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= Rubidium =
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Introduction
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Rubidium is a chemical element; it has symbol Rb and atomic number 37.
It is a very soft, whitish-grey solid in the alkali metal group,
similar to potassium and caesium. Rubidium is the first alkali metal
in the group to have a density higher than water. On Earth, natural
rubidium comprises two isotopes: 72% is a stable isotope (85)Rb, and
28% is slightly radioactive (87)Rb, with a half-life of 48.8 billion
years - more than three times as long as the estimated age of the
universe.
German chemists Robert Bunsen and Gustav Kirchhoff discovered rubidium
in 1861 by the newly developed technique, flame spectroscopy. The name
comes from the Latin word , meaning deep red, the color of its
emission spectrum. Rubidium's compounds have various chemical and
electronic applications. Rubidium metal is easily vaporized and has a
convenient spectral absorption range, making it a frequent target for
laser manipulation of atoms. Rubidium is not a known nutrient for any
living organisms. However, rubidium ions have similar properties and
the same charge as potassium ions, and are actively taken up and
treated by animal cells in similar ways.
Physical properties
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Rubidium is a very soft, ductile, silvery-white metal. It has a
melting point of 39.3 °C and a boiling point of 688 °C. It forms
amalgams with mercury and alloys with gold, iron, caesium, sodium, and
potassium, but not lithium (despite rubidium and lithium being in the
same periodic group). Rubidium and potassium show a very similar
purple color in the flame test, and distinguishing the two elements
requires more sophisticated analysis, such as spectroscopy.
Chemical properties
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Rubidium is the second most electropositive of the stable alkali
metals and has a very low first ionization energy of only 403 kJ/mol.
It has an electron configuration of [Kr]5s1 and is photosensitive. Due
to its strong electropositive nature, rubidium reacts explosively with
water to produce rubidium hydroxide and hydrogen gas. As with all the
alkali metals, the reaction is usually vigorous enough to ignite metal
or the hydrogen gas produced by the reaction, potentially causing an
explosion. Rubidium, being denser than potassium, sinks in water,
reacting violently; caesium explodes on contact with water. However,
the reaction rates of all alkali metals depend upon surface area of
metal in contact with water, with small metal droplets giving
explosive rates. Rubidium has also been reported to ignite
spontaneously in air.
Compounds
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Rubidium chloride (RbCl) is probably the most used rubidium compound:
among several other chlorides, it is used to induce living cells to
take up DNA; it is also used as a biomarker, because in nature, it is
found only in small quantities in living organisms and when present,
replaces potassium. Other common rubidium compounds are the corrosive
rubidium hydroxide (RbOH), the starting material for most
rubidium-based chemical processes; rubidium carbonate (Rb2CO3), used
in some optical glasses, and rubidium copper sulfate,
Rb2SO4·CuSO4·6H2O. Rubidium silver iodide (RbAg4I5) has the highest
room temperature conductivity of any known ionic crystal, a property
exploited in thin film batteries and other applications.
Rubidium forms a number of oxides when exposed to air, including
rubidium monoxide (Rb2O), Rb6O, and Rb9O2; rubidium in excess oxygen
gives the superoxide RbO2. Rubidium forms salts with halogens,
producing rubidium fluoride, rubidium chloride, rubidium bromide, and
rubidium iodide.
Isotopes
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Rubidium in the Earth's crust is composed of two isotopes: the stable
85Rb (72.2%) and the radioactive 87Rb (27.8%). Natural rubidium is
radioactive, with specific activity of about 670 Bq/g, enough to
significantly expose a photographic film in 110 days. Thirty
additional rubidium isotopes have been synthesized with half-lives of
less than 3 months; most are highly radioactive and have few uses.
Rubidium-87 has a half-life of years, which is more than three times
the age of the universe of years, making it a primordial nuclide. It
readily substitutes for potassium in minerals, and is therefore fairly
widespread. Rb has been used extensively in dating rocks; 87Rb beta
decays to stable 87Sr. During fractional crystallization, Sr tends to
concentrate in plagioclase, leaving Rb in the liquid phase. Hence, the
Rb/Sr ratio in residual magma may increase over time, and the
progressing differentiation results in rocks with elevated Rb/Sr
ratios. The highest ratios (10 or more) occur in pegmatites. If the
initial amount of Sr is known or can be extrapolated, then the age can
be determined by measurement of the Rb and Sr concentrations and of
the 87Sr/86Sr ratio. The dates indicate the true age of the minerals
only if the rocks have not been subsequently altered (see
rubidium-strontium dating).
Rubidium-82, one of the element's non-natural isotopes, is produced by
electron-capture decay of strontium-82 with a half-life of 25.36 days.
With a half-life of 76 seconds, rubidium-82 decays by positron
emission to stable krypton-82.
Occurrence
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Rubidium is not abundant, being one of 56 elements that combined make
up 0.05% of the Earth's crust; at roughly the 23rd most abundant
element in the Earth's crust it is more abundant than zinc or copper.
It occurs naturally in the minerals leucite, pollucite, carnallite,
and zinnwaldite, which contain as much as 1% rubidium oxide.
Lepidolite contains between 0.3% and 3.5% rubidium, and is the
commercial source of the element. Some potassium minerals and
potassium chlorides also contain the element in commercially
significant quantities.
Seawater contains an average of 125 μg/L of rubidium compared to the
much higher value for potassium of 408 mg/L and the much lower value
of 0.3 μg/L for caesium. Rubidium is the 18th most abundant element in
seawater.
Because of its large ionic radius, rubidium is one of the
"incompatible elements". During magma crystallization, rubidium is
concentrated together with its heavier analogue caesium in the liquid
phase and crystallizes last. Therefore, the largest deposits of
rubidium and caesium are zone pegmatite ore bodies formed by this
enrichment process. Because rubidium substitutes for potassium in the
crystallization of magma, the enrichment is far less effective than
that of caesium. Zone pegmatite ore bodies containing mineable
quantities of caesium as pollucite or the lithium minerals lepidolite
are also a source for rubidium as a by-product.
Two notable sources of rubidium are the rich deposits of pollucite at
Bernic Lake, Manitoba, Canada, and the rubicline found as impurities
in pollucite on the Italian island of Elba, with a rubidium content of
17.5%. Both of those deposits are also sources of caesium.
Production
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Although rubidium is more abundant in Earth's crust than caesium, the
limited applications and the lack of a mineral rich in rubidium limits
the production of rubidium compounds to 2 to 4 tonnes per year.
Several methods are available for separating potassium, rubidium, and
caesium. The fractional crystallization of a rubidium and caesium alum
yields after 30 subsequent steps pure rubidium alum. Two other methods
are reported, the chlorostannate process and the ferrocyanide process.
For several years in the 1950s and 1960s, a by-product of potassium
production called Alkarb was a main source for rubidium. Alkarb
contained 21% rubidium, with the rest being potassium and a small
amount of caesium. Today the largest producers of caesium produce
rubidium as a by-product from pollucite.
History
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Rubidium was discovered in 1861 by Robert Bunsen and Gustav Kirchhoff,
in Heidelberg, Germany, in the mineral lepidolite through flame
spectroscopy. Because of the bright red lines in its emission
spectrum, they chose a name derived from the Latin word , meaning
"deep red".
Rubidium is a minor component in lepidolite. Kirchhoff and Bunsen
processed 150 kg of a lepidolite containing only 0.24% rubidium
monoxide (Rb2O). Both potassium and rubidium form insoluble salts with
chloroplatinic acid, but those salts show a slight difference in
solubility in hot water. Therefore, the less soluble rubidium
hexachloroplatinate (Rb2PtCl6) could be obtained by fractional
crystallization. After reduction of the hexachloroplatinate with
hydrogen, the process yielded 0.51 grams of rubidium chloride (RbCl)
for further studies. Bunsen and Kirchhoff began their first
large-scale isolation of caesium and rubidium compounds with of
mineral water, which yielded 7.3 grams of caesium chloride and 9.2
grams of rubidium chloride. Rubidium was the second element, shortly
after caesium, to be discovered by spectroscopy, just one year after
the invention of the spectroscope by Bunsen and Kirchhoff.
The two scientists used the rubidium chloride to estimate that the
atomic weight of the new element was 85.36 (the currently accepted
value is 85.47). They tried to generate elemental rubidium by
electrolysis of molten rubidium chloride, but instead of a metal, they
obtained a blue homogeneous substance, which "neither under the naked
eye nor under the microscope showed the slightest trace of metallic
substance". They presumed that it was a subchloride (); however, the
product was probably a colloidal mixture of the metal and rubidium
chloride. In a second attempt to produce metallic rubidium, Bunsen was
able to reduce rubidium by heating charred rubidium tartrate. Although
the distilled rubidium was pyrophoric, they were able to determine the
density and the melting point. The quality of this research in the
1860s can be appraised by the fact that their determined density
differs by less than 0.1 g/cm3 and the melting point by less than 1 °C
from the presently accepted values.
The slight radioactivity of rubidium was discovered in 1908, but that
was before the theory of isotopes was established in 1910, and the low
level of activity (half-life greater than 1010 years) made
interpretation complicated. The now proven decay of 87Rb to stable
87Sr through beta decay was still under discussion in the late 1940s.
Rubidium had minimal industrial value before the 1920s. Since then,
the most important use of rubidium is research and development,
primarily in chemical and electronic applications. In 1995,
rubidium-87 was used to produce a Bose-Einstein condensate, for which
the discoverers, Eric Allin Cornell, Carl Edwin Wieman and Wolfgang
Ketterle, won the 2001 Nobel Prize in Physics.
Applications
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Rubidium compounds are sometimes used in fireworks to give them a
purple color. Rubidium has also been considered for use in a
thermoelectric generator using the magnetohydrodynamic principle,
whereby hot rubidium ions are passed through a magnetic field. These
conduct electricity and act like an armature of a generator, thereby
generating an electric current. Rubidium, particularly vaporized 87Rb,
is one of the most commonly used atomic species employed for laser
cooling and Bose-Einstein condensation. Its desirable features for
this application include the ready availability of inexpensive diode
laser light at the relevant wavelength and the moderate temperatures
required to obtain substantial vapor pressures. For cold-atom
applications requiring tunable interactions, 85Rb is preferred for its
rich Feshbach spectrum.
Rubidium has been used for polarizing 3He, producing volumes of
magnetized 3He gas, with the nuclear spins aligned rather than random.
Rubidium vapor is optically pumped by a laser, and the polarized Rb
polarizes 3He through the hyperfine interaction. Such spin-polarized
3He cells are useful for neutron polarization measurements and for
producing polarized neutron beams for other purposes.
The resonant element in atomic clocks utilizes the hyperfine structure
of rubidium's energy levels, and rubidium is useful for high-precision
timing. It is used as the main component of secondary frequency
references (rubidium oscillators) in cell site transmitters and other
electronic transmitting, networking, and test equipment. These
rubidium standards are often used with GNSS to produce a "primary
frequency standard" that has greater accuracy and is less expensive
than caesium standards. Such rubidium standards are often
mass-produced for the telecommunications industry.
Other potential or current uses of rubidium include a working fluid in
vapor turbines, as a getter in vacuum tubes, and as a photocell
component. Rubidium is also used as an ingredient in special types of
glass, in the production of superoxide by burning in oxygen, in the
study of potassium ion channels in biology, and as the vapor in atomic
magnetometers. In particular, 87Rb is used with other alkali metals in
the development of spin-exchange relaxation-free (SERF) magnetometers.
Rubidium-82 is used for positron emission tomography. Rubidium is very
similar to potassium, and tissue with high potassium content will also
accumulate the radioactive rubidium. One of the main uses is
myocardial perfusion imaging. As a result of changes in the
blood-brain barrier in brain tumors, rubidium collects more in brain
tumors than normal brain tissue, allowing the use of radioisotope
rubidium-82 in nuclear medicine to locate and image brain tumors.
Rubidium-82 has a very short half-life of 76 seconds, and the
production from decay of strontium-82 must be done close to the
patient.
Rubidium was tested for the influence on manic depression and
depression. Dialysis patients suffering from depression show a
depletion in rubidium, and therefore a supplementation may help during
depression. In some tests the rubidium was administered as rubidium
chloride with up to 720 mg per day for 60 days.
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Precautions and biological effects
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Rubidium reacts violently with water and can cause fires. To ensure
safety and purity, this metal is usually kept under dry mineral oil or
sealed in glass ampoules in an inert atmosphere. Rubidium forms
peroxides on exposure even to a small amount of air diffused into the
oil, and storage is subject to similar precautions as the storage of
metallic potassium.
Rubidium, like sodium and potassium, almost always has +1 oxidation
state when dissolved in water, even in biological contexts. The human
body tends to treat Rb+ ions as if they were potassium ions, and
therefore concentrates rubidium in the body's intracellular fluid
(i.e., inside cells). The ions are not particularly toxic; a 70 kg
person contains on average 0.36 g of rubidium, and an increase in this
value by 50 to 100 times did not show negative effects in test
persons. The biological half-life of rubidium in humans measures 31-46
days. Although a partial substitution of potassium by rubidium is
possible, when more than 50% of the potassium in the muscle tissue of
rats was replaced with rubidium, the rats died.
Further reading
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* Meites, Louis (1963). 'Handbook of Analytical Chemistry' (New York:
McGraw-Hill Book Company, 1963)
*
External links
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*
* [
http://www.periodicvideos.com/videos/037.htm Rubidium] at 'The
Periodic Table of Videos' (University of Nottingham)
License
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http://creativecommons.org/licenses/by-sa/3.0/
Original Article:
http://en.wikipedia.org/wiki/Rubidium
