rubidium n : a soft silvery metallic element of the alkali metal group; burns in air and reacts violently in water; occurs in carnallite and lepidolite and pollucite [syn: Rb, atomic number 37]

see Rubidium

English

Etymology

A word derived by German chemist R. W. Bunsen in 1861, from rubidus because its spectrum has two red lines.

Pronunciation

• a UK /ˌɹʊuˈbɪd.iː.ʌm/ /%ru:"bId.i:.Vm/

Noun

1. A metallic chemical element (symbol Rb) with an atomic number of 37.

Derived terms

• iodrubidium
• rubidium auride
• rubidium chloride
• rubidium hydroxide
• rubidium oxide
• rubidium-strontium

Related terms

• rubidic

Translations

References

External links

For etymology and more information refer to: http://elements.vanderkrogt.net/elem/rb.html (A lot of the translations were taken from that site with permission from the author)

Finnish

Noun

1. rubidium

Melting point 39.31˚ C Rubidium (, /rəˈbɪdiəm/) is a chemical element with the symbol Rb and atomic number 37. Rb is a soft, silvery-white metallic element of the alkali metal group. Rb-87, a naturally occurring isotope, is very slightly radioactive, with a half-life of 4.88×1010 years, which, at 49 billion years, is far longer than the estimated age of the universe. Rubidium is very soft and highly reactive, with properties similar to other elements in group 1, like rapid oxidation in air.

Notable characteristics

Rubidium is the second most electropositive of the stable alkali elements and liquefies at high ambient temperature (102.7 °F = 39.3 °C). Like other group 1 elements this metal reacts violently in water. In common with potassium and caesium this reaction is usually vigorous enough to ignite the liberated hydrogen. Rubidium has also been reported to ignite spontaneously in air. Also like other alkali metals, it forms amalgams with mercury and it can form alloys with gold, caesium, sodium, and potassium. The element gives a reddish-violet color to a flame, hence its name.

Uses

Potential or current uses of rubidium include:

• A Bose-Einstein condensate.
• A working fluid in vapor turbines.
• A getter in vacuum tubes.
• A photocell component.
• The resonant element in atomic clocks. This is due to the hyperfine structure of Rubidium's energy levels.
• An ingredient in special types of glass.
• The production of superoxide by burning in oxygen.
• The study of potassium ion channels in biology.
• Rubidium vapor has been used to make atomic magnetometers. 87Rb is currently being used, with other alkali metals, in the development of spin-exchange relaxation-free (SERF) magnetometers.

Rubidium is easily ionized, so it has been considered for use in ion engines for space vehicles (but caesium and xenon are more efficient for this purpose).

Rubidium compounds are sometimes used in fireworks to give them a purple color.

RbAg4I5 has the highest room temperature conductivity of any known ionic crystal. This property could be useful in thin film batteries and in other applications.

Rubidium has also been considered for use in a thermoelectric generator using the magnetohydrodynamic principle, where rubidium ions are formed by heat at high temperature and passed through a magnetic field. These conduct electricity and act like an armature of a generator thereby generating an electric current.

Rubidium, particularly 87Rb, in the form of vapor, 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.

Rubidium has been used for polarizing 3He (that is, producing volumes of magnetized 3He gas, with the nuclear spins aligned toward a particular direction in space, rather than randomly). Rubidium vapor is optically pumped by a laser and the polarized Rb polarizes 3He by the hyperfine interaction. Spin-polarized 3He cells are becoming popular for neutron polarization measurements and for producing polarized neutron beams for other purposes.

History

Rubidium (L rubidus, deepest red) was discovered in 1861 by Robert Bunsen and Gustav Kirchhoff in the mineral lepidolite through the use of a spectroscope. The extraction of 150 kg of lepidolite yielded only a few grams for analysis. The first rubidium metal was produced by the reaction of rubidium chloride with potassium by Bunsen. However, this element had minimal industrial use until the 1920s. Historically, the most important use for rubidium has been in research and development, primarily in chemical and electronic applications.

In the year 2000 rubidium-87 was used to make a Bose-Einstein condensate, for which the discoverers won the 2001 Nobel Prize in Physics.

Occurrence

Rubidium is about the sixteenth most abundant metal in the Earth's crust, roughly as abundant as zinc and rather more common than copper. It occurs naturally in the minerals leucite, pollucite, and zinnwaldite, which contains traces of up to 1% of its oxide. Lepidolite contains 1.5% rubidium and this is the commercial source of the element. Some potassium minerals and potassium chlorides also contain the element in commercially significant amounts. One notable source is also in the extensive deposits of pollucite at Bernic Lake, Manitoba.

Rubidium metal can be produced by reducing rubidium chloride with calcium among other methods. In 1997 the cost of this metal in small quantities was about US$ 25/gram.

Isotopes

There are 24 isotopes of rubidium known with naturally occurring rubidium being composed of just two isotopes; Rb-85 (72.2%) and the radioactive Rb-87 (27.8%). Natural rubidium is radioactive with specific activity of about 670 Bq/g, enough to fog photographic film in approximately 30 to 60 days.

Rb-87 has a half-life of 4.88×1010 years. It readily substitutes for potassium in minerals, and is therefore fairly widespread. Rb has been used extensively in dating rocks; Rb-87 decays to stable strontium-87 by emission of a negative beta particle. During fractional crystallization, Sr tends to become concentrated in plagioclase, leaving Rb in the liquid phase. Hence, the Rb/Sr ratio in residual magma may increase over time, resulting in rocks with increasing Rb/Sr ratios with increasing differentiation. Highest ratios (10 or higher) occur in pegmatites. If the initial amount of Sr is known or can be extrapolated, the age can be determined by measurement of the Rb and Sr concentrations and the Sr-87/Sr-86 ratio. The dates indicate the true age of the minerals only if the rocks have not been subsequently altered. See Rubidium-Strontium dating for a more detailed discussion.

Compounds

Rubidium chloride is probably the most-used rubidium compound; it is used in biochemistry to induce cells to take up DNA, and as a biomarker since it is readily taken up to replace potassium, and does not normally occur in living organisms. Rubidium hydroxide is the starting material for most rubidium-based chemical processes; rubidium carbonate is used in some optical glasses.

Rubidium has a number of oxides, including Rb6O and Rb9O2 which appear if rubidium metal is left exposed to air; the final product of reacting with oxygen is the superoxide RbO2. Rubidium forms salts with most anions. Some common rubidium compounds are rubidium chloride (RbCl), rubidium monoxide (Rb2O) and rubidium copper sulfate Rb2SO4·CuSO4·6H20). A compound of rubidium, silver and iodine, RbAg4I5, has interesting electrical characteristics and might be useful in thin film batteries.

Precautions

Rubidium reacts violently with water and can cause fires. To ensure both health and safety and purity, this element must be kept under a dry mineral oil, in a vacuum or in an inert atmosphere.

Biological effects

Rubidium, like sodium and potassium, is almost always in its +1 oxidation state. The human body tends to treat Rb+ ions as if they were potassium ions, and therefore concentrates rubidium in the body's electrolytic fluid. The ions are not particularly toxic, and are relatively quickly removed in the sweat and urine. However, taken in excess it can be dangerous.

See also

• Rubidium compounds

References

Sources

• Los Alamos National Laboratory – Rubidium
• Louis Meites, Handbook of Analytical Chemistry (New York: McGraw-Hill Book Company, 1963)

External links

• WebElements.com – Rubidium