Thorium compounds

Many compounds of thorium are known: this is because thorium and uranium are the most stable and accessible actinides and are the only actinides that can be studied safely and legally in bulk in a normal laboratory. As such, they have the best-known chemistry of the actinides, along with that of plutonium, as the self-heating and radiation from them is not enough to cause radiolysis of chemical bonds as it is for the other actinides. While the later actinides from americium onwards are predominantly trivalent and behave more similarly to the corresponding lanthanides, as one would expect from periodic trends, the early actinides up to plutonium (thus including thorium and uranium) have relativistically destabilised and hence delocalised 5f and 6d electrons that participate in chemistry in a similar way to the early transition metals of group 3 through 8: thus, all their valence electrons can participate in chemical reactions, although this is not common for neptunium and plutonium.

A thorium atom has 90 electrons, of which four are valence electrons. Four atomic orbitals are theoretically available for the valence electrons to occupy: 5f, 6d, 7s, and 7p. However, the 7p orbital is greatly destabilised and hence it is not occupied in the ground state of any thorium ion. Despite thorium's position in the f-block of the periodic table, it has an anomalous [Rn]6d²7s² electron configuration in the ground state, as the 5f and 6d subshells in the early actinides are very close in energy, even more so than the 4f and 5d subshells of the lanthanides. However, in metallic thorium, the [Rn]5f¹6d¹7s² configuration is a low-lying excited state and hence the 5f orbitals contribute, existing in a rather broad energy band. In fact, the 5f subshells of the actinides have a larger spatial extent than the 4f orbitals of the lanthanides and thus actinide compounds have greater covalent character than the corresponding lanthanide compounds, leading to a more extensive coordination chemistry for the actinides than the lanthanides.

The ground-state electron configurations of thorium ions are as follows: Th⁺, [Rn]6d²7s¹; Th²⁺, [Rn]5f¹6d¹; Th³⁺, [Rn]5f¹; Th⁴⁺, [Rn]. This shows the increasing stabilisation of the 5f orbital as ion charge increases; however, this stabilisation is insufficient to chemically stabilise Th³⁺ with its lone 5f valence electron, and therefore the stable and most common form of thorium in chemicals is Th⁴⁺ with all four valence electrons lost, leaving behind an inert core of inner electrons with the electron configuration of the noble gas radon. The first ionisation potential of thorium was measured to be (6.08 ± 0.12) eV in 1974; more recent measurements have refined this to 6.3067 eV.

Thorium is a highly reactive and electropositive metal. At standard temperature and pressure, it is slowly attacked by water, but does not readily dissolve in most common acids, with the exception of hydrochloric acid. It dissolves in concentrated nitric acid containing a small amount of catalytic fluoride or fluorosilicate ions; if these are not present, passivation can occur, similarly to uranium and plutonium. At high temperatures, it is easily attacked by oxygen, hydrogen, nitrogen, the halogens, and sulfur. It can also form binary compounds with carbon and phosphorus. When thorium dissolves in hydrochloric acid, a black insoluble residue, probably ThO(OH, Cl)H is left behind, similarly to protactinium and uranium.

Finely divided thorium metal presents a fire hazard due to its pyrophoricity and must therefore be handled carefully. When heated in air, thorium turnings ignite and burn brilliantly with a white light to produce the dioxide. In bulk, the reaction of pure thorium with air is slow, although corrosion may eventually occur after several months; most thorium samples are however contaminated with varying degrees of the dioxide, which greatly accelerates corrosion. Such samples slowly tarnish in air, becoming grey and finally black at the surface. The impermeability of the oxide layer of thorium contrasts with that of the later actinides and conforms to the trend of increasing electropositivity and reactivity as the actinide series is traversed.

The most important oxidation state of thorium is +4, represented in compounds such as thorium dioxide (ThO₂) and thorium tetrafluoride (ThF₄), although some compounds are known with thorium in lower formal oxidation states. Owing to thorium(IV)'s lack of electrons in 6d and 5f orbitals, the tetravalent thorium compounds are colourless. Th³⁺ compounds are uncommon due to the large negative reduction potential of the Th⁴⁺/Th³⁺ couple. In 1997, reports of amber Th³⁺ (aq) being generated from thorium tetrachloride and ammonia were published: the ion was supposedly stable for about an hour before it was oxidised by water. However, the reaction was shown the next year to be thermodynamically impossible and the more likely explanation for the signals was azido-chloro complexes of thorium(IV). In fact, the redox potentials of thorium, protactinium, and uranium are much more similar to those of the d-block transition metals than the lanthanides, reflecting their historic placement prior to the 1940s as the heaviest members of groups 4, 5, and 6 in the periodic table respectively.

In aqueous solution, thorium occurs exclusively as the tetrapositive aqua ion [Th(H
₂O)
₉]⁴⁺
, which has tricapped trigonal prismatic molecular geometry: at pH < 3, the solutions of thorium salts are dominated by this cation. The Th–O bond distance is (245 ± 1) pm, the coordination number of Th⁴⁺ is (10.8 ± 0.5), the effective charge is 3.82 and the second coordination sphere contains 13.4 water molecules. The Th⁴⁺ ion is relatively large and is the largest of the tetrapositive actinide ions, and depending on the coordination number can have a radius between 0.95 and 1.14 Å. The thorium(IV) hydrated ion is quite acidic due to its high charge, slightly stronger than sulfurous acid: thus it tends to undergo hydrolysis and polymerisation, predominantly to [Th₂(OH)₂]⁶⁺ in solutions with pH 3 or below, but in more alkaline solution polymerisation continues until the gelatinous hydroxide is formed and precipitates out (though equilibrium may take weeks to be reached, because the polymerisation usually slows down significantly just before the precipitation): this behaviour is similar to that of plutonium(IV).

Large coordination numbers are the rule: thorium nitrate pentahydrate was the first known example of coordination number 11, the oxalate tetrahydrate has coordination number 10, and the Th(NO
₃)⁻
₆ anion in the calcium and magnesium salts is 12-coordinate. Due to the large size of the Th⁴⁺ cation, thorium salts have a weaker tendency to hydrolyse than that of many multiply charged ions such as Fe³⁺, but hydrolysis happens more readily at pH above 4, forming various polymers of unknown nature, culminating in the formation of the gelatinous hydroxide: this behaviour is similar to that of protactinium, which also hydrolyses readily in water to form colloidal precipitates. The distinctive ability of thorium salts is their high solubility, not only in water, but also in polar organic solvents. As a hard Lewis acid, Th⁴⁺ favours hard ligands with oxygen atoms as donors: complexes with sulfur atoms as donors are less stable.