Thallium – facts and properties

Position in the periodic table

View the full periodic table

Occurrence and extraction

As an element thallium is scarce and almost never forms deposits of its own. It is usually present in tiny amounts in sulfide ores of zinc, lead and copper and in pyrite. The metal is obtained as a by-product of roasting and smelting these ores: it collects in flue dust and smelter residues, is leached out and is reduced electrolytically to the metal. World production is only about ten tonnes a year.

Its discovery is a textbook case of spectral analysis, which Robert Bunsen and Gustav Kirchhoff had developed shortly before and which had already revealed caesium and rubidium. In 1861 the English chemist William Crookes used it in London to examine residues from a sulfuric acid factory. He saw a previously unknown, bright green line and concluded that there was a new element. He named it after the Greek thallos, a green twig, for the fresh colour of the line. The following year Crookes and, independently of him, the Frenchman Claude-Auguste Lamy produced the metal in larger amounts.

Properties and reactions

Thallium is in period 6 and group 13, below indium, to the right of mercury and to the left of lead. Its configuration, [Xe] 4f¹⁴ 5d¹⁰ 6s² 6p¹, ends in a single 6p electron, and its electronegativity is 1.62. The metal melts at 304 °C, boils at 1473 °C and has a density of 11.8 g/cm³. It is so soft that it can be cut with a knife, and fresh cuts quickly turn grey in air.

Thallium has 3 valence electrons, yet oxidation number plus three is not the one it prefers. Because the two 6s electrons are held especially tightly, they often stay out of bonding, an effect known as the inert pair effect. That is why oxidation number plus one is the most stable, and thallium(III) compounds are strong oxidising agents.

The thallium(I) ion is similar in size to the potassium ion and carries the same single charge. The body can barely tell them apart: thallium enters cells by the same transport routes as potassium and disrupts important enzymes there. This explains its toxicity. Poisoning often shows itself only after days, with stomach and bowel complaints and painful nerve damage, and later the hair falls out. It can be fatal. Prussian blue, which binds thallium in the gut, is used as an antidote. Thallium and its compounds have no place in school experiments.

Uses in everyday life and technology

There is no broad everyday use. In the twentieth century thallium sulfate was sold as rat poison and ant poison. Because it also endangers people and pets, it is banned in most countries. Thallium was also misused for murder several times, much like arsenic, which was once nicknamed “inheritance powder”.

Today it is used mainly in measurement technology. Sodium iodide crystals doped with a little thallium give off short flashes of light when gamma rays or energetic particles strike them. Such scintillation counters are found in radiation detectors, in medical technology and in prospecting for raw materials. Crystals of thallium bromide and thallium iodide are transparent to infrared light and serve as windows and lenses in infrared instruments.

In medicine, thallium-201 is used, a radioactive isotope with a half-life of about three days that emits X-rays and gamma rays. Heart muscle cells take up thallium ions as they do potassium ions, so a camera shows which regions of the heart are well supplied with blood. The procedure is called myocardial scintigraphy. It is performed only by specialist staff with small amounts, and today other methods partly replace it.

Quick question

What is the atomic number of thallium (Tl)?

FAQ