|Boiling Point: 961°K, 688°C, 1270°F
Melting Point: 312.79°K, 39.64°C, 103.35°F
Electrons Energy Level: 2, 8, 18, 8, 1
Isotopes: 30 + 1 Stable
Heat of Vaporization: 72.216 kJ/mol
Heat of Fusion: 2.192 kJ/mol
Density: 1.63g/cm3 @ 300°K
Specific Heat: 0.363 J/g°K
Atomic Radius: 2.98Å
Ionic Radius: 1.52Å
Electronegativity: 0.82 (Pauling); 0.89 (Allrod Rochow)
Vapor Pressure: 0.000156 Pa @ 39.64°C
|Rubidium (Latin, rubidus,
deepest red) was discovered by its never
seen before red
spectral emission in 1861 by
German chemists Robert Bunsen and Gustav Kirchloff in the mineral Lepidolite
through the use of a spectroscope. It is the seventeenth most abundant element.
Bunsen was eventually able to isolate samples of rubidium metal. Historically, the most important use for Rubidium has been in research and development, primarily in chemical and electronic applications.
Rb is a soft, silvery-white metallic element of the alkali metal group. Rb-87, a naturally occurring isotope, is (slightly) radioactive. Rubidium is very soft and highly reactive, with properties similar to other elements in group 1, like rapid oxidation in air.
Like Sodium Ions, the presence of Potassium Ions in the body is essential for the correct function of many cells. Rubidium (Latin: rubidius = red) is similar in physical and chemical characteristics to Potassium, but much more reactive. Its melting point is so low you could melt it in your hand if you had a fever (39oC). But that would not be a good idea because it would react violently with the moisture in your skin.
Rubidium was once thought to be quite rare but recent discoveries of large deposits indicate that there is plenty to use. However at present it finds only limited application in the manufacture of cathode ray tubes.
Rubidium reacts violently with water and can cause fires. To ensure both safety and purity, this element must be kept under a dry mineral oil, in a vacuum or in an inert atmosphere. Today, most rubidium is obtained as a byproduct of refining lithium.
1s2 2s2p6 3s2p6d10 4s2p6 5s1
Rubidium is the second most electropositive of the stable alkaline 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 Cesium 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, Cesium, Sodium, and Potassium. The element gives a reddish-violet color to a flame, hence its name.
This element is considered to be the 16th most abundant element in the Earth's crust. 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 mettal can be produced by reducing Rubidium Chloride with Calcium among other methods. Rubidium forms at least four oxides: Rb2O, Rb2O2, Rb2O3, RbO2. In 1997 the cost of this metal in small quantities was about $25/gram.
Potential or current uses of Rubidium include:
Rubidium is easily ionized, so it has been considered for use in ion engines for space vehicles (but Cesium 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 wavelenth, 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.
Since it is easily ionized, it might be used as a propellant in ion engines on spacecraft. Recent discoveries of large deposits of rubidium suggest that its usefulness will increase as its properties become better understood.
Rubidium forms a large number of compounds, although none of them has any significant commercial application. Some of the common rubidium compounds are: Rubidium Chloride (RbCl), Rubidium Monoxide (Rb2O), Rubidium Fluoride (RbF), Rubidium Sulfate (Rb2SO4) 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. Rubidium forms four oxides: Rb2O, Rb2O2, Rb2O3, Rb2O4.
|Rubidium Chloride (RbCl)||Rubidium Fluoride (RbF)|
|Rubidium Sulfate (Rb2SO4)||Rubidium Copper Sulfate (Rb2SO4·CuSO4·6H2O)|
There are 30 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%). Normal mixes of Rubidium are radioactive 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.
|Rb72||71.959||< 1.2 ms|
|Rb73||72.95||< 30 ns|
|Rb87||86.9092||4.75E 10 years|
|Rubidium reacts violently with water and can cause fires. To ensure both safety and purity, this element must be kept under a dry mineral oil, in a vacuum or in an inert atmosphere.|
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.
Rubidium crystals are a key component of the subdermal transponders injected into Capt. James T. Kirk and Cmdr. Spock before they beam down to the surface of Ekos in the Star Trek episode "Patterns of Force".
Atomic Radius (Å): 2.98Å
Electrochemical Equivalents: 3.1888 g/amp-hr
Atomic Mass Average: 85.4678
(L. rubidus, deepest red) Discovered in 1861 by Bunsen and Kirchoff in the mineral lepidolite by use of the spectroscope. The element is much more abundant than was thought several years ago. It is now considered to be the 16th most abundant element in the earth's crust. Rubidium occurs in pollucite, leucite, and zinnwaldite, which contains traces up to 1%, in the form of the oxide. It is found in lepidolite to the extent of about 1.5%, and is recovered commercially from this source. Potassium minerals, such as those found at Searles Lake, California, and potassium chloride recovered from the brines in Michigan also contain the element and are commercial sources. It is also found along with cesium in the extensive deposits of pollucite at Bernic Lake, Manitoba. Rubidium can be liquid at room temperature. It is a soft, silvery-white metallic element of the alkali group and is the second most electropositive and alkaline element. It ignites spontaneously in air and reacts violently in water, setting fire to the liberated hydrogen. As with other alkali metals, it forms amalgams with mercury and it alloys with gold, cesium, sodium, and potassium. It colors a flame yellowish violet. Rubidium metal can be prepared by reducing rubidium chloride with calcium, and by a number of other methods. It must be kept under a dry mineral oil or in a vacuum or inert atmosphere. Twenty four isotopes of rubidium are known. Naturally occurring rubidium is made of two isotopes, 85Rb and 87Rb. Rubidium-87 is present to the extent of 27.85% in natural rubidium and is a beta emitter with a half-life of 4.9 x 1010 years. Ordinary rubidium is sufficiently radioactive to expose a photographic film in about 30 to 60 days. Rubidium forms four oxides: Rb2O, Rb2O2, Rb2O3, Rb2O4. Because rubidium can be easily ionized, it has been considered for use in "ion engines" for space vehicles; however, cesium is somewhat more efficient for this purpose. It is also proposed for use as a working fluid for vapor turbines and 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 amature of a generator thereby generating an electric current. Rubidium is used as a getter in vacuum tubes and as a photocell component. It has been used in making special glasses. RbAg4I5 is important, as it has the highest room conductivity of any known ionic crystal. At 20oC its conductivity is about the same as dilute sulfuric acid. This suggests use in thin film batteries and other applications. The present cost in small quantities is about $25/g.
Source: CRC Handbook of Chemistry and Physics, 1913-1995. David R. Lide, Editor in Chief. Author: C.R. Hammond