7440-17-7

Usage

Uses

Used in Research:
Rubidium is used in research involving magnetohydrodynamics and thermoionic experiments. The beta-emitter rubidium-87 is used to determine the age of some rocks and minerals, providing valuable information for geological studies.
Used in Electronics:
Rubidium is used in photocells, which convert light into electricity. It is also used as a getter of oxygen in vacuum tubes, helping to maintain a vacuum by reacting with any residual gases. In early radios, TVs, and cathode-ray tubes, rubidium was used as a getter in vacuum tubes.
Used in Timekeeping:
Rubidium gas, when placed in sealed glass cells along with an inert gas, becomes a rubidium-gas cell clock. Due to the consistent and exact frequency of its atoms, it serves as a very accurate timekeeper.
Used in Manufacturing:
Rubidium and selenium are used in the manufacture of photoelectric cells, sometimes referred to as electric eyes. These cells are essential components in various electronic devices and systems.
Used in Medicine:
Rubidium salts have applications in pharmaceuticals as soporifics, sedatives, and for treating epilepsy. The iodide salt is used to treat goiters. Rubidium's unique ability to locate brain tumors is due to its weak radioisotope nature, which allows it to attach itself to diseased tissue rather than healthy tissue, making detection possible.

History and Occurrence

Rubidium was discovered in 1861 by Kirchoff and Bunsen. They observed new lines in the dark red spectral region of a sample extract of mineral lepidolite. The element got its name from the Latin word rubidus, which means dark red. Bunsen later succeeded in preparing metallic rubidium in low yield by heating rubidium hydrogen tartrate with carbon. The metal was obtained in higher yield by Hevesy and later by Beketov, Hevesy electrolyzing a melt of rubidium hydroxide and Beketov reducing the hydroxide with aluminum at red heat. Rubidium is widely distributed in nature. Its abundance in the earth’s crust is estimated to be 90 mg/kg. Rubidium occurs at trace levels in many potassium minerals. Often it is associated with cesium. Some rubidium-containing minerals are lepidolite, leucite, petalite, feldspars, pollucite, beryl, and amazonite. The metal is never found as a major constituent in any mineral. Rubidium also occurs in many rocks such as basalts, granites and clay shales. Rubidium is found in seawater at an average concentration of 0.12 mg/L.

Reactions

Rubidium is a highly reactive metal, more reactive than sodium or potassium. Most reactions are similar to sodium or potassium (see Potassium). The metal ignites spontaneously in air forming oxides. It is coated rapidly with a gray-blue oxide film. It forms four oxides, Rb2O, Rb2O2, Rb2O3, and Rb2O4. It reacts violently with water to form rubidium hydroxide, RbOH: 2Rb + 2H2O → 2RbOH + H2 Reaction with dilute mineral acids can proceed with explosive violence, releasing hydrogen. Rubidium combines with hydrogen and nitrogen forming hydride, RbH and nitride, Rb3N, respectively.

Hazard

As a highly reactive metal, its contact with water or acids can produce violent reactions. Skin contact can cause serious burns.

Hazard

The major hazard is from fire and explosions of the elemental metallic form of rubidium.It must be stored in an inert atmosphere or in kerosene. When rubidium contacts skin, itignites and keeps burning and produces a deep, serious wound. Water and blood just make itreact more vigorously.Many of the compounds of rubidium are toxic and strong irritants to the skin and lungs.It is one of the elements best left to experienced handlers.Very small traces of rubidium are found in the leaves of tobacco, tea, and coffee, as well asin several edible plants, but these radiation traces are harmless when used in moderation.

Isotopes

There are 30 isotopes of rubidium, ranging from Rb-75 to Rb-98. Rb-85 is theonly stable form of rubidium and constitutes 72.17% of all rubidium isotopes found inthe Earth’s crust. Rb-87 is radioactive (a half-life of 4.9×1010 years) and makes up about27.83% of the remainder of rubidium found in the Earth’s crust. All the other 28 isotopes make up a tiny fraction of all the rubidium found on Earth and are radioactive withvery short half-lives.

Origin of Name

Rubidium is named for the Latin word rubidus, meaning “reddish.

History

Rubidium was discovered in 1861 by Bunsen and Kirchhoff 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, carnallite, 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 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. Thirty-five isotopes and isomers of rubidium are known. Naturally occurring rubidium is made of two isotopes, 85Rb and 87Rb. Rubidium-87 is present to the extent of 27.83% in natural rubidium and is a beta emitter with a half-life of 4.9 × 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 armature 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-temperature conductivity of any known ionic crystal. At 20°C 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 $50/g (99.8% pure).

Characteristics

Rubidium is located between potassium and cesium in the first group in the periodic table.It is the second most electropositive alkali element and reacts vigorously and explosively in airor water. If placed on concrete on a sunny day, it would melt and then react violently withmoist air to release hydrogen with enough heat to burn the hydrogen. If a chunk of rubidiummetal is left on a table exposed to the air, it combusts spontaneously. Rubidium must be storedin oil, such as kerosene.

Preparation

Although rubidium metals have been prepared by fused salt electrolysis, the highly reactive nature of the metals complicates the collection step and favors the use of other preparative methods where the metals can be removed in vapor form from the reaction mixture. The oxides, hydroxides, carbonates, halides, sulphates, chromates and nitrates of rubidium have been reduced to the metals by strong reducing metals such as sodium, calcium, magnesium, barium, iron, zirconium, aluminum or silicon at moderately high temperatures. The preferred method, however, involves the reduction of the anhydrous metal chlorides with calcium metal under vacuum. Anhydrous rubidium chloride is mixed with a large excess of calcium chips and heated under vacuum at 700- 800°C. As the chloride is reduced, metal vapors issue from the reaction mixture and are led under the vacuum to a cooler portion of the vessel where they condense and drop into a collection vessel.

Production Methods

Rubidium is recovered from its ore lepidolite or pollucite. Mineral lepidolite is a lithium mica having a composition: KRbLi(OH,F)Al2Si3O10. The ore is opened by fusion with gypsum (potassium sulfate) or with a mixture of barium sulfate and barium carbonate. The fused mass is extracted with hot water to leach out water-soluble alums of cesium, rubidium, and potassium. The solution is filtered to remove insoluble residues. Alums of alkali metals are separated from solution by fractional crystallization. Solubility of rubidium alum or rubidium aluminum sulfate dodecahydrate, RbAl(SO4)2?12H2O falls between potassium and cesium alum.Alternatively, the mineral is opened by prolonged heating with sulfuric acid. Often calcium fluoride (fluorspar) is added for removal of silicon. Alkali metals are converted into water-soluble sulfates. After filtering residual solid, the solution is treated with ammonium or potassium carbonate or carbon dioxide. Lithium precipitates as lithium carbonate. Alkali metal carbonates are converted back to alums and separated by fractional crystallization.Rubidium alum obtained by either method above is decomposed by treatment with alkali solutions for removal of aluminum and sulfate. Aluminum is precipitated as aluminum hydroxide. Addition of barium hydroxide to the filtrate removes sulfate, precipitating barium sulfate. Evaporation of the solution crystallizes rubidium as hydroxide.Rubidium also may be recovered by the chlorostannate method. In this method the alkali metal carbonate solution obtained from the mixed alum is treated with carbon dioxide. Most potassium is precipitated as bicarbonate, KHCO3. Addition of hydrochloric acid converts the carbonates to chlorides. The chlorides are converted to chlorostannates by carefully adding stoichiometric quantities of stannic chloride at pH just below 7:2RbCl + SnCl4 → Rb2SnCl6Cesium chlorostannate, Cs2SnCl6, more insoluble than the rubidium salt, precipitates before any rubidium starts to precipitate. Under such controlled addition of stannic chloride, potassium chloride remains in solution in chloride form. Rubidium chlorostannate complex, on thermal decomposition, forms rubidium chloride, RbCl.Rubidium metal may be obtained from its carbonate, hydroxide or chloride by reduction with magnesium or calcium at high temperatures in the presence of hydrogen:Rb2CO3 + 3Mg → 2Rb + 3MgO +C2RbOH + Mg → 2Rb + Mg(OH)22RbCl + Ca → 2Rb + CaCl2Rubidium is a flammable solid. It is stored in dry hexane, isooctane or other saturated hydrocarbon liquids. Alternatively, the metal may be packaged and stored in well-sealed borosilicate glass ampules or stainless-steel containers under vacuum or an inert atmosphere.

Air & Water Reactions

Tarnishes rapidly upon exposure to air. Reacts violently with water to form corrosive RUBIDIUM hydroxide and hydrogen, a flammable gas. The heat of the reaction usually ignites the hydrogen.

Reactivity Profile

RUBIDIUM METAL is a strong reducing agent. Burns spontaneously in dry oxygen [Mellor 2:468 1946-47]. Readily catches fire in air when molten or with a sulfur vapor [Mellor 2: 469 1946-47]. Causes explosive decomposition of maleic anhydride. [Chem Safety Data Sheet SD-88 1962; Chem. Haz. Info. Series C-71 1960] Burns in chlorine [Mellor 2, Supp. 1:380 1956]. Interaction with mercury is exothermic and may be violent, [Mellor, 1941, Vol. 2, 469].

Health Hazard

Inhalation or contact with vapors, substance or decomposition products may cause severe injury or death. May produce corrosive solutions on contact with water. Fire will produce irritating, corrosive and/or toxic gases. Runoff from fire control may cause pollution.

Fire Hazard

Produce flammable gases on contact with water. May ignite on contact with water or moist air. Some react vigorously or explosively on contact with water. May be ignited by heat, sparks or flames. May re-ignite after fire is extinguished. Some are transported in highly flammable liquids. Runoff may create fire or explosion hazard.

Safety Profile

Moderately toxic by intraperitoneal route. A very reactive alkali metal (more reactive than potassium or cesium). In the body, rubidlum substitutes for potassium as an intracellular ion. The ratio of Rb/K intake is important in the toxicology of rubidium. A ratio above 40% is dangerous. In rats, a failure to gain weight is the first symptom, followed by ataxia and hyperirritabhty. Symptoms include: skin ulcers, poor hair coat, sensitivity, and extreme nervousness leading to convulsions and death. hazard when exposed to heat or flame or by chemical reaction with oxidlzers. Igmtes on contact with air, oxygen, and halogens. A very dangerous fire and explosion RUBIDIUM HYDROXIDE RPZOOO 121 5 Ignites spontaneously on contact with water. Reaction with water, moisture, or steam forms explosive hydrogen gas, whch then ignites. Explodes in contact with liquid bromine. Can react explosively with air, halogens, mercury, nonmetals, vanadium chloride oxide, moisture, acids, oxidizers. Violent reaction with vanadium trichloride oxide (at 60℃C), Cl202, P. Molten rubidium ignites in sulfur vapor and reacts vigorously with carbon. RbOH is more basic than KOH. Storage and handling: Keep under benzene, petroleum, or other liquids not containing gaseous O2. When heated to decomposition it emits toxic fumes of RbzO. See also SODIUM and SODIUM POTASSIUM ALLOY.

InChI:InChI=1/Rb

Synthetic route

monensin(1-) ion → C36H61O11(1-)*Rb(1+) (7440-17-7)

Conditions	Yield
With rubidium ion In ethanol at 25℃; Rate constant;

C36H61O11(1-)*Rb(1+) (7440-17-7) → monensin(1-) ion

Conditions	Yield
In ethanol at 25℃; Rate constant;

Relevant academic research and scientific papers

Rates and Equilibria of Alkali Metal and Silver Ion Complex Formation with Monensin in Ethanol

Measurements are reported on the stability constants and the rates of formation and dissociation of alkali metal and silver complexes of monensin in ethanol.Among the alkali metal complexes the order of stability is Li++>K+>Rb+>Cs+ (with Cs+ ca.Li+).Compared with the neutral antibiotic ionophores, the stability constants of monensin complexes are in general higher and show a much sharper peak selectivity.Dissociation rate constants are very sensitive to cation size and reflect a similar (inverse) variation with cation size of the stability constants; the formation rate constants increase monotonically with increasing cation size from Li+ (9.0E7 M-1s-1) to Cs+(2.5E9 M-1S-1.The silver complex, AgMon, has a stability constant slightly higher than that of the most stable alkali metal complex, NaMon, and a remarkably high formation rate constant, kf=3.5E10 M-1s-1, approaching that of a diffusion-controlled reaction, despite the high solvation energy of Ag+.The kinetic and thermodynamic properties of monensin complexes are compared with those of other naturally occuring and synthetic macrocyclic ionophores.