Lanthanide

Lanthanides in the periodic table

The lanthanide /ˈlænθənaɪd/ or lanthanoid /ˈlænθənɔɪd/ series of chemical elements comprises the fifteen metallic chemical elements with atomic numbers 57 through 71, from lanthanum through lutetium. These fifteen lanthanide elements, along with the chemically similar elements scandium and yttrium, are often collectively known as the rare earth elements.

The informal chemical symbol Ln is used in general discussions of lanthanide chemistry to refer to any lanthanide. All but one of the lanthanides are f-block elements, corresponding to the filling of the 4f electron shell; lutetium, a d-block element, is also generally considered to be a lanthanide due to its chemical similarities with the other fourteen. All lanthanide elements form trivalent cations, Ln³⁺, whose chemistry is largely determined by the ionic radius, which decreases steadily from lanthanum to lutetium.

They are termed as lanthanides because the lighter elements in the series are chemically similar to lanthanum. Strictly speaking, both lanthanum and lutetium have been labeled as group 3 elements, because they both have a single valence electron in the d shell. However, both elements are often included in any general discussion of the chemistry of the lanthanide elements.

In presentations of the periodic table, the lanthanides and the actinides are customarily shown as two additional rows below the main body of the table, with placeholders or else a selected single element of each series (either lanthanum and actinium, or lutetium and lawrencium) shown in a single cell of the main table, between barium and hafnium, and radium and rutherfordium, respectively. This convention is entirely a matter of aesthetics and formatting practicality; a rarely used wide-formatted periodic table inserts the lanthanide and actinide series in their proper places, as parts of the table's sixth and seventh rows (periods).

Lanthanides
Lan­thanum 57 La	Cerium 58 Ce	Praseo­dymium 59 Pr	Neo­dymium 60 Nd	Prome­thium 61 Pm	Sama­rium 62 Sm	Europ­ium 63 Eu	Gadolin­ium 64 Gd	Ter­bium 65 Tb	Dyspro­sium 66 Dy	Hol­mium 67 Ho	Erbium 68 Er	Thulium 69 Tm	Ytter­bium 70 Yb	Lute­tium 71 Lu
Primordial From decay Synthetic Border shows natural occurrence of the element

Etymology

Together with scandium and yttrium, the trivial name "rare earths" is sometimes used to describe all the lanthanides. This name arises from the minerals from which they were isolated, which were uncommon oxide-type minerals. However, the use of the name is deprecated by IUPAC, as the elements are neither rare in abundance nor "earths" (an obsolete term for water-insoluble strongly basic oxides of electropositive metals incapable of being smelted into metal using late 18th century technology). Cerium is the 26th most abundant element in the Earth's crust, neodymium is more abundant than gold and even thulium (the least common naturally occurring lanthanide) is more abundant than iodine, which is itself common enough for biology to have evolved critical usages thereof. Despite their abundance, even the technical term "lanthanides" could be interpreted to reflect a sense of elusiveness on the part of these elements, as it comes from the Greek λανθανειν (lanthanein), "to lie hidden". However, if not referring to their natural abundance, but rather to their property of "hiding" behind each other in minerals, this interpretation is in fact appropriate. The etymology of the term must be sought in the first discovery of lanthanum, at that time a so-called new rare earth element "lying hidden" in a cerium mineral, and it is an irony that lanthanum was later identified as the first in an entire series of chemically similar elements and could give name to the whole series. The term "lanthanide" was introduced by Victor Goldschmidt in 1925.

Physical properties of the elements

Chemical element	La	Ce	Pr	Nd	Pm	Sm	Eu	Gd	Tb	Dy	Ho	Er	Tm	Yb	Lu
Atomic number	57	58	59	60	61	62	63	64	65	66	67	68	69	70	71
Image
Density (g/cm3)	6.162	6.770	6.77	7.01	7.26	7.52	5.244	7.90	8.23	8.540	8.79	9.066	9.32	6.90	9.841
Melting point (°C)	920	795	935	1024	1042	1072	826	1312	1356	1407	1461	1529	1545	824	1652
Boiling point (°C)	3464	3443	3520	3074	3000	1794	1529	3273	3230	2567	2720	2868	1950	1196	3402
Atomic electron configuration*	5d1	4f15d1	4f3	4f4	4f5	4f6	4f7	4f75d1	4f9	4f10	4f11	4f12	4f13	4f14	4f145d1
Metal lattice (RT)	dhcp	fcc	dhcp	dhcp	dhcp	**	bcc	hcp	hcp	hcp	hcp	hcp	hcp	hcp	hcp
Metallic radius pm	162	181.8	182.4	181.4	183.4	180.4	208.4	180.4	177.3	178.1	176.2	176.1	175.9	193.3	173.8
Resistivity (25 °C) /μ Ohm cm	57–80 20 °C	73	68	64		88	90	134	114	57	87	87	79	29	79
mag susceptibility χmol /10−6(cm3·mol−1)	+95.9	+2500 (β)	+5530(α)	+5930 (α)		+1278(α)	+30900	+185000 (350 K)	+170000 (α)	+98000	+72900	+48000	+24700	+67 (β)	+183

* Between initial [Xe] and final 6s² electronic shells ** Sm has a close packed structure like the other lanthanides but has an unusual 9 layer repeat

Gschneider and Daane (1988) attribute the trend in melting point which increases across the series, (lanthanum (920 °C) – lutetium (1622 °C)) to the extent of hybridisation of the 6s, 5d and 4f orbitals. The hybridisation is believed to be at its greatest for cerium which has the lowest melting point of all, 795 °C. The lanthanide metals are soft, their hardness increases across the series. Europium stands out as it has the lowest density in the series at 5.24 g/cm³ and the largest metallic radius in the series at 208.4 pm. It can be compared to barium which has a metallic radius of 222 pm. It is believed that the metal contains the larger Eu²⁺ ion and that there are only two electrons in the conduction band. Ytterbium also has large metallic radius and a similar explanation is suggested. The resistivities of the lanthanide metals are relatively high, ranging from 29 to 134 μ Ohm·cm. These values can be compared to a good conductor such as aluminium which has a resistivity of 2.655 μ Ohm·cm. With the exceptions of La, Yb and Lu (which have no unpaired f electrons) the lanthanides are strongly paramagnetic and this is reflected in their magnetic susceptibilities. Gadolinium becomes ferromagnetic at below 16 °C (Curie point). The other heavier lanthanides, terbium, dysprosium, holmium, erbium, thulium and ytterbium become ferromagnetic at much lower temperatures.

Chemistry and compounds

Chemical element	La	Ce	Pr	Nd	Pm	Sm	Eu	Gd	Tb	Dy	Ho	Er	Tm	Yb	Lu
Atomic number	57	58	59	60	61	62	63	64	65	66	67	68	69	70	71
Ln3+ electron configuration*	4f0	4f1	4f2	4f3	4f4	4f5	4f6	4f7	4f8	4f9	4f10	4f11	4f12	4f13	4f14
Ln3+ radius (pm)	103	102	99	98.3	97	95.8	94.7	93.8	92.3	91.2	90.1	89	88	86.8	86.1
Ln4+ ion colour in aqueous solution	—	Orange-yellow	Yellow	Blue-violet	—	—	—	—	Red-brown	Orange-yellow	—	—	—	—	—
Ln3+ ion colour in aqueous solution	Colorless	Colorless	Green	Violet	Pink	Pale yellow	Colorless	Colorless	V. pale pink	Pale yellow	Yellow	Rose	Pale green	Colorless	Colorless
Ln2+ ion colour in aqueous solution	—	—	—	—	—	Blood red	Colorless	—	—	—	—	—	Violet-red	Yellow-green	—

* Between initial [Xe] and final 6s² electronic shells

It should be noted that the colours of lanthanide complexes originate almost entirely from charge transfer interactions between the metal and the ligand. f → f transitions are symmetry forbidden (or Laporte-forbidden), which is also true of transition metals, however t-metals are able to use vibronic coupling to break this rule. The valence orbitals in lanthanides are almost entirely non-bonding and as such little effective vibronic coupling takes, hence the spectra from f → f transitions are much weaker and narrower than those from d → d transitions. In general this makes the colours of lanthanide complexes far fainter than those of transition metal complexes.

Approximate colors of lanthanide ions in aqueous solution

Oxidation state

57

58

59

60

61

62

63

64

65

66

67

68

69

70

71

+2

Sm2+

Eu2+

Tm2+

Yb2+

+3

La3+

Ce3+

Pr3+

Nd3+

Pm3+

Sm3+

Eu3+

Gd3+

Tb3+

Dy3+

Ho3+

Er3+

Tm3+

Yb3+

Lu3+

+4

Ce4+

Pr4+

Nd4+

Tb4+

Dy4+

Effect of 4f orbitals

Going across the lanthanides, in the periodic table, the 4f orbitals are usually being filled. The effect of the 4f orbitals on the chemistry of the lanthanides is profound and is the factor that distinguishes them from the transition metals. There are seven 4f orbitals and there are two different ways in which they are depicted, firstly as a "cubic set" or as a general set. The cubic set is f_z³, f_xz², f_yz², f_xyz, f_z(x²−y²), f_{x(x²−3y²)} and f_{y(3x²−y²)}. The 4f orbitals penetrate the [Xe] core and are isolated and do not participate in bonding. This explains why crystal field effects are small and why they do not form π bonds. As there are seven 4f orbitals the number of unpaired electrons can be as high as 7 which gives rise to the large magnetic moments observed for lanthanide compounds. Measuring the magnetic moment can be used to investigate the 4f electron configuration and this is a useful tool in providing an insight into the chemical bonding. The lanthanide contraction, the reduction in size of the Ln³⁺ ion from La³⁺(103 pm)- Lu³⁺(86.1 pm) is often explained by the poor shielding of the 5s and 5p electrons by the 4f electrons.

Lanthanide oxides: clockwise from top center: praseodymium, cerium, lanthanum, neodymium, samarium and gadolinium.

The electronic structure of the lanthanide elements, with minor exceptions, is [Xe]6s²4fⁿ. The chemistry of the lanthanides is dominated by the +3 oxidation state and in Ln^III compounds the 6s electrons and (usually) one 4f electron are lost and the ions have the configuration [Xe]4f^m. All the lanthanide elements exhibit the oxidation state +3. In addition Ce³⁺ can lose its single f electron to form Ce⁴⁺ with the stable electronic configuration of xenon. Also, Eu³⁺ can gain an electron to form Eu²⁺ with the f⁷ configuration which has the extra stability of a half-filled shell. Other than Ce(IV) and Eu(II), none of the lanthanides are stable in oxidation states other than +3 in aqueous solution. Promethium is effectively a man-made element as all its isotopes are radioactive with half-lives shorter than 20 years.

In terms of reduction potentials, the Ln^0/3+ couples are nearly the same for all lanthanides, ranging from −1.99 (for Eu) to −2.35 V (for Pr). Thus, these metals are highly reducing, with reducing power similar to alkaline earth metals such as Mg (−2.36 V).

Lanthanide oxidation states

All of the lanthanide elements are commonly known to have the +3 oxidation state and it was thought that only samarium, europium, and ytterbium had the +2 oxidation readily accessible in solution. Now, it is known that all of the lanthanides can form +2 complexes in solution.

Ionization energies and reduction potentials of the elements

Chemical element	La	Ce	Pr	Nd	Pm	Sm	Eu	Gd	Tb	Dy	Ho	Er	Tm	Yb	Lu
Atomic number	57	58	59	60	61	62	63	64	65	66	67	68	69	70	71
electron configuration above [Xe] core	5d16s2	4f15d16s2	4f36s2	4f46s2	4f56s2	4f66s2	4f76s2	4f75d16s2	4f96s2	4f106s2	4f116s2	4f126s2	4f136s2	4f146s2	4f145d16s2
E° Ln4+/Ln3+		1.72
E° Ln3+/Ln2+	57			−2.6		−1.55	−0.35			−2.5			−2.3	−1.05
E° Ln3+/Ln	−2.38	−2.34	−2.35	−2.32	−2.29	−2.30	−1.99	−2.28	−2.31	−2.29	−2.33	−2.32	−2.32	−2.22	−2.30
1st Ionisation energy (kJ·mol−1)	538	541	522	530	536	542	547	595	569	567	574	581	589	603	513
2nd Ionisation energy (kJ·mol−1)	1067	1047	1018	1034	1052	1068	1085	1172	1112	1126	1139	1151	1163	1175	1341
1st + 2nd Ionisation energy (kJ·mol−1)	1605	1588	1540	1564	1588	1610	1632	1767	1681	1693	1713	1732	1752	1778	1854
3rd Ionisation energy (kJ·mol−1)	1850	1940	2090	2128	2140	2285	2425	1999	2122	2230	2221	2207	2305	2408	2054
1st+2d+3rd Ionisation energy (kJ·mol−1)	3455	3528	3630	3692	3728	3895	4057	3766	3803	3923	3934	3939	4057	4186	3908
4th Ionisation energy (kJ·mol−1)	4819	3547	3761	3900	3970	3990	4120	4250	3839	3990	4100	4120	4120	4203	4370

The ionisation energies for the lanthanides can be compared with aluminium. In aluminium the sum of the first three ionisation energies is 5139 kJ·mol⁻¹, whereas the lanthanides fall in the range 3455 – 4186 kJ·mol⁻¹. This correlates with the highly reactive nature of the lanthanides.

The sum of the first two ionisation energies for europium, 1632 kJ·mol⁻¹ can be compared with that of barium 1468.1 kJ·mol⁻¹ and europium's third ionisation energy is the highest of the lanthanides. The sum of the first two ionisation energies for ytterbium are the second lowest in the series and its third ionisation energy is the second highest. The high third ionisation energy for Eu and Yb correlate with the half filling 4f⁷ and complete filling 4f¹⁴ of the 4f sub shell, and the stability afforded by such configurations due to exchange energy. Europium and ytterbium form salt like compounds with Eu²⁺ and Yb²⁺, for example the salt like dihydrides,. Both europium and ytterbium dissolve in liquid ammonia forming solutions of Ln²⁺(NH₃)_x again demonstrating their similarities to the alkaline earth metals.

The relative ease with which the 4th electron can be removed in cerium and (to a lesser extent praseodymium) indicates why Ce(IV) and Pr(IV) compounds can be formed, for example CeO₂ is formed rather than Ce₂O₃ when cerium reacts with oxygen.

Separation of lanthanides

The similarity in ionic radius between adjacent lanthanide elements makes it difficult to separate them from each other in naturally occurring ores and other mixtures. Historically, the very laborious processes of cascading and fractional crystallization were used. Because the lanthanide ions have slightly different radii, the lattice energy of their salts and hydration energies of the ions will be slightly different, leading to a small difference in solubility. Salts of the formula Ln(NO₃)₃·2NH₄NO₃·4H₂O can be used. Industrially, the elements are separated from each other by solvent extraction. Typically an aqueous solution of nitrates is extracted into kerosene containing tri-n-butylphosphate. The strength of the complexes formed increases as the ionic radius decreases, so solubility in the organic phase increases. Complete separation can be achieved continuously by use of countercurrent exchange methods. The elements can also be separated by ion-exchange chromatography, making use of the fact that the stability constant for formation of EDTA complexes increases for log K ≈ 15.5 for [La(EDTA)]⁻ to log K ≈ 19.8 for [Lu(EDTA)]⁻.

Coordination chemistry and catalysis

When in the form of coordination complexes, lanthanides exist overwhelmingly in their +3 oxidation state, although particularly stable 4f configurations can also give +4 (Ce, Tb) or +2 (Eu, Yb) ions. All of these forms are strongly electropositive and thus lanthanide ions are hard Lewis acids. The oxidation states are also very stable and with the exception of SmI₂ and cerium(IV) salts lanthanides are not used for redox chemistry. 4f electrons have a high probability of being found close to the nucleus and are thus strongly affected as the nuclear charge increases across the series; this results in a corresponding decrease in ionic radii referred to as the lanthanide contraction.

The low probability of the 4f electrons existing at the outer region of the atom or ion permits little effective overlap between the orbitals of a lanthanide ion and any binding ligand. Thus lanthanide complexes typically have little or no covalent character and are not influenced by orbital geometries. The lack of orbital interaction also means that varying the metal typically has little effect on the complex (other than size), especially when compared to transition metals. Complexes are held together by weaker electrostatic forces which are omni-directional and thus the ligands alone dictate the symmetry and coordination of complexes. Steric factors therefore dominate, with coordinative saturation of the metal being balanced against inter-ligand repulsion. This results in a diverse range of coordination geometries, many of which are irregular, and also manifests itself in the highly fluxional nature of the complexes. As there is no energetic reason to be locked into a single geometry, rapid intramolecular and intermolecular ligand exchange will take place. This typically results in complexes that rapidly fluctuate between all possible configurations.

Many of these features make lanthanide complexes effective catalysts. Hard Lewis acids are able to polarise bonds upon coordination and thus alter the electrophilicity of compounds, with a classic example being the Luche reduction. The large size of the ions coupled with their labile ionic bonding allows even bulky coordinating species to bind and dissociate rapidly, resulting in very high turnover rates; thus excellent yields can often be achieved with loadings of only a few mol%. The lack of orbital interactions combined with the lanthanide contraction means that the lanthanides change in size across the series but that their chemistry remains much the same. This allows for easy tuning of the steric environments and examples exist where this has been used to improve the catalytic activity of the complex and change the nuclearity of metal clusters.

Despite this, the use of lanthanide coordination complexes as homogeneous catalysts is largely restricted to the laboratory and there are currently few examples them being used on an industrial scale. It should be noted however, that lanthanides exist in many forms other that coordination complexes and many of these are industrially useful. In particular lanthanide metal oxides are used as heterogeneous catalysts in various industrial processes.

Ln(III) compounds

The trivalent lanthanides mostly form ionic salts. The trivalent ions are hard acceptors and form more stable complexes with oxygen-donor ligands than with nitrogen-donor ligands. The larger ions are 9-coordinate in aqueous solution, [Ln(H₂O)₉]³⁺ but the smaller ions are 8-coordinate, [Ln(H₂O)₈]³⁺. There is some evidence that the later lanthanides have more water molecules in the second coordination sphere. Complexation with monodentate ligands is generally weak because it is difficult to displace water molecules from the first coordination sphere. Stronger complexes are formed with chelating ligands because of the chelate effect, such as the tetra-anion derived from 1,4,7,10-tetraazacyclododecane-1,4,7,10-tetraacetic acid (DOTA).

Ln(II) and Ln(IV) compounds

The most common divalent derivatives of the lanthanides are for Eu(II), which achieves a favorable f⁷ configuration. Divalent halide derivatives are known for all of the lanthanides. They are either conventional salts or are Ln(III) "electride"-like salts. The simple salts include YbI₂, EuI₂, and SmI₂. The electride-like salts, described as Ln³⁺, 2I⁻, e⁻, include LaI₂, CeI₂ and GdI₂. Many of the iodides form soluble complexes with ethers, e.g. TmI₂(dimethoxyethane)₃. Samarium(II) iodide is a useful reducing agent. Ln(II) complexes can be synthesized by transmetalation reactions.

Ce(IV) in ceric ammonium nitrate is a useful oxidizing agent. Otherwise tetravalent lanthanides are rare. The Ce(IV) is the exception owing to the tendency to form an unfilled f shell.

Hydrides

Chemical element	La	Ce	Pr	Nd	Pm	Sm	Eu	Gd	Tb	Dy	Ho	Er	Tm	Yb	Lu
Atomic number	57	58	59	60	61	62	63	64	65	66	67	68	69	70	71
Metal lattice (RT)	dhcp	fcc	dhcp	dhcp	dhcp	r	bcc	hcp	hcp	hcp	hcp	hcp	hcp	hcp	hcp
Dihydride	LaH2+x	CeH2+x	PrH2+x	NdH2+x		SmH2+x	EuH2 o "salt like"	GdH2+x	TbH2+x	DyH2+x	HoH2+x	ErH2+x	TmH2+x	YbH2+x o, fcc "salt like"	LuH2+x
Structure	CaF2	CaF2	CaF2	CaF2	CaF2	CaF2	*PbCl2	CaF2	CaF2	CaF2	CaF2	CaF2	CaF2		CaF2
metal sub lattice	fcc	fcc	fcc	fcc	fcc	fcc	o	fcc	fcc	fcc	fcc	fcc	fcc	o fcc	fcc
Trihydride	LaH3-x	CeH3-x	PrH3-x	NdH3-x		SmH3-x	EuH3-x	GdH3-x	TbH3-x	DyH3-x	HoH3-x	ErH3-x	TmH3-x		LuH3-x
metal sub lattice	fcc	fcc	fcc	hcp	hcp	hcp	fcc	hcp	hcp	hcp	hcp	hcp	hcp	hcp	hcp
Trihydride properties transparent insulators (color where recorded)	red	bronze to grey	PrH3-x fcc	NdH3-x hcp		golden greenish	EuH3-x fcc	GdH3-x hcp	TbH3-x hcp	DyH3-x hcp	HoH3-x hcp	ErH3-x hcp	TmH3-x hcp		LuH3-x hcp

Lanthanide metals react exothermically with hydrogen to form LnH₂, dihydrides. With the exception of Eu and Yb which resemble the Ba and Ca hydrides (non conducting,transparent salt like compounds) they form black pyrophoric, conducting compounds where the metal sub-lattice is face centred cubic and the H atoms occupy tetrahedral sites. Further hydrogenation produces a trihydride which is non-stoichiometric, non-conducting, more salt like. The formation of trihydride is associated with and increase in 8–10% volume and this is linked to greater localisation of charge on the hydrogen atoms which become more anionic (H⁻ hydride anion) in character.

Halides

Lanthanide halides

Chemical element	La	Ce	Pr	Nd	Pm	Sm	Eu	Gd	Tb	Dy	Ho	Er	Tm	Yb	Lu
Atomic number	57	58	59	60	61	62	63	64	65	66	67	68	69	70	71
Tetrafluoride		CeF4	PrF4						TbF4
Color m.p. °C		white dec	white dec						white dec
Structure C.N.		UF4 8	UF4 8						UF4 8
Trifluoride	LaF3	CeF3	PrF3	NdF3	PmF3	SmF3	EuF3	GdF3	TbF3	DyF3	HoF3	ErF3	TmF3	YbF3	LuF3
Color m.p. °C	white 1493	white 1430	green 1395	violet 1374	green 1399	white 1306	white 1276	white 1231	white 1172	green 1154	pink 1143	pink 1140	white 1158	white 1157	white 1182
Structure C.N.	LaF3 9	LaF3 9	LaF3 9	LaF3 9		YF3 8	YF3 8	YF3 8	YF3 8	YF3 8	YF3 8	YF3 8	YF3 8	YF3 8	YF3 8
Trichloride	LaCl3	CeCl3	PrCl3	NdCl3	PmCl3	SmCl3	EuCl3	GdCl3	TbCl3	DyCl3	HoCl3	ErCl3	TmCl3	YbCl3	LuCl3
Color m.p. °C	La	white 817	green 786	mauve 758	green 786	yellow 682	yellow dec	white 602	white 582	white 647	yellow 720	violet 776	yellow 824	white 865	white 925
Structure C.N.	UCl3 9	UCl3 9	UCl3 9	UCl3 9		UCl3 9	UCl3 9	UCl3 9	PuBr3 8	PuBr3 8	YCl3 6	YCl3 6	YCl3 6	YCl3 6	YCl3 6
Tribromide	LaBr3	CeBr3	PrBr3	NdBr3	PmF3	SmBr3	EuBr3	GdBr3	TbBr3	DyBr3	HoBr3	ErBr3	TmBr3	YbBr3	LuBr3
Color m.p. °C		white 733	green 691	violet 682	green 693	yellow 640	grey dec	white 770	white 828	white 879	yellow 919	violet 923	white 954	white dec	white 1025
Structure C.N.	UCl3 9	UCl3 9	UCl3 9	PuBr3 8		PuBr3 8	PuBr3 8	6	6	6	6	6	6	6	6
Triiodide	LaI3	CeI3	PrI3	NdI3	PmI3	SmI3	EuI3	GdI3	TbI3	DyI3	HoI3	ErI3	TmI3	YbI3	LuI3
Color m.p. °C		yellow 766	PrF4	green 784	green 737	orange 850	dec.	yellow 925	957	green 978	yellow 994	violet 1015	yellow 1021	white dec	brown 1050
Structure C.N.	PuBr3 8	PuBr3 8	PuBr3 8	PuBr3 8		BiI3 6	BiI3 6	BiI3 6	BiI3 6	BiI3 6	BiI3 6	BiI3 6	BiI3 6	BiI3 6	BiI3 6
Difluoride						SmF2	EuF2							YbF2
Color m.p. °C						purple 1417	yellow 1416							grey
Structure C.N.						CaF2 8	CaF2 8							CaF2 8
Dichloride				NdCl2		SmCl2	EuCl2			DyCl2			TmCl2	YbCl2
Color m.p. °C				green 841		brown 859	white 731			black dec.			green 718	green 720
Structure C.N.				PbCl2 9		PbCl2 9	PbCl2 9			SrBr2			SrI2 7	SrI2 7
Dibromide				NdBr2		SmBr2	EuBr2			DyBr2			TmBr2	YbBr2
Color m.p. °C				green 725		brown 669	white 731			black			green	yellow 673
Structure C.N.				PbCl2 9		SrBr2 8	SrBr2 8			SrI2 7			SrI2 7	SrI2 7
Diiodide	LaI2 metallic	CeI2 metallic	PrI2 metallic	NdI2 high pressure metallic		SmI2	EuI2	GdI2 metallic		DyI2			TmI2	YbI2
Color m.p. °C		bronze 808	bronze 758	violet 562		green 520	green 580	bronze 831		purple 721			black 756	yellow 780	Lu
Structure C.N.	CuTi2 8	CuTi2 8	CuTi2 8	SrBr2 8 CuTi2 8		EuI2 7	EuI2 7	2H-MoS2 6					CdI2 6	CdI2 6
Ln7I12	La7I12		Pr7I12	'					Tb7I12
Sesquichloride	La2Cl3							Gd2Cl3	Tb2Cl3				Tm2Cl3		Lu2Cl3
Structure								Gd2Cl3	Gd2Cl3
Sesquibromide								Gd2Br3	Tb2Br3
Structure								Gd2Cl3	Gd2Cl3
Monoiodide	LaI
Structure	NiAs type

The only tetrahalides known are those of cerium, praseodymium and terbium mirroring the formation of the dioxides. All of the lanthanides form trihalides with fluorine, chlorine, bromine and iodine. They are all high melting and predominantly ionic in nature. The fluorides are only slightly soluble in water and are not sensitive to air, and this contrasts with the other halides which are air sensitive, readily soluble in water and react at high temperature to form oxohalides. The trihalides were important as pure metal can be prepared from them. In the gas phase the trihalides are planar or approximately planar, the lighter lanthanides have a lower % of dimers, the heavier lanthanides a higher proportion. The dimers have a similar structure to Al₂Cl₆

Some of the dihalides are conducting while the rest are insulators. The conducting forms can be considered as Ln^III electride compounds where the electron is delocalised into a conduction band, Ln³⁺ (X⁻)₂(e⁻). All of the diodides have relatively short metal-metal separations. The CuTi₂ structure of the lanthanum, cerium and praseodymium diodides along with HP-NdI₂ contain 4⁴ nets of metal and iodine atoms with short metal-metal bonds (393-386 La-Pr). these compounds should be considered to be two-dimensional metals (two-dimensional in the same way that graphite is). The salt like dihalides include those of Eu Dy Tm and Yb. The formation of a relatively stable +2 oxidation state for Eu and Yb is usually explained by the stability (exchange energy) of half filled (f⁷) and fully filled f¹⁴. GdI₂ possesses the layered MoS₂ structure, is ferromagnetic and exhibits colossal magnetoresistance The sesquihalides Ln₂X₃ and the Ln₇I₁₂ compounds listed in the table contain metal clusters, discrete Ln₆I₁₂ clusters in Ln₇I₁₂ and condensed clusters forming chains in the sesquihalides. Scandium forms a similar cluster compound with chlorine, Sc₇Cl₁₂ Unlike many transition metal clusters these lanthanide clusters do not have strong metal-metal interactions and this is due to the low number of valence electrons involved, but instead are stabilised by the surrounding halogen atoms.

LaI is the only known monohalide. Prepared from the reaction of LaI₃ and La metal, it has a NiAs type structure and can be formulated La³⁺ (I⁻)(e⁻)₂.

Oxides and hydroxides

All of the lanthanides form sesquioxides, Ln₂O₃. The lighter/larger lanthanides adopt a hexagonal 7-coordinate structure while the heavier/smaller ones adopt a cubic 6-coordinate "C-M₂O₃" structure. All of the sesquioxides are basic, and absorb water and carbon dioxide from air to form carbonates, hydroxides and hydroxycarbonates. They dissolve in acids to form salts.

Cerium forms a stoichiometric dioxide, CeO₂, where cerium has an oxidation state of +4. CeO₂ is basic and dissolves with difficulty in acid to form Ce⁴⁺ solutions, from which Ce^IV salts can be isolated, for example the hydrated nitrate Ce(NO₃)₄.5H₂O. CeO₂ is used as an oxidation catalyst in catalytic converters. Praseodymium and terbium form non-stoichiometric oxides containing Ln^IV, although more extreme reaction conditions can produce stoichiometric (or near stoichiometric) PrO₂ and TbO₂.

Europium and ytterbium form salt-like monoxides, EuO and YbO, which have a rock salt structure. EuO is ferromagnetic at low temperatures, and is a semiconductor with possible applications in spintronics. A mixed Eu^II/Eu^III oxide Eu₃O₄ can be produced by reducing Eu₂O₃ in a stream of hydrogen. Neodymium and samarium also form monoxides, but these are shiny conducting solids, although the existence of samarium monoxide is considered dubious.

All of the lanthanides form hydroxides, Ln(OH)₃. With the exception of lutetium hydroxide, which has a cubic structure, they have the hexagonal UCl₃ structure. The hydroxides can be precipitated from solutions of Ln^III. They can also be formed by the reaction of the sesquioxide, Ln₂O₃, with water, but although this reaction is thermodynamically favourable it is kinetically slow for the heavier members of the series. Fajan's rules indicate that the smaller Ln³⁺ ions will be more polarizing and their salts correspondingly less ionic. The hydroxides of the heavier lanthanides become less basic, for example Yb(OH)₃ and Lu(OH)₃ are still basic hydroxides but will dissolve in hot concentrated NaOH.

Chalcogenides (S, Se, Te)

All of the lanthanides form Ln₂Q₃ (Q= S, Se, Te). The sesquisulfides can be produced by reaction of the elements or (with the exception of Eu₂S₃) sulfidizing the oxide (Ln₂O₃) with H₂S. The sesquisulfides, Ln₂S₃ generally lose sulfur when heated and can form a range of compositions between Ln₂S₃ and Ln₃S₄. The sesquisulfides are insulators but some of the Ln₃S₄ are metallic conductors (e.g. Ce₃S₄) formulated (Ln³⁺)₃ (S²⁻)₄ (e⁻), while others (e.g. Eu₃S₄ and Sm₃S₄) are semiconductors. Structurally the sesquisulfides adopt structures that vary according to the size of the Ln metal. The lighter and larger lanthanides favouring 7 coordinate metal atoms, the heaviest and smallest lanthanides (Yb and Lu) favouring 6 coordination and the rest structures with a mixture of 6 and 7 coordination. Polymorphism is common amongst the sesquisulfides. The colors of the sesquisulfides vary metal to metal and depend on the polymorphic form. The colors of the γ-sesquisulfides are La₂S₃, white/yellow; Ce₂S₃, dark red; Pr₂S₃, green; Nd₂S₃, light green; Gd₂S₃, sand; Tb₂S₃, light yellow and Dy₂S₃, orange. The shade of γ-Ce₂S₃ can be varied by doping with Na or Ca with hues ranging from dark red to yellow, and Ce₂S₃ based pigments are used commercially and are seen as low toxicity substitutes for cadmium based pigments.

All of the lanthanides form monochalcogenides, LnQ, (Q= S, Se, Te). The majority of the monochalcogenides are conducting, indicating a formulation Ln^IIIQ²⁻(e-) where the electron is in conduction bands. The exceptions are SmQ, EuQ and YbQ which are semiconductors or insulators but exhibit a pressure induced transition to a conducting state. Compounds LnQ₂ are known but these do not contain Ln^IV but are Ln^III compounds containing polychalcogenide anions.

Oxysulfides Ln₂O₂S are well known, they all have the same structure with 7 coordinate Ln atoms with 3 sulfur atoms and 4 oxygen as near neighbours. Doping these with other lanthanide elements produces phosphors. As an example, gadolinium oxysulfide, Gd₂O₂S doped with Tb³⁺ produces visible photons when irradiated with high energy X-rays and is used as a scintillator in flat panel detectors. When mischmetal, an alloy of lanthanide metals, is added to molten steel to remove oxygen and sulfur, stable oxysulfides are produced that form an immiscible solid.

Pnictides (group 15)

All of the lanthanides form a mononitride, LnN, with the rock salt structure. The mononitrides have attracted interest because of their unusual physical properties. SmN and EuN are reported as being "half metals". NdN, GdN, TbN and DyN are ferromagnetic, SmN is antiferromagnetic. Applications in the field of spintronics are being investigated. CeN is unusual as it is a metallic conductor, contrasting with the other nitrides also with the other cerium pnictides. A simple description is Ce⁴⁺ N³⁻ (e–) but the interatomic distances are a better match for the trivalent state rather than for the tetravalent state. A number of different explanations have been offered. The nitrides can be prepared by the reaction of lanthanum metals with nitrogen. Some nitride is produced along with the oxide, when lanthanum metals are ignited in air. Alternative methods of synthesis are a high temperature reaction of lanthanide metals with ammonia or the decomposition of lanthanide amides, Ln(NH₂)₃. Achieving pure stoichiometric compounds, and crystals with low defect density has proved difficult. The lanthanide nitrides are sensitive to air and hydrolyse producing ammonia.

The other pnictides phosphorus, arsenic, antimony and bismuth also react with the lanthanide metals to form monopnictides, LnQ. Additionally a range of other compounds can be produced with varying stoichiometries, such as LnP₂, LnP₅, LnP₇ Ln₃As, Ln₅As₃ and LnAs₂.

Carbides

Carbides of varying stoichiometries are known for the lanthanides. Non-stoichiometry is common. All of the lanthanides form LnC₂ and Ln₂C₃ which both contain C₂ units. The dicarbides with exception of EuC₂, are metallic conductors with the calcium carbide structure and can be formulated as Ln³⁺C₂²⁻(e–). The C-C bond length is longer than that in CaC₂, which contains the C₂²⁻ anion, indicating that the antibonding orbitals of the C₂²⁻ anion are involved in the conduction band. These dicarbides hydrolyse to form hydrogen and a mixture of hydrocarbons. EuC₂ and to a lesser extent YbC₂ hydrolyse differently producing a higher percentage of acetylene (ethyne). The sesquicarbides, Ln₂C₃ can be formulated as Ln₄(C₂)₃. These compounds adopt the Pu₂C₃ structure which has been described as having C₂²⁻ anions in bisphenoid holes formed by eight near Ln neighbours. The lengthening of the C-C bond is less marked in the sesquicarbides than in the dicarbides, with the exception of Ce₂C₃. Other carbon rich stoichiometries are known for some lanthanides. Ln₃C₄ (Ho-Lu) containing C, C₂ and C₃ units; Ln₄C₇ (Ho- Lu) contain C atoms and C₃ units and Ln₄C₅ (Gd-Ho) containing C and C₂ units. Metal rich carbides contain interstitial C atoms and no C₂ or C₃ units. These are Ln₄C₃ (Tb and Lu); Ln₂C (Dy, Ho, Tm) and Ln₃C (Sm-Lu).

Borides

All of the lanthanides form a number of borides. The "higher" borides (LnB_x where x > 12) are insulators/semiconductors whereas the lower borides are typically conducting. The lower borides have stoichiometries of LnB₂, LnB₄, LnB₆ and LnB₁₂. Applications in the field of spintronics are being investigated. The range of borides formed by the lanthanides can be compared to those formed by the transition metals. The boron rich borides are typical of the lanthanides (and groups 1–3) whereas for the transition metals tend to form metal rich, "lower" borides. The lanthanide borides are typically grouped together with the group 3 metals with which they share many similarities of reactivity, stoichiometry and structure. Collectively these are then termed the rare earth borides.

Many methods of producing lanthanide borides have been used, amongst them are direct reaction of the elements; the reduction of Ln₂O₃ with boron; reduction of boron oxide, B₂O₃, and Ln₂O₃ together with carbon; reduction of metal oxide with boron carbide, B₄C. Producing high purity samples has proved to be difficult. Single crystals of the higher borides have been grown in a low melting metal (e.g. Sn, Cu, Al).

Diborides, LnB₂, have been reported for Sm, Gd, Tb, Dy, Ho, Er, Tm, Yb and Lu. All have the same, AlB₂, structure containing a graphitic layer of boron atoms. Low temperature ferromagnetic transitions for Tb, Dy, Ho and Er. TmB₂ is ferromagnetic at 7.2 K.

Tetraborides, LnB₄ have been reported for all of the lanthanides except EuB₄, all have the same UB₄ structure. The structure has a boron sub-lattice consists of chains of octahedral B₆ clusters linked by boron atoms. The unit cell decreases in size successively from LaB₄ to LuB₄. The tetraborides of the lighter lanthanides melt with decomposition to LnB₆. Attempts to make EuB₄ have failed. The LnB₄ are good conductors and typically antiferromagnetic.

Hexaborides, LnB₆ have been reported for all of the lanthanides. They all have the CaB₆ structure, containing B₆ clusters. They are non-stoichiometric due to cation defects. The hexaborides of the lighter lanthanides (La – Sm) melt without decomposition, EuB₆ decomposes to boron and metal and the heavier lanthanides decompose to LnB₄ with exception of YbB₆ which decomposes forming YbB₁₂. The stability has in part been correlated to differences in volatility between the lanthanide metals. In EuB₆ and YbB₆ the metals have an oxidation state of +2 whereas in the rest of the lanthanide hexaborides it is +3. This rationalises the differences in conductivity, the extra electrons in the Ln^III hexaborides entering conduction bands. EuB₆ is a semiconductor and the rest are good conductors. LaB₆ and CeB₆ are thermionic emitters, used, for example, in scanning electron microscopes.

Dodecaborides, LnB₁₂, are formed by the heavier smaller lanthanides, but not by the lighter larger metals, La – Eu. With the exception YbB₁₂ (where Yb takes an intermediate valence and is a Kondo insulator), the dodecaborides are all metallic compounds. They all have the UB₁₂ structure containing a 3 dimensional framework of cubooctahedral B₁₂ clusters.

The higher boride LnB₆₆ is known for all lanthanide metals. The composition is approximate as the compounds are non-stoichiometric. They all have similar complex structure with over 1600 atoms in the unit cell. The boron cubic sub lattice contains super icosahedra made up of a central B₁₂ icosahedra surrounded by 12 others, B₁₂(B₁₂)₁₂. Other complex higher borides LnB₅₀ (Tb, Dy, Ho Er Tm Lu) and LnB₂₅ are known (Gd, Tb, Dy, Ho, Er) and these contain boron icosahedra in the boron framework.

Organometallic compounds

Lanthanide-carbon σ bonds are well known; however as the 4f electrons have a low probability of existing at the outer region of the atom there is little effective orbital overlap, resulting in bonds with significant ionic character. As such organo-lanthanide compounds exhibit carbanion-like behaviour, unlike the behaviour in transition metal organometallic compounds. Because of their large size, lanthanides tend to form more stable organometallic derivatives with bulky ligands to give compounds such as Ln[CH(SiMe₃)₃]. Analogues of uranocene are derived from dilithiocyclooctatetraene, Li₂C₈H₈. Organic lanthanide(II) compounds are also known, such as Cp*₂Eu.

Physical properties

Magnetic and spectroscopic

All the trivalent lanthanide ions, except lanthanum and lutetium, have unpaired f electrons. However, the magnetic moments deviate considerably from the spin-only values because of strong spin-orbit coupling. The maximum number of unpaired electrons is 7, in Gd³⁺, with a magnetic moment of 7.94 B.M., but the largest magnetic moments, at 10.4–10.7 B.M., are exhibited by Dy³⁺ and Ho³⁺. However, in Gd³⁺ all the electrons have parallel spin and this property is important for the use of gadolinium complexes as contrast reagent in MRI scans.

A solution of 4% holmium oxide in 10% perchloric acid, permanently fused into a quartz cuvette as a wavelength calibration standard

Crystal field splitting is rather small for the lanthanide ions and is less important than spin-orbit coupling in regard to energy levels. Transitions of electrons between f orbitals are forbidden by the Laporte rule. Furthermore, because of the "buried" nature of the f orbitals, coupling with molecular vibrations is weak. Consequently, the spectra of lanthanide ions are rather weak and the absorption bands are similarly narrow. Glass containing holmium oxide and holmium oxide solutions (usually in perchloric acid) have sharp optical absorption peaks in the spectral range 200–900 nm and can be used as a wavelength calibration standard for optical spectrophotometers, and are available commercially.

As f-f transitions are Laporte-forbidden, once an electron has been excited, decay to the ground state will be slow. This makes them suitable for use in lasers as it makes the population inversion easy to achieve. The Nd:YAG laser is one that is widely used. Europium-doped yttrium vanadate was the first red phosphor to enable the development of color television screens. Lanthanide ions have notable luminescent properties due to their unique 4f orbitals. Laporte forbidden f-f transitions can be activated by excitation of a bound "antenna" ligand. This leads to sharp emission bands throughout the visible, NIR, and IR and relatively long luminescence lifetimes.

Occurrence

Main article: Rare earth element § Geological distribution

The lanthanide contraction is responsible for the great geochemical divide that splits the lanthanides into light and heavy-lanthanide enriched minerals, the latter being almost inevitably associated with and dominated by yttrium. This divide is reflected in the first two "rare earths" that were discovered: yttria (1794) and ceria (1803). The geochemical divide has put more of the light lanthanides in the Earth's crust, but more of the heavy members in the Earth's mantle. The result is that although large rich ore-bodies are found that are enriched in the light lanthanides, correspondingly large ore-bodies for the heavy members are few. The principal ores are monazite and bastnäsite. Monazite sands usually contain all the lanthanide elements, but the heavier elements are lacking in bastnäsite. The lanthanides obey the Oddo-Harkins rule – odd-numbered elements are less abundant than their even-numbered neighbors.

Three of the lanthanide elements have radioactive isotopes with long half-lives (¹³⁸La, ¹⁴⁷Sm and ¹⁷⁶Lu) that can be used to date minerals and rocks from Earth, the Moon and meteorites.

Applications

Industrial

Lanthanide elements and their compounds have many uses but the quantities consumed are relatively small in comparison to other elements. About 15000 ton/year of the lanthanides are consumed as catalysts and in the production of glasses. This 15000 tons corresponds to about 85% of the lanthanide production. From the perspective of value, however, applications in phosphors and magnets are more important.

The devices lanthanide elements are used in include superconductors, samarium-cobalt and neodymium-iron-boron high-flux rare-earth magnets, magnesium alloys, electronic polishers, refining catalysts and hybrid car components (primarily batteries and magnets). Lanthanide ions are used as the active ions in luminescent materials used in optoelectronics applications, most notably the Nd:YAG laser. Erbium-doped fiber amplifiers are significant devices in optical-fiber communication systems. Phosphors with lanthanide dopants are also widely used in cathode ray tube technology such as television sets. The earliest color television CRTs had a poor-quality red; europium as a phosphor dopant made good red phosphors possible. Yttrium iron garnet (YIG) spheres can act as tunable microwave resonators. Lanthanide oxides are mixed with tungsten to improve their high temperature properties for welding, replacing thorium, which was mildly hazardous to work with. Many defense-related products also use lanthanide elements such as night vision goggles and rangefinders. The SPY-1 radar used in some Aegis equipped warships, and the hybrid propulsion system of Arleigh Burke-class destroyers all use rare earth magnets in critical capacities. The price for lanthanum oxide used in fluid catalytic cracking has risen from $5 per kilogram in early 2010 to $140 per kilogram in June 2011.

Most lanthanides are widely used in lasers, and as (co-)dopants in doped-fiber optical amplifiers; for example, in Er-doped fiber amplifiers, which are used as repeaters in the terrestrial and submarine fiber-optic transmission links that carry internet traffic. These elements deflect ultraviolet and infrared radiation and are commonly used in the production of sunglass lenses. Other applications are summarized in the following table:

Application	Percentage
Catalytic converters	45
Petroleum refining catalysts	25
Permanent magnets	12
Glass polishing and ceramics	7
Metallurgical	7
Phosphors	3
Other	1

The complex Gd(DOTA) is used in magnetic resonance imaging.

Life science

As mentioned in the industrial applications section above, lanthanide metals are particularly useful in technologies that take advantage of their reactivity to specific wavelengths of light. Certain life science applications take advantage of the unique fluorescence properties of lanthanide ion complexes (Ln(III) chelates or cryptates). These are well-suited for this application due to their large Stokes shifts and extremely long emission lifetimes (from microseconds to milliseconds) compared to more traditional fluorophores (e.g., fluorescein, allophycocyanin, phycoerythrin, and rhodamine). The biological fluids or serum commonly used in these research applications contain many compounds and proteins which are naturally fluorescent. Therefore, the use of conventional, steady-state fluorescence measurement presents serious limitations in assay sensitivity. Long-lived fluorophores, such as lanthanides, combined with time-resolved detection (a delay between excitation and emission detection) minimizes prompt fluorescence interference.

Time-resolved fluorometry (TRF) combined with fluorescence resonance energy transfer (FRET) offers a powerful tool for drug discovery researchers: Time-Resolved Fluorescence Resonance Energy Transfer or TR-FRET. TR-FRET combines the low background aspect of TRF with the homogeneous assay format of FRET. The resulting assay provides an increase in flexibility, reliability and sensitivity in addition to higher throughput and fewer false positive/false negative results.

This method involves two fluorophores: a donor and an acceptor. Excitation of the donor fluorophore (in this case, the lanthanide ion complex) by an energy source (e.g. flash lamp or laser) produces an energy transfer to the acceptor fluorophore if they are within a given proximity to each other (known as the Förster’s radius). The acceptor fluorophore in turn emits light at its characteristic wavelength.

The two most commonly used lanthanides in life science assays are shown below along with their corresponding acceptor dye as well as their excitation and emission wavelengths and resultant Stokes shift (separation of excitation and emission wavelengths).

Donor	Excitation⇒Emission λ (nm)	Acceptor	Excitation⇒Emission λ (nm)	Stoke's Shift (nm)
Eu3+	340⇒615	Allophycocyanin	615⇒660	320
Tb3+	340⇒545	Phycoerythrin	545⇒575	235

Upcoming Medical Uses

Currently there is research showing that lanthanides elements can be used as anticancer agents. The main role of the lanthanides in these studies is to inhibit proliferation of the cancer cells. Specifically cerium and lanthanum have been studied for their role as anti-cancer agents.

One of the specific elements from the lanthanide group that has been tested and used is cerium (Ce). There have been studies that use a protein-cerium complex to observe the effect of cerium on the cancer cells. The hope was to inhibit cell proliferation and promote cytotoxicity. Transferrin receptors in cancer cells, such as those in breast cancer cells and epithelial cervical cells, promote the cell proliferation and malignancy of the cancer. Transferrin is a protein used to transport iron into the cells and is needed to aid the cancer cells in DNA replication. Transferrin acts as a growth factor for the cancerous cells and is dependent on iron. Cancer cells have much higher levels of transferrin receptors than normal cells and are very dependent on iron for their proliferation. Cerium has shown results as an anti-cancer agent due to its similarities in structure and biochemistry to iron. Cerium may bind in the place of iron on to the transferrin and then be brought into the cancer cells by transferrin-receptor mediated endocytosis. The cerium binding to the transferrin in place of the iron inhibits the transferrin activity in the cell. This creates a toxic environment for the cancer cells and causes a decrease in cell growth. This is the proposed mechanism for cerium’s effect on cancer cells, though the real mechanism may be more complex in how cerium inhibits cancer cell proliferation. Specifically in HeLa cancer cells studied in vitro, cell viability was decreased after 48 to 72 hours of cerium treatments. Cells treated with just cerium had decreases in cell viability, but cells treated with both cerium and transferrin had more significant inhibition for cellular activity.

Another specific element that has be tested and used as an anti-cancer agent is lanthanum, more specifically lanthanum chloride (LaCl3 ). The lanthanum ion is used to affect the levels of let-7a and microRNAs miR-34a in a cell throughout the cell cycle. When the lanthanum ion was introduced to the cell in vivo or in vitro, it inhibited the rapid growth and induced apoptosis of the cancer cells (specifically cervical cancer cells). This effect was caused by the regulation of the let-7a and microRNAs by the lanthanum ions. The mechanism for this effect is still unclear but it is possible that the lanthanum is acting in a similar way as the cerium and binding to a ligand necessary for cancer cell proliferation.

Biological effects

Due to their sparse distribution in the earth's crust and low aqueous solubility, the lanthanides have a low availability in the biosphere, and are not known to naturally form part of any biological molecules. Compared to most other nondietary elements, non-radioactive lanthanides are classified as having low toxicity.

See also

• Actinide
• Group 3 element
• Lanthanide contraction
• Rare earth element
• Lanthanide probes

External links

• lanthanide Sparkle Model, used in the computational chemistry of lanthanide complexes
• USGS Rare Earths Statistics and Information
• Ana de Bettencourt-Dias: Chemistry of the lanthanides and lanthanide-containing materials