Lanthanides: lanthanides periodic table facts

What are the lanthanides? the lanthanide series of metals

A lanthanide is any one of the elements from lanthanum (La, element 57) through lutetium (Lu, element 71). Textbooks use the informal symbol Ln to stand for any member of the group, so a formula such as Ln(III) means "any lanthanide in the +3 state." All 15 sit in period 6, and IUPAC treats the lanthanides as one collective set because they are so alike: every lanthanide forms a trivalent ion, and what that ion does is governed mostly by its size rather than by its outer shell. From hexachloroplatinic acid to chemotherapy agents, reactions in medicine and industry are listed in platinum uses along with notes on its chemical behaviour.

Calling them the lanthanide series of metals also says something practical. Lanthanides are silvery, reactive solids that tarnish in air, and in trade and geology each one is simply sold as a rare-earth metal. Because the atomic number rises by one at each step while the chemistry barely moves, the table below is the quickest way to see the whole family at once, with the ion configuration and size that drive most of its behavior.

Lanthanides and actinides in the f-block of the periodic table

Most printed versions of the periodic table pull the lanthanides and actinides out as two rows below the main grid. That is a layout choice, not a chemical one: dropped back into place, the f-block elements would stretch the table to 32 columns, which is too wide for most pages. In that wide form lanthanum follows barium (56), the row it starts runs straight on to hafnium (72), and the row underneath fills the 5f subshell in the same way that these 4f elements fill theirs.

Lutetium is the awkward member. Its 4f shell is already full, so its last electron enters a 5d orbital, which strictly makes it a d-block element and one of the transition metals.

For that reason some tables put lutetium, not lanthanum, under scandium and yttrium in group 3, and an IUPAC project has been reviewing the question. Whichever layout your textbook uses, it behaves like the other fourteen.

Lanthanoids or lanthanides: where the name comes from

The word comes from the Greek lanthanein, "to lie hidden." Lanthanum was first found in 1839 hiding inside a cerium mineral, and Victor Goldschmidt coined "lanthanide" in 1925 for the whole family of elements that resemble it. Because the name literally means "like lanthanum," purists argue that lanthanum itself should not count; IUPAC includes it anyway, on the basis of common usage. Alongside its discovery near Ytterby, the ytterbium metal entry explains the work of Jean Charles Galissard de Marignac.

IUPAC's current recommendation is lanthanoid, because the suffix "-ide" normally marks a negative ion (as in chloride), while "-oid" signals a likeness. Journals and textbooks still overwhelmingly say lanthanides, so you will meet both forms, and IUPAC's own documents accept the older one.

Electron configuration and the lanthanide contraction

Almost everything distinctive about the lanthanides traces back to electron placement. Each step from La to Lu adds one proton to the nucleus and one more negative charge, but that newcomer mostly lands in a deep inner subshell instead of the outermost electron shell (6s), which stays the same for the whole row. The outside of the atom barely changes, while the inside quietly fills up.

How the 4f orbitals fill across the series

Beyond a xenon core, the ground-state electron configuration of most gas-phase lanthanide atoms follows one simple pattern:

$$[\text{Xe}]\,4f^{n}\,6s^{2}, \qquad n = Z - 56$$

So neodymium, with \(Z = 60\), is \([\text{Xe}]\,4f^{4}\,6s^{2}\). La, Ce, Gd and Lu break the rule by putting one electron in 5d instead, which gives gadolinium a half-filled shell of seven 4f electrons.

These seven orbitals sit inside the filled 5s and 5p shells, so they hardly take part in bonding. Up to seven of them can hold unpaired electrons, and that is the source of the very large magnetic moments measured for lanthanide compounds.

Ionic radius from La to Lu

The lanthanide contraction is the steady shrink in the ionic radius of the +3 ion across the series: from 103 pm for La to 86.1 pm for Lu in the table above, a drop of about 16%. The 4f shell shields the growing nuclear charge poorly, so each added proton pulls the outer shells in a little tighter.

Two practical effects follow. Neighboring lanthanide ions differ in size by only about 1 pm, which is why they are so hard to pull apart.

And by holmium the Ln ion is almost exactly the size of yttrium's, so it travels with the heavy lanthanides in nature. The same squeeze carries past the series, which is why hafnium ends up nearly the same size as zirconium.

Physical and magnetic properties of lanthanides

The properties of lanthanides change gradually along the row rather than abruptly, which makes the exceptions easy to spot. Among the lanthanides, Eu and Yb keep breaking the trends for the same reason: in the solid they contribute only two bonding electrons each instead of three.

Density, melting point and hardness

From the first member to the last, values climb from about 6.2 to 9.8 g/cm³ as the atoms shrink and get heavier. Eu and Yb dip well below their neighbors, with the largest metallic radius in the row (208.4 pm for europium), because each atom puts only two into the conduction band.

Cerium melts lowest of all the lanthanides, which is linked to unusually strong mixing of its 6s, 5d and 4f levels, and the values then rise toward the heavy end. Hardness rises across the row as well, but all of them are soft enough to cut with a steel blade.

Magnetism, color and spectra of the ions

Apart from La, Yb and Lu, which have no unpaired f electrons in the metal, the lanthanides are strongly paramagnetic, and some show the highest magnetic susceptibility of any element. Gadolinium becomes ferromagnetic below a Curie point of about 20 °C, close to room temperature, and dysprosium and holmium order magnetically only at much lower temperatures.

Because the 4f orbitals are tucked inside the atom, f–f transitions are weak and very narrow. In aqueous solution the colors are pale: the +3 ions of La, Ce, Gd and Lu are colorless, Pr is green, Nd is violet and Er is rose pink.

Each one also has a sharp-lined absorption spectrum that spectroscopy can use as a fingerprint. Those same narrow lines are why lanthanide ions emit such pure light in lasers and lighting.

Lanthanide chemistry: oxidation states and bonding

Lanthanide chemistry is far more uniform than the chemistry of the transition elements next door. Lanthanides are highly reactive, strongly electropositive metals: they burn in air, react slowly with cold water and quickly with hot water, giving off hydrogen. Their compounds are mostly ionic, and the same few patterns repeat from one end of the row to the other.

Why the +3 oxidation state dominates

Every member shows the +3 oxidation state, and for most it is the only one stable in water. Stripping the two 6s electrons plus one more costs relatively little: the sum of the first three ionization energy values runs from 3455 kJ/mol for lanthanum to 4186 kJ/mol for Yb, well below the 5139 kJ/mol aluminium needs.

The reduction potentials for the Ln(III)/Ln couple all fall between about −1.99 V (Eu) and −2.35 V (Pr), close to magnesium's −2.36 V, so the metals are reducing agents about as strong as the alkaline earth metals. Once the ion has formed, its buried 4f shell changes very little from element to element, so size is the main thing that varies.

Exceptions: Ce(IV), Eu(II) and Yb(II)

A few members step out of +3 when doing so reaches an especially stable configuration. Cerium gives up a fourth electron to reach an empty 4f shell like xenon: cerium metal burns to CeO₂ rather than Ce₂O₃, and ceric ammonium nitrate is a standard oxidizing agent in organic labs. Praseodymium and terbium can also reach +4, the latter because the result is 4f⁷.

Going the other way, Eu(II) (4f⁷) and Yb(II) (a full 4f¹⁴) are the common +2 ions, and samarium(II) iodide, SmI₂, is a popular reducing agent. Europium and ytterbium even dissolve in liquid ammonia like barium, and their simple salts, such as EuI₂, behave much like alkaline-earth salts. Outside Ce(IV) and Eu(II), no other oxidation state is stable in water.

Coordination chemistry of trivalent cations

In solution and in complexes, the trivalent cations act as hard Lewis acids. They bind oxygen donors such as water, carboxylates and phosphates more strongly than nitrogen donors, and the bonding is almost entirely electrostatic, with little covalent character.

Because the 4f shell barely overlaps with anything, the ligands rather than the metal decide the shape. Coordination numbers are high and flexible: the larger early cations hold nine water molecules, [Ln(H₂O)₉]³⁺, and the smaller late ones hold eight.

Lanthanide complexes also swap ligands very quickly. Chelating ligands bind far more tightly than single-point ones; the stability constant for EDTA rises from log K ≈ 15.5 for La to about 19.8 for Lu, and that steady climb, a direct result of the contraction, is what made the ion-exchange separations described below possible. In practice this coordination chemistry is also why Ln ions are such useful catalysts and tags: chemists can tune a complex simply by picking a larger or smaller ion.

Lanthanide compounds: oxides, halides, hydrides and borides

Because every member prefers +3, the same families of compounds turn up across the whole row, with only small shifts in structure as the ions shrink:

Oxides and hydroxides

All form basic sesquioxides, Ln₂O₃, which slowly absorb water and carbon dioxide from air. Cerium's stable dioxide, ceria (CeO₂), is the main exception. The hydroxides, Ln(OH)₃, precipitate from solution and get less basic toward the heavy end.

Halides

Every member forms trihalides with fluorine, chlorine, bromine and iodine. They are high-melting and ionic; the fluorides barely dissolve in water, while the others are air-sensitive and soluble. Molten trihalides are the usual starting point for making pure Ln metal.

Hydrogen compounds

The metals react with hydrogen to form black, conducting dihydrides, LnH₂, and then salt-like trihydrides that expand the solid by up to 10%.

Borides, carbides and nitrides

The hexaborides LaB₆ and CeB₆ are electron emitters in electron microscopes. Dicarbides, LnC₂, react with water to release hydrocarbons, and the mononitrides, LnN, have a rock-salt crystal structure that is studied for spintronics.

How the rare earths were discovered and separated

For more than a century, chemists struggled to tell the rare earths apart. The near-identical size that makes their chemistry so uniform made every new element hard to isolate and even harder to prove.

From Ytterby to Moseley: counting the elements

The story starts in 1787 with a heavy black mineral from a quarry at Ytterby, near Stockholm, later named gadolinite after the Finnish chemist Johan Gadolin. Cerium was isolated from another Swedish ore in 1803. Between 1839 and 1843 Carl Mosander split lanthanum, a so-called didymium and two more "earths" out of those samples, and the little village eventually gave its name to four elements: yttrium, ytterbium, terbium and erbium.

Didymium was not an element at all; in 1885 Carl Auer von Welsbach separated it into praseodymium and neodymium. Glass tinted with that mixture is still used in safety goggles for glassblowers, because it filters out the bright yellow glare of hot sodium.

New spectral methods helped identify more members, but nobody knew how many to expect, and Mendeleev never found a satisfying place for them in his table. The count was settled in 1913, when Henry Moseley showed that X-ray frequencies track atomic number: there had to be exactly 15 members, with element 61 still missing.

Separation of lanthanides by solvent extraction

Early separation of lanthanides relied on fractional crystallization: dissolve, crystallize, repeat, and let a tiny difference in solubility add up. In 1911 Charles James reported about 15,000 recrystallisations to purify a single thulium salt. The wartime Manhattan Project brought ion-exchange chromatography, in which complexing agents wash each element off a resin column one by one, in order of size.

Today the industrial method is solvent extraction. A nitrate solution of the mixed lanthanides is contacted with kerosene containing tributyl phosphate, and hundreds of countercurrent stages each nudge the ratio a little further, until the elements come out one at a time. It is still one of the most laborious processes in inorganic chemistry, and it is a large part of why separated material costs what it does.

Where lanthanide metals occur in nature

Despite the label, the rare earth elements are not especially rare. Cerium is about as common in Earth's crust as copper, and even lutetium, one of the scarcest stable members, is roughly 200 times more abundant than gold.

What is rare is a concentrated deposit: lanthanide metals almost always occur mixed together in minerals such as bastnäsite and monazite, or in ion-adsorption clays that are richer in the heavy members. Their abundance also zig-zags, since even-numbered elements are more common than their odd-numbered neighbors.

The one true exception is promethium. It has no stable isotopes; its longest-lived form, Pm-145, has a half-life of 17.7 years, so this radioactive element exists naturally only as traces made by the nuclear fission of uranium in ore.

It was finally confirmed in 1945 among fission products at Oak Ridge. Most of the world's rare earth metals are now mined and refined in China.

Uses of lanthanides in technology and medicine

Broadly, the uses of lanthanides fall into three groups (magnetism, light and chemistry), and each one exploits a property covered above.

Magnets, phosphors and catalysis

Neodymium-iron-boron (Nd₂Fe₁₄B) permanent magnets are the strongest commercial type, used in electric-vehicle motors, wind turbines, hard drives and the speaker in every smartphone; a little dysprosium keeps them working at high temperature. Samarium-cobalt alloys cover even hotter jobs. Europium and terbium ions supply the red and green phosphors in screens and energy-saving lamps, and erbium-doped fiber boosts the light signals that carry internet traffic.

Catalysis accounts for roughly a tenth of the economic value of the group. Lanthanum-loaded zeolites crack crude oil in refineries, and cerium oxide stores and releases oxygen in catalytic converters.

Mischmetal, an alloy of mostly cerium and lanthanum, strips oxygen and sulfur from molten steel and makes lighter flints. Other industrial uses include polishing glass and lenses with cerium oxide and oxysulfide screens that serve as an X-ray scintillator.

Life science and medical research

In medicine and life science, these elements earn their place through the same 4f electrons:

• Gadolinium chelates are the most widely used MRI contrast agents; its seven unpaired electrons speed up the relaxation of nearby water protons, which brightens the image.
• Lanthanum carbonate is prescribed as a phosphate binder for people with kidney failure.
• Lutetium-177 compounds deliver targeted radiation to some prostate and neuroendocrine tumors.
• Long-lived europium luminescence is used as a label in medical research assays.

Biology was long thought to ignore the group entirely, until microbiologists found in 2011 that some methanol-eating bacteria build lanthanum and cerium into a key enzyme, a finding that is now reshaping how biologists think about these metals.

Biological effects and toxicology of the lanthanide elements

The lanthanide elements are generally of low to moderate toxicity. Their salts are poorly absorbed from the gut, and most of what does enter the blood ends up in the liver and bone. Because a +3 ion is close in size to the calcium ion, Ln ions can block calcium channels and interfere with calcium-dependent enzymes, which is the main mechanism behind their biological effects.

Toxicology references usually treat the group together. The best-documented hazards are lung disease in workers who breathe rare earths dust over many years, and nephrogenic systemic fibrosis, a rare condition linked to free gadolinium released from older, less stable MRI agents in patients with severe kidney disease. That second case is the reason modern contrast agents use the most tightly binding chelates available.