Polonium is a chemical element with symbol Po and atomic number
84, discovered in 1898 by Marie Curie and Pierre Curie. A rare and
highly radioactive element with no stable isotopes, polonium is
chemically similar to bismuth and tellurium, and it occurs in
uranium ores. Applications of polonium are few. They include
heaters in space probes, antistatic devices, and sources of
neutrons and alpha particles. Because of its position in the
periodic table, polonium is sometimes classified as a metalloid.[3]
Other sources say that on the basis of its properties and behavior,
it is "unambiguously a metal."Also tentatively called "radium F",
polonium was discovered by Marie and Pierre Curie in 1898, and was
named after Marie Curie's native land of Poland (Latin: Polonia).
Polonium may be the first element named to highlight a political
controversy.This element was the first one discovered by the Curies
while they were investigating the cause of pitchblende
radioactivity. Pitchblende, after removal of the radioactive
elements uranium and thorium, was more radioactive than the uranium
and thorium combined. This spurred the Curies to search for
additional radioactive elements. They first separated out polonium
from pitchblende in July 1898, and five months later, also isolated
radium. The Atomic Energy Commission and the Manhattan Project
funded human experiments using polonium on 5 people at the
University of Rochester between 1943 and 1947. The people were
administered between 9 and 22 microcuries (330 and 810kBq) of
polonium to study its excretion. Polonium has 33 known isotopes,
all of which are radioactive. They have atomic masses that range
from 188 to 220 u. 210Po (half-life 138.376days) is the most widely
available. The longer-lived 209Po (half-life 125.2 3.3 years,
longest-lived of all polonium isotopes)[2] and 208Po (half-life
2.9years) can be made through the alpha, proton, or deuteron
bombardment of lead or bismuth in a cyclotron.[5]210Po is an alpha
emitter that has a half-life of 138.4days; it decays directly to
its stable daughter isotope, 206Pb. A milligram (5curies) of 210Po
emits about as many alpha particles per second as 5grams of
226Ra.[6] A few curies (1curie equals 37gigabecquerels, 1Ci =
37GBq) of 210Po emit a blue glow which is caused by excitation of
surrounding air.About one in 100,000alpha emissions causes an
excitation in the nucleus which then results in the emission of a
gamma ray with a maximum energy of 803keVPolonium is a radioactive
element that exists in two metallic allotropes. The alpha form is
the only known example of a simple cubic crystal structure in a
single atom basis, with an edge length of 335.2 picometers; the
beta form is rhombohedral.[9][10][11] The structure of polonium has
been characterized by X-ray diffraction[12][13] and electron
diffraction.[14]210Po (in common with 238Pu) has the ability to
become airborne with ease: if a sample is heated in air to 55C
(131F), 50% of it is vaporized in 45 hours to form diatomic Po2
molecules, even though the melting point of polonium is 254C (489F)
and its boiling point is 962C (1,764F).[15][16][1] More than one
hypothesis exists for how polonium does this; one suggestion is
that small clusters of polonium atoms are spalled off by the alpha
decay.ChemistryThe chemistry of polonium is similar to that of
tellurium and bismuth. Polonium dissolves readily in dilute acids,
but is only slightly soluble in alkalis. Polonium solutions are
first colored in pink by the Po2+ ions, but then rapidly become
yellow because alpha radiation from polonium ionizes the solvent
and converts Po2+ into Po4+. This process is accompanied by
bubbling and emission of heat and light by glassware due to the
absorbed alpha particles; as a result, polonium solutions are
volatile and will evaporate within days unless
sealed.[17][18]CompoundsPolonium has no common compounds, and
almost all of its compounds are synthetically created; more than 50
of those are known.[19] The most stable class of polonium compounds
are polonides, which are prepared by direct reaction of two
elements. Na2Po has the antifluorite structure, the polonides of
Ca, Ba, Hg, Pb and lanthanides form a NaCl lattice, BePo and CdPo
have the wurtzite and MgPo the nickel arsenide structure. Most
polonides decompose upon heating to about 600C, except for HgPo
that decomposes at ~300C and the lanthanide polonides, which do not
decompose but melt at temperatures above 1000C. For example PrPo
melts at 1250C and TmPo at 2200C.[20] PbPo is one of the very few
naturally occurring polonium compounds, as polonium alpha decays to
form lead.[21]Polonium hydride (PoH2) is a volatile liquid at room
temperature prone to dissociation.[20] The two oxides PoO2 and PoO3
are the products of oxidation of polonium.[22]Halides of the
structure PoX2, PoX4 and PoX6 are known. They are soluble in the
corresponding hydrogen halides, i.e., PoClX in HCl, PoBrX in HBr
and PoI4 in HI.[23] Polonium dihalides are formed by direct
reaction of the elements or by reduction of PoCl4 with SO2 and with
PoBr4 with H2S at room temperature. Tetrahalides can be obtained by
reacting polonium dioxide with HCl, HBr or HI.[24]Other polonium
compounds include acetate, bromate, carbonate, citrate, chromate,
cyanide, formate, hydroxide, nitrate, selenate, monosulfide,
sulfate and disulfate.Krypton (from Greek: kryptos "the hidden
one") is a chemical element with symbolKr and atomic number36. It
is a member of group 18 (noble gases) elements. A colorless,
odorless, tasteless noble gas, krypton occurs in trace amounts in
the atmosphere, is isolated by fractionally distilling liquefied
air, and is often used with other rare gases in fluorescent lamps.
Krypton is inert for most practical purposes.A whitish, largely
inert gaseous element used chiefly in gas discharge lamps and
fluorescent lamps. Solid krypton is a white crystalline substance
with a face-centered cubic structure which is common to all the
'rare gases'. The Atomic Number of this element is 36 and the
Element Symbol is Kr.Krypton, like the other noble gases, can be
used in lighting and photography. Krypton light has a large number
of spectral lines, and krypton's high light output in plasmas
allows it to play an important role in many high-powered gas lasers
(krypton ion and excimer lasers), which pick out one of the many
spectral lines to amplify. There is also a specific krypton
fluoride laser. The high power and relative ease of operation of
krypton discharge tubes caused (from 1960 to 1983) the official
length of a meter to be defined in terms of the wavelength of the
605nm (orange) spectral line of krypton-86.Krypton was discovered
in Britain in 1898 by Sir William Ramsay, a Scottish chemist, and
Morris Travers, an English chemist, in residue left from
evaporating nearly all components of liquid air. Neon was
discovered by a similar procedure by the same workers just a few
weeks later.[5] William Ramsay was awarded the 1904 Nobel Prize in
Chemistry for discovery of a series of noble gases, including
krypton.In 1960, the International Conference on Weights and
Measures defined the meter as 1,650,763.73 wavelengths of light
emitted by the krypton-86 isotope.[6][7] This agreement replaced
the 1889 international prototype meter located in Paris, which was
a metal bar made of a platinum-iridium alloy (one of a series of
standard meter bars, originally constructed to be one ten-millionth
of a quadrant of the Earth's polar circumference). This also
obsoleted the 1927 definition of the ngstrm based on the red
cadmium spectral line,[8] replacing it with 1 = 1010m. The
krypton-86 definition lasted until the October 1983 conference,
which redefined the meter as the distance that light travels in a
vacuum during 1/299,792,458s.[9][10][11]CharacteristicsKrypton is
characterized by several sharp emission lines (spectral signatures)
the strongest being green and yellow.[12] It is one of the products
of uranium fission.[13] Solidified krypton is white and crystalline
with a face-centered cubic crystal structure, which is a common
property of all noble gases (except helium, with a hexagonal
close-packed crystal structure).IsotopesMain article: Isotopes of
kryptonNaturally occurring krypton is made of six stable isotopes.
In addition, about thirty unstable isotopes and isomers are
known.[14] 81Kr, the product of atmospheric reactions, is produced
with the other naturally occurring isotopes of krypton. Being
radioactive, it has a half-life of 230,000 years. Krypton is highly
volatile when it is near surface waters but 81Kr has been used for
dating old (50,000800,000 years) groundwater.[15]85Kr is an inert
radioactive noble gas with a half-life of 10.76 years. It is
produced by the fission of uranium and plutonium, such as in
nuclear bomb testing and nuclear reactors. 85Kr is released during
the reprocessing of fuel rods from nuclear reactors. Concentrations
at the North Pole are 30% higher than at the South Pole due to
convective mixing.[16]ChemistryLike the other noble gases, krypton
is highly chemically unreactive. However, following the first
successful synthesis of xenon compounds in 1962, synthesis of
krypton difluoride (KrF2) was reported in 1963. In fact, before the
1960s, no noble gas compounds had been discovered.[17] Under
extreme conditions, krypton reacts with fluorine to form KrF2
according to the following equation:Kr + F2 KrF2In the same year,
KrF4 was reported by Grosse, et al.,[18] but was subsequently shown
to be a mistaken identification.[19] There are also unverified
reports of a barium salt of a krypton oxoacid.[20] ArKr+ and KrH+
polyatomic ions have been investigated and there is evidence for
KrXe or KrXe+.[21]Compounds with krypton bonded to atoms other than
fluorine have also been discovered. The reaction of KrF2 with
B(OTeF5)3 produces an unstable compound, Kr(OTeF5)2, that contains
a krypton-oxygen bond. A krypton-nitrogen bond is found in the
cation [HCNKrF]+, produced by the reaction of KrF2 with
[HCNH]+[AsF6] below 50C.[22][23] HKrCN and HKrCCH (krypton
hydride-cyanide and hydrokryptoacetylene) were reported to be
stable up to 40 K.Lead (/ld/) is a chemical element in the carbon
group with symbolPb (from Latin: plumbum) and atomic number82. Lead
is a soft, malleable and heavy post-transition metal. Metallic lead
has a bluish-white color after being freshly cut, but it soon
tarnishes to a dull grayish color when exposed to air. Lead has a
shiny chrome-silver luster when it is melted into a liquid. It is
also one of the heaviest non-radioactive elements.Lead is used in
building construction, lead-acid batteries, bullets and shot,
weights, as part of solders, pewters, fusible alloys, and as a
radiation shield. Lead has the highest atomic number of all of the
stable elements, although the next higher element, bismuth, has one
isotope with a half-life that is so long (over one billion times
the estimated age of the universe) that it can be considered
stable. Lead's four stable isotopes have 82 protons, a magic number
in the nuclear shell model of atomic nuclei. The isotope lead-208
also has 126 neutrons, another magic number, and is hence double
magic, a property that grants it enhanced stability: lead-208 is
the heaviest known stable isotope.If ingested, lead is poisonous to
animals and humans, damaging the nervous system and causing brain
disorders. Excessive lead also causes blood disorders in mammals.
Like the element mercury, another heavy metal, lead is a neurotoxin
that accumulates both in soft tissues and the bones. Lead poisoning
has been documented from ancient Rome, ancient Greece, and ancient
China.
CharacteristicsLead is a bright and silvery metal with a very
slight shade of blue in a dry atmosphere.[1] Upon contact with air,
it begins to tarnish by forming a complex mixture of compounds
depending on the conditions. The color of the compounds can vary.
The tarnish layer can contain significant amounts of carbonates and
hydroxycarbonates.[2][3] Lead's characteristic properties include
high density, softness, ductility and malleability, poor electrical
conductivity compared to other metals, high corrosion resistance,
and ability to react with organic chemicals.[1]Various traces of
other metals change its properties significantly: the addition of
small amounts of antimony or copper to lead increases the alloy's
hardness and improves corrosion resistance to sulfuric acid.[1]
Some other metals, such as cadmium, tin, and tellurium, also
improve hardness and fight metal fatigue. Sodium and calcium also
have this ability, but they reduce the alloy's chemical
stability.[1] Finally, zinc and bismuth simply impair the corrosion
resistance (0.1% bismuth content is the industrial usage
threshold).[1] Conversely, lead impurities mostly worsen the
quality of industrial materials, although there are exceptions: for
example, small amounts of lead improve the ductility of
steel.[1]Lead has only one common allotrope, which is face-centered
cubic, with the leadlead distance being 349pm.[4] At 327.5C
(621.5F),[5] lead melts; the melting point is above that of tin
(232C, 449.5F),[5] but significantly below that of germanium (938C,
1721F).[6] The boiling point of lead is 1749C (3180F),[7] which is
below those of both tin (2602C, 4716F)[5] and germanium (2833C,
5131F).[6] Densities increase down the group: the Ge and Sn values
(5.23[8] and 7.29gcm3,[9] respectively) are significantly below
that of lead: 11.32gcm3.[8]A lead atom has 82 electrons, having an
electronic configuration of [Xe]4f145d106s26p2. In its compounds,
lead (unlike the other group 14 elements) most commonly loses its
two and not four outermost electrons, becoming lead(II) ions, Pb2+.
Such unusual behavior is rationalized by considering the inert pair
effect, which occurs because of the stabilization of the 6s-orbital
due to relativistic effects, which are stronger closer to the
bottom of the periodic table.[10] Tin shows a weaker such effect:
tin(II) is still a reducer.[10]The figures for electrode potential
show that lead is only slightly easier to oxidize than hydrogen.
Lead thus can dissolve in acids, but this is often impossible due
to specific problems (such as the formation of insoluble
salts).[11] Powdered lead burns with a bluish-white flame. As with
many metals, finely divided powdered lead exhibits
pyrophoricity.[12] Toxic fumes are released when lead is
burned.Isotopes[edit]Main article: Isotopes of leadLead occurs
naturally on Earth exclusively in the form of four isotopes:
lead-204, -206, -207, and -208.[13] All four can be radioactive as
the hypothetical alpha decay of any would be exothermic, but the
lower half-life limit has been put only for lead-204: over
1.41017years.[14] This effect is, however, so weak that natural
lead poses no radiation hazard. Three isotopes are also found in
three of the four major decay chains: lead-206, -207 and -208 are
final decay products of uranium-238, uranium-235, and thorium-232,
respectively. Since the amounts of them in nature depend also on
other elements' presence, the isotopic composition of natural lead
varies by sample: in particular, the relative amount of lead-206
varies between 20.84% and 27.78%.[13]Aside from the stable ones,
thirty-four radioisotopes have been synthesized: they have mass
numbers of 178215.[14] Lead-205 is the most stable radioisotope of
lead, with a half-life of over 107years. 47 nuclear isomers
(long-lived excited nuclear states), corresponding to 24 lead
isotopes, have been characterized. The most long-lived isomer is
lead-204m2 (half-life of about 1.1hours).[14]Chemical
reactivity[edit]Lead is classified as a post-transition metal and
is also a member of the carbon group. Massive lead forms a
protective oxide layer, but finely powdered highly purified lead
can ignite in air. Melted lead is oxidized in air to lead monoxide.
All chalcogens oxidize lead upon heating.[15]Fluorine does not
oxidize cold lead. Hot lead can be oxidized, but the formation of a
protective halide layer lowers the intensity of the reaction above
100C (210F). The reaction with chlorine is similar: thanks to the
chloride layer, lead persistence against chlorine surpasses that of
copper or steel up to 300C (570F).[15]Water in the presence of
oxygen attacks lead to start an accelerating reaction. The presence
of carbonates or sulfates results in the formation of insoluble
lead salts, which protect the metal from corrosion. So does carbon
dioxide, as the insoluble lead carbonate is formed; however, an
excess of the gas leads to the formation of the soluble
bicarbonate, which makes the use of lead pipes dangerous.[11] Lead
dissolves in organic acids (in the presence of oxygen) and
concentrated (80%) sulfuric acid thanks to complexation; however,
it is only weakly affected by hydrochloric acid and is stable
against hydrofluoric acid, as the corresponding halides are weakly
soluble. Lead also dissolves in quite concentrated alkalis (10%)
because of the amphoteric character and solubility of
plumbites.[11]Compounds[edit]Lead compounds exist mainly in two
main oxidation states, +2 and +4. The former is more common.
Inorganic lead(IV) compounds are typically strong oxidants or exist
only in highly acidic solutions.[
