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Cerium,  ₅₈Ce

General properties
Pronunciation	/ˈsɪəriəm/ ​(SEER-ee-əm)
Appearance	silvery white
Standard atomic weight.mw-parser-output .nobold{font-weight:normal} (Ar, standard)	7002140116000000000♠140.116(1)[1]
Cerium in the periodic table
Hydrogen Helium Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson – ↑ Ce ↓ Th lanthanum ← cerium → praseodymium
Atomic number (Z)	58
Group	group n/a
Period	period 6
Block	f-block
Element category	lanthanide
Electron configuration	[Xe] 4f1 5d1 6s2[2]
Electrons per shell	2, 8, 18, 19, 9, 2
Physical properties
Phase at STP	solid
Melting point	1068 K ​(795 °C, ​1463 °F)
Boiling point	3716 K ​(3443 °C, ​6229 °F)
Density (near r.t.)	6.770 g/cm3
when liquid (at m.p.)	6.55 g/cm3
Heat of fusion	5.46 kJ/mol
Heat of vaporization	398 kJ/mol
Molar heat capacity	26.94 J/(mol·K)
Vapor pressure P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 1992 2194 2442 2754 3159 3705
Atomic properties
Oxidation states	+1, +2, +3, +4 (a mildly basic oxide)
Electronegativity	Pauling scale: 1.12
Ionization energies	1st: 534.4 kJ/mol 2nd: 1050 kJ/mol 3rd: 1949 kJ/mol (more)
Atomic radius	empirical: 181.8 pm
Covalent radius	204±9 pm
Spectral lines of cerium
Other properties
Crystal structure	​double hexagonal close-packed (dhcp) β-Ce
Crystal structure	​face-centered cubic (fcc) γ-Ce
Speed of sound thin rod	2100 m/s (at 20 °C)
Thermal expansion	γ, poly: 6.3 µm/(m·K) (at r.t.)
Thermal conductivity	11.3 W/(m·K)
Electrical resistivity	β, poly: 828 nΩ·m (at r.t.)
Magnetic ordering	paramagnetic[3]
Magnetic susceptibility	(β) +2450.0·10−6 cm3/mol (293 K)[4]
Young's modulus	γ form: 33.6 GPa
Shear modulus	γ form: 13.5 GPa
Bulk modulus	γ form: 21.5 GPa
Poisson ratio	γ form: 0.24
Mohs hardness	2.5
Vickers hardness	210–470 MPa
Brinell hardness	186–412 MPa
CAS Number	7440-45-1
History
Naming	after dwarf planet Ceres, itself named after Roman deity of agriculture Ceres
Discovery	Martin Heinrich Klaproth, Jöns Jakob Berzelius, Wilhelm Hisinger (1803)
First isolation	Carl Gustaf Mosander (1838)
Main isotopes of cerium
Iso­tope Abun­dance Half-life (t1/2) Decay mode Pro­duct 134Ce syn 3.16 d ε 134La 136Ce 0.186% stable 138Ce 0.251% stable 139Ce syn 137.640 d ε 139La 140Ce 88.449% stable 141Ce syn 32.501 d β− 141Pr 142Ce 11.114% stable 143Ce syn 33.039 d β− 143Pr 144Ce syn 284.893 d β− 144Pr
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Cerium is a chemical element with symbol Ce and atomic number 58. Cerium is a soft, ductile and silvery-white metal that tarnishes when exposed to air, and it is soft enough to be cut with a knife. Cerium is the second element in the lanthanide series, and while it often shows the +3 oxidation state characteristic of the series, it also exceptionally has a stable +4 state that does not oxidize water. It is also considered one of the rare-earth elements. Cerium has no biological role and is not very toxic.

Despite always occurring in combination with the other rare-earth elements in minerals such as those of the monazite and bastnäsite groups, cerium is easy to extract from its ores, as it can be distinguished among the lanthanides by its unique ability to be oxidized to the +4 state. It is the most common of the lanthanides, followed by neodymium, lanthanum, and praseodymium. It is the 26th-most abundant element, making up 66 ppm of the Earth's crust, half as much as chlorine and five times as much as lead.

Cerium was the first of the lanthanides to be discovered, in Bastnäs, Sweden by Jöns Jakob Berzelius and Wilhelm Hisinger in 1803, and independently by Martin Heinrich Klaproth in Germany in the same year. In 1839 Carl Gustaf Mosander became the first to isolate the metal. Today, cerium and its compounds have a variety of uses: for example, cerium(IV) oxide is used to polish glass and is an important part of catalytic converters. Cerium metal is used in ferrocerium lighters for its pyrophoric properties. Cerium-doped YAG phosphor is used in blue light-emitting diodes to produce white light.

Characteristics

Physical

Cerium is the second element of the lanthanide series. In the periodic table, it appears between the lanthanides lanthanum to its left and praseodymium to its right, and above the actinide thorium. It is a ductile metal with a hardness similar to that of silver.^[5] Its 58 electrons are arranged in the configuration [Xe]4f¹5d¹6s², of which the four outer electrons are valence electrons. Immediately after lanthanum, the 4f orbitals suddenly contract and are lowered in energy to the point that they participate readily in chemical reactions; however, this effect is not yet strong enough at cerium and thus the 5d subshell is still occupied.^[6] Most lanthanides can use only three electrons as valence electrons, as afterwards the remaining 4f electrons are too strongly bound: cerium is an exception because of the stability of the empty f-shell in Ce⁴⁺ and the fact that it comes very early in the lanthanide series, where the nuclear charge is still low enough until neodymium to allow the removal of the fourth valence electron by chemical means.^[7]

Phase diagram of cerium

Four allotropic forms of cerium are known to exist at standard pressure, and are given the common labels of α to δ:^[8]

• The high-temperature form, δ-cerium, has a bcc (body-centred cubic) crystal structure and exists above 726 °C.


• The stable form below 726 °C to approximately room temperature is γ-cerium, with an fcc (face-centred cubic) crystal structure.


• The dhcp (double hexagonal close-packed) form β-cerium is the equilibrium structure approximately from room temperature to −150 °C.


• The fcc form α-cerium is stable below about −150 °C; it has a density of 8.16 g/cm³.


• Other solid phases occurring only at high pressures are shown on the phase diagram.


• Both γ and β forms are quite stable at room temperature, although the equilibrium transformation temperature is estimated at around 75 °C.^[8]
  

Cerium has a variable electronic structure. The energy of the 4f electron is nearly the same as that of the outer 5d and 6s electrons that are delocalized in the metallic state, and only a small amount of energy is required to change the relative occupancy of these electronic levels. This gives rise to dual valence states. For example, a volume change of about 10% occurs when cerium is subjected to high pressures or low temperatures. It appears that the valence changes from about 3 to 4 when it is cooled or compressed.^[9]

At lower temperatures the behavior of cerium is complicated by the slow rates of transformation. Transformation temperatures are subject to substantial hysteresis and values quoted here are approximate. Upon cooling below −15 °C, γ-cerium starts to change to β-cerium, but the transformation involves a volume increase and, as more β forms, the internal stresses build up and suppress further transformation.^[8] Cooling below approximately −160 °C will start formation of α-cerium but this is only from remaining γ-cerium. β-cerium does not significantly transform to α-cerium except in the presence of stress or deformation.^[8] At atmospheric pressure, liquid cerium is more dense than its solid form at the melting
point.^[5][10][11]

Isotopes

Naturally occurring cerium is made up of four isotopes: ¹³⁶Ce (0.19%), ¹³⁸Ce (0.25%), ¹⁴⁰Ce (88.4%), and ¹⁴²Ce (11.1%). All four are observationally stable, though the light isotopes ¹³⁶Ce and ¹³⁸Ce are theoretically expected to undergo inverse double beta decay to isotopes of barium, and the heaviest isotope ¹⁴²Ce is expected to undergo double beta decay to ¹⁴²Nd or alpha decay to ¹³⁸Ba. Additionally, ¹⁴⁰Ce would release energy upon spontaneous fission. None of these decay modes have yet been observed, though the double beta decay of ¹³⁶Ce, ¹³⁸Ce, and ¹⁴²Ce has been experimentally searched for. The current experimental limits for their half-lives are:^[12]

¹³⁶Ce: >3.8×10¹⁶ y

¹³⁸Ce: >1.5×10¹⁴ y

¹⁴²Ce: >5×10¹⁶ y

All other cerium isotopes are synthetic and radioactive. The most stable of them are ¹⁴⁴Ce with a half-life of 284.9 days, ¹³⁹Ce with a half-life of 137.6 days, ¹⁴³Ce with a half-life of 33.04 days, and ¹⁴¹Ce with a half-life of 32.5 days. All other radioactive cerium isotopes have half-lives under four days, and most of them have half-lives under ten minutes.^[12] The isotopes between ¹⁴⁰Ce and ¹⁴⁴Ce inclusive occur as fission products of uranium.^[12] The primary decay mode of the isotopes lighter than ¹⁴⁰Ce is inverse beta decay or electron capture to isotopes of lanthanum, while that of the heavier isotopes is beta decay to isotopes of praseodymium.^[12]

The great rarity of the proton-rich ¹³⁶Ce and ¹³⁸Ce is explained by the fact that they cannot be made in the most common processes of stellar nucleosynthesis for elements beyond iron, the s-process (slow neutron capture) and the r-process (rapid neutron capture). This is so because they are bypassed by the reaction flow of the s-process, and the r-process nuclides are blocked from decaying to them by more neutron-rich stable nuclides. Such nuclei are called p-nuclei, and their origin is not yet well understood: some speculated mechanisms for their formation include proton capture as well as photodisintegration.^[13]140Ce is the most common isotope of cerium, as it can be produced in both the s- and r-processes, while ¹⁴²Ce can only be produced in the r-process. Another reason for the abundance of ¹⁴⁰Ce is that it is a magic nucleus, having a closed neutron shell (it has 82 neutrons), and hence it has a very low cross-section towards further neutron capture. Although its proton number of 58 is not magic, it is granted additional stability, as its eight additional protons past the magic number 50 enter and complete the 1 g_7/2 proton orbital.^[13] The abundances of the cerium isotopes may differ very slightly in natural sources, because ¹³⁸Ce and ¹⁴⁰Ce are the daughters of the long-lived primordial radionuclides ¹³⁸La and ¹⁴⁴Nd, respectively.^[12]

Chemistry

Cerium tarnishes in air, forming a spalling oxide layer like iron rust; a centimeter-sized sample of cerium metal corrodes completely in about a year.^[14] It burns readily at 150 °C to form the pale-yellow cerium(IV) oxide, also known as ceria:^[15]

Ce + O₂ → CeO₂

This may be reduced to cerium(III) oxide with hydrogen gas.^[16] Cerium metal is highly pyrophoric, meaning that when it is ground or scratched, the resulting shavings catch fire.^[17] This reactivity conforms to periodic trends, since cerium is one of the first and hence one of the largest lanthanides.^[18] Cerium(IV) oxide has the fluorite structure, similarly to the dioxides of praseodymium and terbium. Many nonstoichiometric chalcogenides are also known, along with the trivalent Ce₂Z₃ (Z = S, Se, Te). The monochalcogenides CeZ conduct electricity and would better be formulated as Ce³⁺Z²⁻e⁻. While CeZ₂ are known, they are polychalcogenides with cerium(III): cerium(IV) chalcogenides remain unknown.^[16]

Cerium(IV) oxide

Cerium is a highly electropositive metal and reacts with water. The reaction is slow with cold water but speeds up with increasing temperature, producing cerium(III) hydroxide and hydrogen gas:^[15]

2 Ce (s) + 6 H₂O (l) → 2 Ce(OH)₃ (aq) + 3 H₂ (g)

Cerium metal reacts with all the halogens to give trihalides:^[15]

2 Ce (s) + 3 F₂ (g) → 2 CeF₃ (s) [white]

2 Ce (s) + 3 Cl₂ (g) → 2 CeCl₃ (s) [white]

2 Ce (s) + 3 Br₂ (g) → 2 CeBr₃ (s) [white]

2 Ce (s) + 3 I₂ (g) → 2 CeI₃ (s) [yellow]

Reaction with excess fluorine produces the stable white tetrafluoride CeF₄; the other tetrahalides are not known. Of the dihalides, only the bronze diiodide CeI₂ is known; like the diiodides of lanthanum, praseodymium, and gadolinium, this is a cerium(III) electride compound.^[19] True cerium(II) compounds are restricted to a few unusual organocerium complexes.^[20][21]

Cerium dissolves readily in dilute sulfuric acid to form solutions containing the colorless Ce³⁺ ions, which exist as a [Ce(H₂O)₉]³⁺ complexes:^[15]

2 Ce (s) + 3 H₂SO₄ (aq) → 2 Ce³⁺ (aq) + 3 SO²⁻
₄ (aq) + 3 H₂ (g)

The solubility of cerium is much higher in methanesulfonic acid.^[22] Cerium(III) and terbium(III) have ultraviolet absorption bands of relatively high intensity compared with the other lanthanides, as their configurations (one electron more than an empty or half-filled f-subshell respectively) make it easier for the extra f electron to undergo f→d transitions instead of the forbidden f→f transitions of the other lanthanides.^[23]Cerium(III) sulfate is one of the few salts whose solubility in water decreases with rising temperature.^[24]

Ceric ammonium nitrate

Cerium(IV) aqueous solutions may be prepared by reacting cerium(III) solutions with the strong oxidising agents peroxodisulfate or bismuthate. The value of E^⦵(Ce⁴⁺/Ce³⁺) varies widely depending on conditions due to the relative ease of complexation and hydrolysis with various anions, though +1.72 V is a usually representative value; that for E^⦵(Ce³⁺/Ce) is −2.34 V. Cerium is the only lanthanide which has important aqueous and coordination chemistry in the +4 oxidation state.^[25] Due to ligand-to-metal charge transfer, aqueous cerium(IV) ions are orange-yellow.^[26] Aqueous cerium(IV) is metastable in water^[27] and is a strong oxidising agent that oxidizes hydrochloric acid to give chlorine gas.^[25] For example, ceric ammonium nitrate is a common oxidising agent in organic chemistry, releasing organic ligands from metal carbonyls.^[28] In the Belousov–Zhabotinsky reaction, cerium oscillates between the +4 and +3 oxidation states to catalyse the reaction.^[29] Cerium(IV) salts, especially cerium(IV) sulfate, are often used as standard reagents for volumetric analysis in cerimetric titrations.^[30]

The nitrate complex [Ce(NO₃)₆]²⁻ is the most common cerium complex encountered when using cerium(IV) is an oxidising agent: it and its cerium(III) analogue [Ce(NO₃)₆]³⁻ have 12-coordinate icosahedral molecular geometry, while [Ce(NO₃)₅]²⁻ has 10-coordinate bicapped dodecadeltahedral molecular geometry. Cerium nitrates also form 4:3 and 1:1 complexes with 18-crown-6 (the ratio referring to that between cerium and the crown ether). Halogen-containing complex ions such as CeF⁴⁻
₈, CeF²⁻
₆, and the orange CeCl²⁻
₆ are also known.^[25]Organocerium chemistry is similar to that of the other lanthanides, being primarily that of the cyclopentadienyl and cyclooctatetraenyl compounds. The cerium(III) cyclooctatetraenyl compound has the uranocene structure.^[31]

Cerium(IV)

Despite the common name of cerium(IV) compounds, the Japanese spectroscopist Akio Kotani wrote "there is no genuine example of cerium(IV)". The reason for this can be seen in the structure of ceria itself, which always contains some octahedral vacancies where oxygen atoms would be expected to go and could be better considered a non-stoichiometric compound with chemical formula CeO_2−x. Furthermore, each cerium atom in ceria does not lose all four of its valence electrons, but retains a partial hold on the last one, resulting in an oxidation state between +3 and +4.^[32][33] Even supposedly purely tetravalent compounds such as CeRh₃, CeCo₅, or ceria itself have X-ray photoemission and X-ray absorption spectra more characteristic of intermediate-valence compounds.^[34] The 4f electron in cerocene, Ce(C₈H₈)₂, is poised ambiguously between being localized and delocalized and this compound is also considered intermediate-valent.^[33]

History

The dwarf planet Ceres, after which cerium is named

Cerium was discovered in Bastnäs in Sweden by Jöns Jakob Berzelius and Wilhelm Hisinger, and independently in Germany by Martin Heinrich Klaproth, both in 1803.^[35] Cerium was named by Berzelius after the dwarf planet Ceres, discovered two years earlier.^[35][36] The dwarf planet itself is named after the Roman goddess of agriculture, grain crops, fertility and motherly relationships, Ceres.^[35]

Cerium was originally isolated in the form of its oxide, which was named ceria, a term that is still used. The metal itself was too electropositive to be isolated by then-current smelting technology, a characteristic of rare-earth metals in general. After the development of electrochemistry by Humphry Davy five years later, the earths soon yielded the metals they contained. Ceria, as isolated in 1803, contained all of the lanthanides present in the cerite ore from Bastnäs, Sweden, and thus only contained about 45% of what is now known to be pure ceria. It was not until Carl Gustaf Mosander succeeded in removing lanthana and "didymia" in the late 1830s that ceria was obtained pure. Wilhelm Hisinger was a wealthy mine-owner and amateur scientist, and sponsor of Berzelius. He owned and controlled the mine at Bastnäs, and had been trying for years to find out the composition of the abundant heavy gangue rock (the "Tungsten of Bastnäs", which despite its name contained no tungsten), now known as cerite, that he had in his mine.^[36] Mosander and his family lived for many years in the same house as Berzelius, and Mosander was undoubtedly persuaded by Berzelius to investigate ceria further.^[37]

Occurrence and production

Cerium is the most abundant of all the lanthanides, making up 66 ppm of the Earth's crust; this value is just behind that of copper (68 ppm), and cerium is even more abundant than common metals such as lead (13 ppm) and tin (2.1 ppm). Thus, despite its position as one of the so-called rare-earth metals, cerium is actually not rare at all.^[38] Cerium content in the soil varies between 2 and 150 ppm, with an average of 50 ppm; seawater contains 1.5 parts per trillion of cerium.^[36] Cerium occurs in various minerals, but the most important commercial sources are the minerals of the monazite and bastnäsite groups, where it makes up about half of the lanthanide content. Monazite-(Ce) is the most common representative of the monazites, with "-Ce" being the Levinson suffix informing on the dominance of the particular REE element representative.^[39][40][41]). Also the cerium-dominant bastnäsite-(Ce) is the most important of the bastnäsites.^[42][43] Cerium is the easiest lanthanide to extract from its minerals because it is the only one that can reach a stable +4 oxidation state in aqueous solution.^[44] Because of the decreased solubility of cerium in the +4 oxidation state, cerium is sometimes depleted from rocks relative to the other rare-earth elements and is incorporated into zircon, since Ce⁴⁺ and Zr⁴⁺ have the same charge and similar ionic radii.^[45] In extreme cases, cerium(IV) can form its own minerals separated from the other rare-earth elements, such as cerianite (correctly named cerianite-(Ce)^[46][47][48]), (Ce,Th)O₂.^[49][50][51]

Crystal structure of bastnäsite-(Ce). Color code: carbon, C, blue-gray; fluorine, F, green; cerium, Ce, white; oxygen, O, red.

Bastnäsite, Ln^IIICO₃F, is usually lacking in thorium and the heavy lanthanides beyond samarium and europium, and hence the extraction of cerium from it is quite direct. First, the bastnäsite is purified, using dilute hydrochloric acid to remove calcium carbonate impurities. The ore is then roasted in the air to oxidize it to the lanthanide oxides: while most of the lanthanides will be oxidized to the sesquioxides Ln₂O₃, cerium will be oxidized to the dioxide CeO₂. This is insoluble in water and can be leached out with 0.5 M hydrochloric acid, leaving the other lanthanides behind.^[44]

The procedure for monazite, (Ln,Th)PO₄, which usually contains all the rare earths, as well as thorium, is more involved. Monazite, because of its magnetic properties, can be separated by repeated electromagnetic separation. After separation, it is treated with hot concentrated sulfuric acid to produce water-soluble sulfates of rare earths. The acidic filtrates are partially neutralized with sodium hydroxide to pH 3–4. Thorium precipitates out of solution as hydroxide and is removed. After that, the solution is treated with ammonium oxalate to convert rare earths to their insoluble oxalates. The oxalates are converted to oxides by annealing. The oxides are dissolved in nitric acid, but cerium oxide is insoluble in HNO₃ and hence precipitates out.^[11] Care must be taken when handling some of the residues as they contain ²²⁸Ra, the daughter of ²³²Th, which is a strong gamma emitter.^[44]

Applications

Carl Auer von Welsbach, who discovered many applications of cerium

The first use of cerium was in gas mantles, invented by the Austrian chemist Carl Auer von Welsbach. In 1885, he had previously experimented with mixtures of magnesium, lanthanum, and yttrium oxides, but these gave green-tinted light and were unsuccessful.^[52] Six years later, he discovered that pure thorium oxide produced a much better, though blue, light, and that mixing it with cerium dioxide resulted in a bright white light.^[53] Additionally, cerium dioxide also acts as a catalyst for the combustion of thorium oxide. This resulted in great commercial success for von Welsbach and his invention, and created great demand for thorium; its production resulted in a large amount of lanthanides being simultaneously extracted as by-products.^[54] Applications were soon found for them, especially in the pyrophoric alloy known as "mischmetall" composed of 50% cerium, 25% lanthanum, and the remainder being the other lanthanides, that is used widely for lighter flints.^[54] Usually, iron is also added to form an alloy known as ferrocerium, also invented by von Welsbach.^[55] Due to the chemical similarities of the lanthanides, chemical separation is not usually required for their applications, such as the mixing of mischmetall into steel to improve its strength and workability, or as catalysts for the cracking of petroleum.^[44] This property of cerium saved the life of writer Primo Levi at the Auschwitz concentration camp, when he found a supply of ferrocerium alloy and bartered it for food.^[56]

Ceria is the most widely used compound of cerium. The main application of ceria is as a polishing compound, for example in chemical-mechanical planarization (CMP). In this application, ceria has replaced other metal oxides for the production of high-quality optical surfaces.^[55] Major automotive applications for the lower sesquioxide are as a catalytic converter for the oxidation of CO and NO_x emissions in the exhaust gases from motor vehicles,^[57][58] Ceria has also been used as a substitute for its radioactive congener thoria, for example in the manufacture of electrodes used in gas tungsten arc welding, where ceria as an alloying element improves arc stability and ease of starting while decreasing burn-off.^[59]Cerium(IV) sulfate is used as an oxidising agent in quantitative analysis. Cerium(IV) in methanesulfonic acid solutions is applied in industrial scale electrosynthesis as a recyclable oxidant.^[60] Ceric ammonium nitrate is used as an oxidant in organic chemistry and in etching electronic components, and as a primary standard for quantitative analysis.^[5][61]

The photostability of pigments can be enhanced by the addition of cerium. It provides pigments with light fastness and prevents clear polymers from darkening in sunlight. Television glass plates are subject to electron bombardment, which tends to darken them by creation of F-center color centers. This effect is suppressed by addition of cerium oxide. Cerium is also an essential component of phosphors used in TV screens and fluorescent lamps.^[62][63] Cerium sulfide forms a red pigment that stays stable up to 350 °C. The pigment is a nontoxic alternative to cadmium sulfide pigments.^[36]

Cerium is used as alloying element in aluminum to create castable eutectic alloys, Al-Ce alloys with 6–16 wt.% Ce, to which Mg and/or Si can be further added; these alloys have excellent high temperature strength.^[64]

Biological role and precautions

Cerium

Hazards
GHS pictograms
GHS signal word	Danger
GHS hazard statements	H228, H302, H312, H332, H315, H319, H335
GHS precautionary statements	P210, P261, P280, P301, P312, P330, P305, P351, P338, P370, P378[65]
NFPA 704	0 2 0

Cerium has no known biological role in humans, but is not very toxic either; it does not accumulate in the food chain to any appreciable extent. Because it often occurs together with calcium in phosphate minerals, and bones are primarily calcium phosphate, cerium can accumulate in bones in small amounts that are not considered dangerous. Cerium, like the other lanthanides, is known to affect human metabolism, lowering cholesterol levels, blood pressure, appetite, and risk of blood coagulation. Cerium nitrate is an effective topical antimicrobial treatment for third-degree burns,^[36][66] although large doses can lead to cerium poisoning and methemoglobinemia.^[67] The early lanthanides act as essential cofactors for the methanol dehydrogenase of the methanotrophic bacterium Methylacidiphilum fumariolicum SolV, for which lanthanum, cerium, praseodymium, and neodymium alone are about equally effective.^[68]

Like all rare-earth metals, cerium is of low to moderate toxicity. A strong reducing agent, it ignites spontaneously in air at 65 to 80 °C. Fumes from cerium fires are toxic. Water should not be used to stop cerium fires, as cerium reacts with water to produce hydrogen gas. Workers exposed to cerium have experienced itching, sensitivity to heat, and skin lesions. Cerium is not toxic when eaten, but animals injected with large doses of cerium have died due to cardiovascular collapse.^[36] Cerium is more dangerous to aquatic organisms, on account of being damaging to cell membranes, but this is not an important risk because it is not very soluble in water.^[36]

References

Bibliography

• 
  Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0-08-037941-9.