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.
General chemistry
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 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.