Cerium is a chemical element with symbol
Ce and atomic number 58. It is a soft,
silvery, ductile metal which easily
oxidizes in air. Cerium was named after
the dwarf planet Ceres. Cerium is the
most abundant of the rare earth
elements, making up about 0.0046% of the
Earth's crust by weight. It is found in
a number of minerals, the most important
being monazite and bastnäsite.
Commercial applications of cerium are
numerous. They include catalysts,
additives to fuel to reduce emissions
and to glass and enamels to change their
color. Cerium oxide is an important
component of glass polishing powders and
phosphors used in screens and
fluorescent lamps. It is also used in
the "flint" of lighters.
Characteristics
= Physical properties=
Cerium is a silvery metal, belonging to
the lanthanide group. It resembles iron
in color and luster, but is soft, and
both malleable and ductile. Cerium has
the second-longest liquid range of any
non-radioactive element: 2648 °C or 4766
°F.
Cerium has a variable electronic
structure. The energy of the inner 4f
level is nearly the same as that of the
outer or valence electrons, and only
small energy is required to change the
relative occupancy of these electronic
levels. This gives rise to dual valency
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.
Four allotropic forms of cerium are
known to exist at standard pressure, and
are given the common labels of α to δ:.
The high-temperature form, δ-cerium, has
a bcc crystal structure and exists above
726 °C.
The stable form below 726 °C to
approximately room temperature is
γ-cerium, with an fcc crystal structure.
The dhcp form of β-cerium is the
equilibrium structure approximately from
room temperature to −150 °C
The fcc α-cerium exists 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 ca. 75 °C.
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. 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.
At atmospheric pressure, liquid cerium
is more dense than its solid form at the
melting point.
= Chemical properties=
Cerium metal tarnishes slowly in air and
burns readily at 150 °C to form
cerium(IV) oxide:
Ce + O2 → CeO2
Cerium metal is highly pyrophoric,
meaning that when it is ground or
scratched, the resulting shavings catch
fire.
Cerium is quite electropositive and
reacts slowly with cold water and quite
quickly with hot water to form cerium
hydroxide:
2 Ce + 6 H2O → 2 Ce(OH)3 + 3 H2
Cerium metal reacts with all the
halogens:
2 Ce + 3 F2 → 2 CeF3 [white]
2 Ce + 3 Cl2 → 2 CeCl3 [white]
2 Ce + 3 Br2 → 2 CeBr3 [white]
2 Ce + 3 I2 → 2 CeI3 [yellow]
Cerium dissolves readily in dilute
sulfuric acid to form solutions
containing the colorless Ce(III) ions,
which exist as a [Ce(OH2)9]3+ complexes:
2 Ce + 3 H2SO4 → 2 Ce3+ + 3 SO2−
4 + 3 H2
The solubility of cerium is much higher
in methanesulfonic acid.
= Compounds=
Cerium(IV) salts are orange red or
yellowish, whereas cerium(III) salts are
usually white or colorless. Both
oxidation states absorb ultraviolet
light strongly. Cerium(III) can be used
to make glasses that are colorless, yet
absorb ultraviolet light almost
completely. Cerium can be readily
detected in rare earth mixtures by a
very sensitive qualitative test:
addition of ammonia and hydrogen
peroxide to an aqueous solution of
lanthanides produces a characteristic
dark brown color if cerium is present.
Cerium exhibits three oxidation states,
+2, +3 and +4. The +2 state is rare and
is observed in some organometallic
complexes. The compounds CeH2, CeI2 and
CeS have electride-like structure with
Ce3+ and delocalized electrons. The most
common compound of cerium is cerium(IV)
oxide, which is used as "Jeweller's
rouge" as well as in the walls of some
self-cleaning ovens. Two common
oxidizing agents used in titrations are
ammonium cerium(IV) sulfate2Ce(SO4)3)
and ammonium cerium(IV)
nitrate2Ce(NO3)6). Cerium also forms a
chloride, CeCl3 or cerium(III) chloride,
used to facilitate reactions at carbonyl
groups in organic chemistry. Other
compounds include cerium(III)
carbonate3), cerium(III) fluoride,
cerium(III) oxide, as well as cerium(IV)
sulfate2) and cerium(III) triflate3).
The differing properties of the two
oxidation states of cerium allows cerium
to be the most readily purified of all
the lanthanides. Otherwise these
elements are notoriously difficult to
separate. A wide range of procedures
have been devised that exploit the
difference. These include:
Leaching the mixed hydroxides with
dilute nitric acid: the trivalent
lanthanides dissolve in cerium-free
condition, and tetravalent cerium
remains in the insoluble residue as a
concentrate to be further purified by
other means. A variation on this uses
hydrochloric acid and the calcined
oxides from bastnasite, but the
separation is less sharp.
Precipitating cerium from a nitrate or
chloride solution using potassium
permanganate and sodium carbonate in a
1:4 molar ratio.
Boiling rare-earth nitrate solutions
with potassium bromate and marble chips.
Formerly used commercially was a method
whereby a solution of cerium(IV) in
nitric acid would be added to dilute
sulfuric acid. This step caused
cerium(IV) to largely precipitate as a
basic salt, leaving trivalent lanthanide
in solution. However, the finely divided
precipitate was difficult to filter from
the highly corrosive medium. Using the
classical methods of rare-earth
separation, there was a considerable
advantage to a strategy of removing
cerium from the mixture at the
beginning. Cerium typically comprised
45% of the cerite or monazite rare
earths, and removing it early greatly
reduced the bulk of what needed to be
further processed. However, not all
cerium purification methods relied on
basicity. Ceric ammonium nitrate
[ammonium hexanitratocerate(IV)]
crystallization from nitric acid was one
purification method. Cerium(IV) nitrate
was more readily extractable into
certain solvents than the trivalent
lanthanides. However, modern practice in
China seems to be to do purification of
cerium by counter-current solvent
extraction, in its trivalent form, just
like the other lanthanides.
Cerium(IV) is a strong oxidant under
acidic conditions, but stable under
alkaline conditions, when it is
cerium(III) that becomes a strong
reductant, easily oxidized by
atmospheric oxygen. This ease of
oxidation under alkaline conditions
leads to the occasional geochemical
parting of the ways between cerium and
the trivalent light lanthanides under
supergene weathering conditions, leading
variously to the "negative cerium
anomaly" or to the formation of the
mineral cerianite. Air-oxidation of
alkaline cerium(III) is the most
economical way to get to cerium(IV),
which can then be handled in acid
solution.
= Isotopes=
Four isotopes of cerium occur naturally.
These are 140Ce, which makes up 88.5% of
cerium on earth; 142Ce, which makes up
11% of cerium on earth; 138Ce, which
makes up 0.3% of cerium on earth; and
136Ce, which makes up 0.2% of cerium on
Earth. However, only 140Ce is stable.
142Ce, 138Ce, and 136Ce have half-lives
of 5×1016 years, 1.5×1015 years, and 70
trillion years respectively.
Additionally, some isotopes of cerium
occur naturally in trace amounts as the
result of fission of uranium.
39 isotopes of cerium have been
discovered. They range from 119Ce to
157Ce. Cerium also has 10 known nuclear
isomers. Aside from the nearly-stable
isotopes, the longest-lived isotopes of
cerium are 139Ce and 144Ce, which have
half-lives on the order of 107 seconds.
An additional eight isotopes have
half-lives greater than 1000 seconds.
All but 16 isotopes of cerium have
half-lives greater than 1 second. The
least stable isotopes whose half-lives
are known are 153Ce, 154Ce and 155Ce.
They have half-lives on the order of
10−7 seconds. All the isotopes lighter
than 140Ce decay by electron capture or
β+ decay. All the isotopes heavier than
140Ce decay by β− decay.
History
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. Cerium was named by
Berzelius after the dwarf planet Ceres,
discovered two years earlier. As
originally isolated, cerium was in the
form of its oxide, and 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 or controlled the mine at Bastnäs,
and had been trying for years to find
out the composition of the abundant
heavy gangue rock, now known as cerite,
that he had in his mine. 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.
When the rare earths were first
discovered, since they were strong bases
like the oxides of calcium or magnesium,
they were thought to be divalent. Thus,
"ceric" cerium was thought to be
trivalent, and the oxidation state ratio
was therefore thought to be 1.5.
Berzelius was annoyed to keep on getting
the correct ratio 1.33.
In the late 1950s, the Lindsay Chemical
Division of American Potash and Chemical
Corporation of West Chicago, Illinois,
then the largest producer of rare earths
in the world, was offering cerium
compounds in two purity ranges,
"commercial" at 94–97% purity, and
"purified", at a reported 99.9+% purity.
In their October 1, 1958 price list,
one-pound quantities of the oxides were
priced at $3.30 or $8.10 respectively
for the two purities; the per-pound
price for 50-pound quantities were
respectively $1.95 or $4.95 for the two
grades. Cerium salts were
proportionately cheaper, reflecting
their lower net content of oxide.
Occurrence
Cerium is the most abundant of the rare
earth elements, making up about 0.0046%
of the Earth's crust by weight. It is
found in a number of minerals including
lanite—(Ca,Ce,La,Y)2(Al,Fe)3(SiO4)3(OH),
monazitePO4, bastnasiteCO3F,
hydroxylbastnasiteCO3(OH,F),
rhabdophanePO4·H2O, zircon, and
synchysite Ca(Ce,La,Nd,Y)(CO3)2F.
Monazite and bastnasite are presently
the two most important sources of
cerium. Large deposits of monazite,
allanite, and bastnasite will supply
cerium, thorium, and other rare-earth
metals for many years to come.
Cerium content in the soil varies
between 2 and 150 parts per million,
with an average of 50 ppm. Seawater
contains 1.5 parts per trillion of
cerium. There is almost no cerium in the
atmosphere.
Production
The mineral mixtures are crushed, ground
and 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 into
their insoluble oxalates. The oxalates
are converted to oxides by annealing.
The oxides are dissolved in nitric acid
that excludes one of the main
components, cerium, whose salts are
insoluble in HNO3. Metallic cerium is
prepared by metallothermic reduction
techniques, such as by reducing cerium
fluoride or chloride with calcium, or by
electrolysis of molten cerous chloride
or other cerous halides. The
metallothermic technique is used to
produce high-purity cerium. The annual
world production of cerium is in the
region of 24,000 mt.
Ce atoms can be doubly encapsulated in
C82 fullerenes and then in carbon
nanotubes. This procedure allows
studying individual Ce atoms. It can
also be used for controlled doping of
carbon peapods with Ce electrons for
electronic applications, owing to the
strong electron affinity of fullerenes.
Applications
A traditional use of cerium was in the
pyrophoric ferrocerium alloy used for
lighter flints.
Because of the high affinity of cerium
to sulfur and oxygen, cerium itself or
as mischmetal is used in various
aluminum alloys, and iron alloys. In
steels, cerium degasifies and can help
reduce sulfides and oxides, and it is a
precipitation hardening agent in
stainless steel. Adding cerium to cast
irons opposes graphitization and
produces a malleable iron. Addition of
3–4% of cerium to magnesium alloys,
along with 0.2–0.6% zirconium, helps
refine the grain and give sound casting
of complex shapes. It also adds heat
resistance to magnesium castings. Cerium
metal is sometimes added to aluminum to
improve aluminum's corrosion resistance.
Cerium alloys are used in permanent
magnets, and in tungsten electrodes for
gas tungsten arc welding. Cerium is used
in carbon-arc lighting, especially in
the motion picture industry.
= CMP=
Ceria is the most widely used compound
of cerium. The main application of ceria
is as a polishing compounds, e.g.
chemical-mechanical planarization. In
this application, ceria dioxide has
replaced other metal oxides for the
production of high quality optical
surfaces.
= Catalysis=
Another important use of cerium oxide is
in a hydrocarbon catalyst in self
cleaning ovens, incorporated into oven
walls and as a petroleum cracking
catalyst in petroleum refining.
Major automotive applications for
cerium(III) oxide are as a catalytic
converter for the oxidation of CO and
NOx emissions in the exhaust gases from
motor vehicles,
= Niche applications=
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. Cerium
sulfide forms a red pigment that stays
stable up to 350 °C. The pigment is a
nontoxic alternative to cadmium sulfide
pigments.
Cerium(IV) oxide is used in incandescent
gas mantles, such as the Welsbach
mantle, where it was combined with
thorium, lanthanum, magnesium, or
yttrium oxides.
Cerium(IV) sulfate is used as a
volumetric oxidizing agent in
quantitative analysis. Ceric ammonium
nitrate is an oxidant in organic
chemistry and in etching electronic
components, and as a primary standard
for quantitative analysis.
Biological role
Cerium can act similar to calcium in
organisms, so accumulates in bones in
small amounts. Cerium is also found in
small amounts in tobacco plants, barley,
and the wood of beech trees. However,
very little cerium accumulates in the
food chain. Human blood contains 0.001
ppm, human bones contain 3 ppm, and
human tissue contains 0.3 ppm of cerium.
There is a total of 40 milligrams of
cerium in a typical 70-kilogram human.
Humans typically consume less than a
milligram per day of cerium. Cerium are
the cofactor for the methanol
dehydrogenase of the methanotrophic
bacterium Methylacidiphilum fumariolicum
SolV.
Cerium salts can stimulate metabolism.
Precautions
Cerium, like all rare-earth metals, is
of low to moderate toxicity. Cerium is a
strong reducing agent and 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 consumed orally, but animals
injected with large doses of cerium have
died due to cardiovascular collapse.
Cerium is more dangerous to aquatic
organisms, on account of being damaging
to cell membranes. Cerium(IV) oxide is a
powerful oxidizing agent at high
temperatures and will react with
combustible organic materials. While
cerium is not radioactive, the impure
commercial grade may contain traces of
thorium, which is weakly radioactive.
See also
Organocerium chemistry
References
External links
It's Elemental – The Element Cerium
Cerium Properties and Applications
Chemistry in its element podcast from
the Royal Society of Chemistry's
Chemistry World: Cerium
