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.Greenwood and Earnshaw, p. 1265 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.Greenwood and Earnshaw, p. 1266
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.Wickleder et al., pp. 59–60Golub et al., pp. 222–7 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.David R. Lide (ed), CRC Handbook of Chemistry and Physics, 84th Edition. CRC Press. Boca Raton, Florida, 2003; Section 10, Atomic, Molecular, and Optical Physics; Ionization Potentials of Atoms and Atomic Ions
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.Wickleder et al., pp. 61–63 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.Greenwood and Earnshaw, p. 1264
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.Wickleder et al., pp. 64–6Wickleder et al., pp. 70–7Wickleder et al., pp. 78–94 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 , which has tricapped trigonal prismatic molecular geometry:Wickleder et al., pp. 117–134 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).Greenwood and Earnshaw, p. 1275–7
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 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.
The standard reduction potentials in acidic aqueous solution for some common thorium ions are given below:Greenwood and Earnshaw, p. 1263
| Th4+ + e−|| Th3+ || E0 = −3.8 V |- | Th4+ + 4e−|| Th || E0 = −1.83 V
Oxides and hydroxides
In air, thorium burns to form the binary oxide thorium dioxide, ThO2, also called thoria or thorina. Thoria, a refractory material, has the highest melting point (3390 °C) of all known oxides. It is somewhat hygroscopic and reacts readily with water and many gases, but dissolves easily in concentrated nitric acid in the presence of fluoride, like neptunium dioxide and plutonium dioxide do. When heated, it emits intense blue light, which becomes white when mixed with its lighter homologue cerium dioxide (CeO2, ceria): this is the basis for its previously common application in gas mantles.
Reports of thorium peroxide, initially supposed to be Th2O7 and be formed from reacting thorium salts with hydrogen peroxide, were later discovered to contain both peroxide anions and the anions of the reacting thorium salt. The mixed hydroxide-hydroperoxide has been prepared. It decomposes at 120 °C to yield thorium dioxide and water.
Thorium monoxide has been produced through laser ablation of thorium in the presence of oxygen. This highly polar molecule is calculated to have one of the largest known internal electric fields."The ACME EDM Experiment." electronedm.org
Thorium hydroxide, Th(OH)4, can be prepared by adding a hydroxide of ammonium or an alkali metal to a thorium salt solution, where it appears as a gelatinous precipitate that will dissolve in dilute acids, among other substances. It can also be prepared by electrolysis of thorium nitrates. It is stable from 260 to 450 °C; at 470 °C and above it continuously decomposes to become thoria. It easily absorbs atmospheric carbon dioxide to form the hydrated carbonate ThOCO3·xH2O and, under high-pressure conditions in a carbon dioxide atmosphere, or . Several mixed oxides are known, such as BaThO3, which has the perovskite structure.Greenwood and Earnshaw, p. 1269
Halides
All four thorium tetrahalides are known, as are some low-valent bromides and iodides: the tetrahalides are all hygroscopic compounds. However, with the exception of thorium tetrafluoride, which is sparingly soluble in water, the other thorium tetrahalides readily dissolve in polar solvents such as water. Additionally, many related polyhalide ions are also known. Thorium tetrafluoride (ThF4, white, m.p. 1068 °C) is most easily produced by reacting various thorium salts, thoria, or thorium hydroxide with hydrogen fluoride: methods that involve steps in the aqueous phase are more difficult because they result in hydroxide and oxide fluorides that have to be reduced with hydrogen fluoride or fluorine gas. It has a monoclinic crystal structure and is isotypic with zirconium tetrafluoride and hafnium tetrafluoride, where the Th4+ ions are coordinated with F− ions in somewhat distorted square antiprisms. It is a white, hygroscopic powder: at temperatures above 500 °C, it reacts with atmospheric moisture to produce the oxyfluoride ThOF2.
Thorium tetrachloride (ThCl4, white, m.p. 770 °C) is produced by heating thoria in an organochloride compound such as carbon tetrachloride.Greenwood and Earnshaw, p. 1271 The usual method of purification is crystallisation from an aqueous solution and then heating the product above 100 °C to dehydrate it. Further purification can be achieved by subliming it. Its melting and boiling points are respectively 770 °C and 921 °C. It undergoes a phase transition at 405 °C, with a low-temperature α phase and high-temperature β phase. Nevertheless, the β phase usually persists below the transition temperature. Both phases crystallise in the tetragonal crystal system and the structural differences are small. Below −203 °C, a low-temperature form exists with a complex structure.