Cerium is achemical element; it hassymbolCe andatomic number 58, and is asoft,ductile, and silvery-whitemetal that tarnishes when exposed to air. Cerium is the second element in thelanthanide series, and while it often shows theoxidation state of +3 characteristic of the series, it also has a stable +4 state that does not oxidize water. It is considered one of therare-earth elements. Cerium has no known biological role in humans but is not particularly toxic, except with intense or continued exposure.
Despite always occurring in combination with the other rare-earth elements in minerals such as those of themonazite andbastnäsite 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 in aqueous solution. It is the most common of the lanthanides, followed byneodymium,lanthanum, andpraseodymium. Its estimatedabundance in the Earth's crust is 68 ppm.
Cerium is the second element of thelanthanide series. In the periodic table, it appears between the lanthanideslanthanum to its left andpraseodymium to its right, and above theactinidethorium. It is aductile metal with a hardness similar to that ofsilver.[9] Its 58 electrons are arranged in theconfiguration [Xe]4f15d16s2, of which the four outer electrons arevalence electrons.[10] The 4f, 5d, and 6s energy levels are very close to each other, and the transfer of one electron to the 5d shell is due to strong interelectronic repulsion in the compact 4f shell. This effect is overwhelmed when the atom is positively ionised; thus Ce2+ on its own has instead the regular configuration [Xe]4f2, although in some solid solutions it may be [Xe]4f15d1.[11] 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 untilneodymium to allow the removal of the fourth valence electron by chemical means.[12]
Cerium has a variableelectronic 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. In its high pressure phase (α-Cerium), the 4f electrons are also delocalized and itinerate, as opposed to localized 4f electrons in low pressure phase (γ-Cerium).[13] It appears that the valence changes from about 3 to 4 when it is cooled or compressed.[14]
Like the other lanthanides, cerium metal is a goodreducing agent, havingstandard reduction potential ofE⦵ = −2.34 V for the Ce3+/Ce couple.[15] It tarnishes in air, forming apassivating oxide layer likeiron rust. A centimeter-sized sample of cerium metal corrodes completely in about a year. More dramatically, metallic cerium can be highlypyrophoric:[16]
Ce + O2 → CeO2
Being highlyelectropositive, cerium reacts with water. The reaction is slow with cold water but speeds up with increasing temperature, producing cerium(III) hydroxide and hydrogen gas:[17]
Fourallotropic forms of cerium are known to exist at standard pressure and are given the common labels of α to δ:[18]
The high-temperature form, δ-cerium, has a bcc (body-centered 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-centered cubic) crystal structure.
The DHCP (doublehexagonal 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 75 °C.[18]
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.[18] 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.[18] At atmospheric pressure, liquid cerium is more dense than its solid form at the meltingpoint.[9][19][20]
All nuclear data not otherwise stated is from the standard source:[21]
Naturally occurring cerium is made up of four isotopes:136Ce (0.19%),138Ce (0.25%),140Ce (88.45%), and142Ce (11.11%). All areobservationally stable, though the light isotopes136Ce and138Ce are theoretically expected to undergodouble electron capture to isotopes ofbarium, and the heaviest isotope142Ce is expected to undergo double beta decay to142Nd or alpha decay to138Ba. Thus,140Ce is the only theoreticallystable isotope. None of these decay modes have yet been observed, though the double beta decays of136Ce,138Ce, and142Ce have been experimentally searched for. The current experimental limits for their half-lives are about 1×1017, 4×1017, and 3×1018 years - all short compared to known double-beta half-lives.
All other cerium isotopes aresynthetic andradioactive. The most stable of them are144Ce with a half-life of 284.9 days,139Ce with a half-life of 137.6 days, and141Ce 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. The isotopes between140Ce and144Ce inclusive occur asfission products ofuranium.[22] The primary decay mode of the isotopes lighter than140Ce isinverse beta decay orelectron capture toisotopes of lanthanum, while that of the heavier isotopes isbeta decay toisotopes of praseodymium. Someisotopes of neodymium canalpha decay or are predicted to decay to isotopes of cerium.[23]
Cerium exists in two main oxidation states, Ce(III) and Ce(IV). This pair of adjacent oxidation states dominates several aspects of the chemistry of this element. Cerium(IV) aqueous solutions may be prepared by reacting cerium(III) solutions with the strong oxidizing agentsperoxodisulfate orbismuthate. The value ofE⦵(Ce4+/Ce3+) varies widely depending on conditions due to the relative ease of complexation and hydrolysis with various anions, although +1.72 V is representative. Cerium is the only lanthanide which has important aqueous and coordination chemistry in the +4 oxidation state.[15]
Cerium forms all four trihalides CeX3 (X = F, Cl, Br, I) usually by reaction of the oxides with the hydrogen halides. The anhydrous halides are pale-colored, paramagnetic, hygroscopic solids. Upon hydration, the trihalides convert to complexes containing aquo complexes [Ce(H2O)8-9]3+. Unlike most lanthanides, Ce forms a tetrafluoride, a white solid. It also forms a bronze-colored diiodide, which has metallic properties.[24] Aside from the binary halide phases, a number of anionic halide complexes are known. The fluoride gives the Ce(IV) derivativesCeF4−8 andCeF2−6. The chloride gives the orangeCeCl2−6.[15]
Cerium(IV) oxide ("ceria") has thefluorite structure, similarly to the dioxides of praseodymium andterbium. Ceria is anonstoichiometric compound, meaning that the real formula is CeO2−x, where x is about 0.2. Thus, the material is not perfectly described as Ce(IV). Ceria reduces tocerium(III) oxide with hydrogen gas.[25] Manynonstoichiometricchalcogenides 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) derivatives of S, Se, and Te are unknown.[25]
The compoundceric ammonium nitrate (CAN)(NH4)2[Ce(NO3)6] is the most common cerium compound encountered in the laboratory. The six nitrate ligands bind asbidentate ligands. The complex[Ce(NO3)6]2− is 12-coordinate, a high coordination number which emphasizes the large size of the Ce4+ ion. CAN is a popular oxidant inorganic synthesis, both as a stoichiometric reagent[26] and as a catalyst.[27] It is inexpensive, stable in air, easily handled, and of low toxicity.[27] It operates by one-electron redox. Cerium nitrates also form 4:3 and 1:1 complexes with18-crown-6 (the ratio referring to that between the nitrate and thecrown ether). Classically, CAN is a primary standard for quantitative analysis.[9][28] Cerium(IV) salts, especiallycerium(IV) sulfate, are often used as standard reagents forvolumetric analysis incerimetric titrations.[29]
A white LED in operation: the diode produces monochromatic blue light but the Ce:YAG phosphor converts some of it into yellow light; the combination is perceived as white by the human eye.
Cerium was originally isolated in the form of its oxide, which was namedceria, 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 ofelectrochemistry byHumphry 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 untilCarl 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 notungsten), now known as cerite, that he had in his mine.[38] 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.[39][40][41][42]
The element played a role in theManhattan Project, where cerium compounds were investigated in theBerkeley site as materials forcrucibles foruranium andplutonium casting.[43] For this reason,new methods for the preparation and casting of cerium were developed within the scope of theAmes daughter project (now theAmes Laboratory).[44] Production of extremely pure cerium in Ames commenced in mid-1944 and continued until August 1945.[44]
Cerium is the most abundant of all the lanthanides and the 25th most abundant element, making up 68 ppm of the Earth's crust.[45] This value is the same ofcopper, and cerium is even more abundant than common metals such aslead (13 ppm) andtin (2.1 ppm). Thus, despite its position as one of the so-calledrare-earth metals, cerium is actually not rare at all.[46] 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.[38] Cerium occurs in various minerals, but the most important commercial sources are the minerals of themonazite andbastnäsite 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.[47][48][49] Also the cerium-dominant bastnäsite-(Ce) is the most important of the bastnäsites.[50][47] 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.[51] 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 intozircon, since Ce4+ andZr4+ have the same charge and similar ionic radii.[52] In extreme cases, cerium(IV) can form its own minerals separated from the other rare-earth elements, such ascerianite-(Ce)[53][49][47] and(Ce,Th)O2.[54]
Crystal structure of bastnäsite-(Ce). Color code: carbon, C, blue-gray; fluorine, F, green; cerium, Ce, white; oxygen, O, red.
Bastnäsite, LnIIICO3F, is usually lacking inthorium and the heavy lanthanides beyondsamarium andeuropium, and hence the extraction of cerium from it is quite direct. First, the bastnäsite is purified, using dilutehydrochloric acid to removecalcium 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 sesquioxidesLn2O3, 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.[51]
The procedure formonazite,(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 withsodium hydroxide to pH 3–4. Thorium precipitates out of solution as hydroxide and is removed. After that, the solution is treated withammonium oxalate to convert rare earths to their insolubleoxalates. 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.[20] Care must be taken when handling some of the residues as they contain228Ra, the daughter of232Th, which is a strong gamma emitter.[51]
Carl Auer von Welsbach, who discovered many applications of cerium
Cerium has two main applications, both of which use CeO2. The industrial application of ceria is for polishing, especiallychemical-mechanical planarization (CMP). In its other main application, CeO2 is used to decolorize glass. It functions by converting green-tinted ferrous impurities to nearly colorless ferric oxides.[55] Ceria has also been used as a substitute for its radioactive congenerthoria, for example in the manufacture of electrodes used ingas tungsten arc welding, where cerium as an alloying element improves arc stability and ease of starting while decreasing burn-off.[56]
The first use of cerium was ingas mantles, invented by Austrian chemistCarl Auer von Welsbach. In 1885, he had previously experimented with mixtures ofmagnesium, lanthanum, and yttrium oxides, but these gave green-tinted light and were unsuccessful.[57] Six years later, he discovered that purethorium oxide produced a much better, though blue, light, and that mixing it with cerium dioxide resulted in a bright white light.[58] Cerium dioxide also acts as a catalyst for the combustion of thorium oxide.[citation needed]
This resulted in 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.[59] Applications were soon found for them, especially in the pyrophoric alloy known as "mischmetal" composed of 50% cerium, 25% lanthanum, and the remainder being the other lanthanides, that is used widely for lighter flints.[59] Usually iron is added to form the alloyferrocerium, also invented by von Welsbach.[60] Due to the chemical similarities of the lanthanides, chemical separation is not usually required for their applications, such as the addition of mischmetal to steel as an inclusion modifier to improve mechanical properties, or as catalysts for the cracking of petroleum.[51] This property of cerium saved the life of writerPrimo Levi at theAuschwitz concentration camp, when he found a supply of ferrocerium alloy and bartered it for food.[61]
The photostability ofpigments can be enhanced by the addition of cerium, as it provides pigments withlightfastness and prevents clear polymers from darkening in sunlight.[62] An example of a cerium compound used on its own as aninorganic pigment is the vivid redcerium(III) sulfide (cerium sulfide red), which stays chemically inert up to very high temperatures. The pigment is a safer alternative to lightfast but toxiccadmium selenide-based pigments.[38] The addition of cerium oxide to oldercathode-ray tube television glass plates was beneficial, as it suppresses the darkening effect from the creation ofF-center defects due to the continuous electron bombardment during operation. Cerium is also an essential component as adopant forphosphors used in CRT TV screens, fluorescent lamps, and laterwhite light-emitting diodes.[63][64] The most commonly used example iscerium(III)-doped yttrium aluminium garnet (Ce:YAG) which emits yellow light (530–540 nm)[65] and also behaves as ascintillator.[66]
Cerium salts, such as the sulfidesCe2S3 and Ce3S4, were considered during theManhattan Project as advancedrefractory materials for the construction of crucibles which could withstand the high temperatures and stronglyreducing conditions when casting plutonium metal.[43][44] Despite desirable properties, these sulfides were never widely adopted due to practical issues with their synthesis.[43] Cerium is used as alloying element in aluminium to create castable eutecticaluminium alloys with 6–16 wt.% Ce, to which other elements such as Mg, Ni, Fe and Mn can be added. These Al-Ce alloys have excellent high temperature strength and are suitable for automotive applications (e.g. incylinder heads).[67] Other alloys of cerium include Pu-Ce and Pu-Ce-Coplutonium alloys, which have been used asnuclear fuel.[68]
Cerium oxide can be used incatalytic converters to increase the efficiency of oxidation of CO andNOx emissions during low-oxygen conditions in the exhaust gases from motor vehicles.[69][70]
Cerium nitrate is an effective topical antimicrobial treatment forthird-degree burns,[38][78] although large doses can lead to cerium poisoning andmethemoglobinemia.[79] Like all rare-earth metals, cerium is of low to moderate toxicity.[80] A strong reducing agent, it ignites spontaneously in air at 65 to 80 °C. Fumes from cerium fires are toxic.[38] Cerium reacts with water to produce hydrogen gas, and thus cerium fires can only be effectively extinguished usingclass D dry powder extinguishing media.[81] 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. Cerium is more dangerous to aquatic organisms because it damages cell membranes; it is not very soluble in water and can cause environmental contamination.[38]
^abcdArblaster, John W. (2018).Selected Values of the Crystallographic Properties of Elements. Materials Park, Ohio: ASM International.ISBN978-1-62708-155-9.
^Jørgensen, Christian (1973). "The Loose Connection between Electron Configuration and the Chemical Behavior of the Heavy Elements (Transuranics)".Angewandte Chemie International Edition.12 (1):12–19.doi:10.1002/anie.197300121.
^Holleman, Arnold Frederik; Wiberg, Egon (2001), Wiberg, Nils (ed.),Inorganic Chemistry, translated by Eagleson, Mary; Brewer, William, San Diego/Berlin: Academic Press/De Gruyter, pp. 1703–5,ISBN0-12-352651-5
^abcdKoskimaki, D. C.; Gschneidner, K. A.; Panousis, N. T. (1974). "Preparation of single phase β and α cerium samples for low temperature measurements".Journal of Crystal Growth.22 (3):225–229.Bibcode:1974JCrGr..22..225K.doi:10.1016/0022-0248(74)90098-0.
^Nair, Vijay; Deepthi, Ani (2007). "Cerium(IV) Ammonium Nitrate: A Versatile Single-Electron Oxidant".Chemical Reviews.107 (5):1862–1891.doi:10.1021/cr068408n.PMID17432919.
^abSridharan, Vellaisamy; Menéndez, J. Carlos (2010). "Cerium(IV) Ammonium Nitrate as a Catalyst in Organic Synthesis".Chemical Reviews.110 (6):3805–3849.doi:10.1021/cr100004p.PMID20359233.
^B. P. Belousov (1959). "Периодически действующая реакция и ее механизм" [Periodically acting reaction and its mechanism].Сборник рефератов по радиационной медицине (in Russian).147: 145.
^Mikhail N. Bochkarev (2004). "Molecular compounds of "new" divalent lanthanides".Coordination Chemistry Reviews.248 (9–10):835–851.doi:10.1016/j.ccr.2004.04.004.
^M. Cristina Cassani; Yurii K. Gun'ko; Peter B. Hitchcock; Alexander G. Hulkes; Alexei V. Khvostov; Michael F. Lappert; Andrey V. Protchenko (2002). "Aspects of non-classical organolanthanide chemistry".Journal of Organometallic Chemistry.647 (1–2):71–83.doi:10.1016/s0022-328x(01)01484-x.
^Weeks, Mary Elvira (1932). "The Discovery of the Elements: XI. Some Elements Isolated with the Aid of Potassium and Sodium: Zirconium, Titanium, Cerium and Thorium".The Journal of Chemical Education.9 (7):1231–1243.Bibcode:1932JChEd...9.1231W.doi:10.1021/ed009p1231.
^abcHadden, Gavin, ed. (1946). "Chapter 11 - Ames Project".Manhattan District History. Vol. 4. Washington, D.C.: United States Army Corps of Engineers.
^Thomas, J. B.; Bodnar, R. J.; Shimizu, N.; Chesner, C. A. (2003). "Melt inclusions in zircon".Reviews in Mineralogy and Geochemistry.53 (1):63–87.Bibcode:2003RvMG...53...63T.doi:10.2113/0530063.
^Klaus Reinhardt and Herwig Winkler in "Cerium Mischmetal, Cerium Alloys, and Cerium Compounds" in Ullmann's Encyclopedia of Industrial Chemistry 2000, Wiley-VCH, Weinheim.doi:10.1002/14356007.a06_139
^Gleń, Marta; Grzmil, Barbara; Sreńscek-Nazzal, Joanna; Kic, Bogumił (2011-04-01). "Effect of CeO2 and Sb2O3 on the phase transformation and optical properties of photostable titanium dioxide".Chemical Papers.65 (2):203–212.doi:10.2478/s11696-010-0103-x.ISSN1336-9075.S2CID94458692.
_NIST_ASD_emission_spectrum.png)
.jpg)
.jpg)
