(I) hydroxide and salts (notably the very soluble nitrate TlNO₃ and the insoluble sulfate Tl₂SO₄), and noted that thallium(III) compounds were strong oxidizers.

Historically, thallium compounds gained notoriety as poisons. In the early 20th century, thallium sulfate was widely used as an insecticide and rat poison (cheaper and more odorless than arsenic). It earned monikers like “the poisoner’s poison” and “inheritance powder,” as it was colorless and tasteless and caused a gruesome poisoning syndrome (hair loss, etc.). Many criminal and accidental thallium poisonings occurred before its dangers were fully appreciated. By the 1970s, thallium compounds were banned for pesticide or rodenticide use in most countries.

In technology, thallium’s story is more recent. Its use in optics and electronics grew in the late 20th century. For example, thallium-based semiconductors and scintillators came into use only after mid-century. The role of thallium in high-Tc superconductors emerged in 1987 when a thallium-containing copper oxide was found to superconduct at 115 K (Nobel Prize-winning work followed). This showed thallium could be a key component in layered perovskite superconductors.

The etymology remains straightforward: Thallium (Latinized) from thallos (Greek for green shoot). The choice of the “-ium” ending reflects its metal status. In historical footnotes, the symbolism of a “green shoot” evokes thallium’s green spectral line and is an early example of chemical nomenclature linking color to element names (like cesium from kaïkos for “sky-blue”).

Data Table

Property	Value
Atomic number (Z)	81
Symbol	Tl
Element category	Poor (post-transition) metal
Group / Period / Block	13 (IIIA) / 6 / p-block
Standard atomic weight	204.38 (unified mass units)
Phase at STP	Solid
Density (20 °C)	≈ 11.85 g·cm⁻³
Melting point	302.9 °C (≈ 576 K)
Boiling point	1457 °C (≈ 1730 K)
Crystal structure	HCP (hexagonal close-packed)
Molar volume	≈ 17.22 cm³·mol⁻¹ (at 25 °C)
Electrical resistivity	~180 nΩ·m (at 20 °C)
Thermal conductivity	46 W·m⁻¹·K⁻¹ (at 300 K)
Magnetic ordering	Diamagnetic
Electron configuration	[Xe] 4f¹⁴ 5d¹⁰ 6s² 6p¹
Common oxidation states	+1 (major), +3 (minor)
Electronegativity (Pauling)	1.62
Ionization energy (1st)	6.108 eV (589.00 kJ·mol⁻¹)
Atomic radius (covalent)	~148 pm (single-bond)
Isotopes (stable)	^203Tl (29.5% natural), ^205Tl (70.5%)
Nuclear spins (stable)	^203Tl, ^205Tl each I = 1/2
CAS Registry Number	7440-28-0

This concise data table summarizes thallium’s essential physical constants, isotopic data, and classification for quick reference. All values are approximate and given in SI units or standard atomic mass units as appropriate.