rubidium

krypton ← rubidium → strontium K ↑ Rb ↓ Cs 37Rb Periodic table
Appearance
grey white
General properties
Name, symbol, number	rubidium, Rb, 37
Element category	alkali metal
Group, period, block	1, 5, s
Standard atomic weight	85.4678(3) g·mol−1
Electron configuration	[Kr] 5s1
Electrons per shell	2, 8, 18, 8, 1 (Image)
Physical properties
Phase	solid
Density (near r.t.)	1.532 g·cm−3
Liquid density at m.p.	1.46 g·cm−3
Melting point	312.46 K, 39.31 °C, 102.76 °F
Boiling point	961 K, 688 °C, 1270 °F
Critical point	(extrapolated) 2093 K, 16 MPa
Heat of fusion	2.19 kJ·mol−1
Heat of vaporization	75.77 kJ·mol−1
Specific heat capacity	(25 °C) 31.060 J·mol−1·K−1
Vapor pressure
P/Pa 1 10 100 1 k 10 k 100 k at T/K 434 486 552 641 769 958
Atomic properties
Oxidation states	1 (strongly basic oxide)
Electronegativity	0.82 (Pauling scale)
Atomic radius	248 pm
Covalent radius	220±9 pm
Van der Waals radius	303 pm
Miscellanea
Crystal structure	body-centered cubic
Magnetic ordering	paramagnetic[1]
Electrical resistivity	(20 °C) 128Ω·m
Thermal conductivity	(300 K) 58.2 W·m−1·K−1
Speed of sound (thin rod)	(20 °C) 1300 m/s
Young's modulus	2.4 GPa
Bulk modulus	2.5 GPa
Mohs hardness	0.3
Brinell hardness	0.216 MPa
CAS registry number	7440-17-7
Most stable isotopes
Main article: Isotopes of rubidium
iso NA half-life DM DE (MeV) DP 83Rb syn 86.2 d ε - 83Kr γ 0.52, 0.53, 0.55 - 84Rb syn 32.9 d ε - 84Kr β+ 1.66, 0.78 84Kr γ 0.881 - β− 0.892 84Sr 85Rb 72.168% 85Rb is stable with 48 neutrons 86Rb syn 18.65 d β− 1.775 86Sr γ 1.0767 - 87Rb 27.835% 4.88 × 1010 y β− 0.283 87Sr
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Rubidium (pronounced /rʊˈbɪdiəm/, roo-BID-ee-əm) 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.

Rubidium is very soft and highly reactive, with properties similar to other elements in group 1, such as very rapid oxidation in air. Its compounds have 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 known to be necessary for any living organisms. However, like caesium, rubidium ions are handled by living organisms in a manner similar to potassium: it is actively taken up by plants and by living animal cells.

Rubidium has one stable isotope,⁸⁵Rb. The isotope ⁸⁷Rb which composes almost 28% of naturally occurring rubidium is slightly radioactive, with a half-life of 49 billion years—more than three times longer than the estimated age of the universe.

Characteristics

Rubidium is the second most electropositive of the stable alkali elements and liquefies at a temperature of 39.3 °C (102.7 °F). 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.

History

Rubidium (Latin: rubidus, deepest red) was discovered in 1861 by Robert Bunsen and Gustav Kirchhoff in the mineral lepidolite through the use of a spectroscope.^[2] Processing 150 kg of lepidolite yielded only a few grams for analysis. Rubidium metal was first produced by the reaction of rubidium chloride with potassium by Bunsen.

Occurrence

Rubidium is about the twenty-third most abundant element in the Earth's crust, roughly as abundant as zinc and rather more common than copper.^[3] It occurs naturally in the minerals leucite, pollucite, carnallite and zinnwaldite, which contain 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 (also a source of the related element caesium).

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

Rubidium-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; ⁸⁷Rb 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 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 for a more detailed discussion.

Uses and applications

Rubidium had minimal industrial use until the 1930s. Historically, the most important use for rubidium has been in research and development, primarily in chemical and electronic applications.

In 1999 rubidium-87 was used to make a Bose-Einstein condensate^[4], for which the discoverers won the 2001 Nobel Prize in Physics^[5].

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.^[6]

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 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.

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.^[7] Spin-polarized ³He cells are becoming popular for neutron polarization measurements and for producing polarized neutron beams for other purposes.^[8]

Rubidium is the primary compound used in secondary frequency references (rubidium oscillators) to maintain frequency accuracy in cell site transmitters and other electronic transmitting, networking and test equipment. Rubidium references are often used with GPS to produce a "primary frequency standard" that has greater accuracy but is less expensive than caesium standards. Rubidium references such as the LPRO series from Datum were mass-produced for the Telecom industry. The general life expectancy is 10 years or better for most designs.

Other potential or current uses of rubidium include:

• 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 is used to locate brain tumors, due to its slight radioactivity.^[3]
• 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.^[9]

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 form if rubidium metal is 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·CuSO₄·6H₂O. A compound of rubidium, silver and iodine, RbAg₄I₅, has interesting electrical characteristics and might be useful in thin film batteries.^[10]

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, and in practice is usually sealed in glass ampoules in an inert atmosphere. Rubidium forms peroxides on exposure to even air diffusing into oil, and is thus subject to some of the same peroxide precautions as storage of metallic potassium.

Biological effects

Rubidium, like sodium and potassium, is almost always in its +1 oxidation state when dissolved in water, and this includes all biological systems. 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, and are relatively quickly removed in the sweat and urine. As a result of changes in the blood brain barrier in brain tumors, rubidium collects more in brain tumors than normal brain tissue, allowing short-lived radioisotopes of rubidium to be used in nuclear medicine to locate and image brain tumors.

References

Sources

• Louis Meites, Handbook of Analytical Chemistry (New York: McGraw-Hill Book Company, 1963)
• Daniel A. Steck. "Rubidium-87 D Line Data". Los Alamos National Laboratory (technical report LA-UR-03-8638). http://george.ph.utexas.edu/~dsteck/alkalidata/rubidium87numbers.pdf. 

External links

Wikimedia Commons has media related to: Rubidium

Look up rubidium in Wiktionary, the free dictionary.

• WebElements.com – Rubidium