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.^[8] 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 ⁸⁵Rb, and 28% is slightly radioactive ⁸⁷Rb, with a half-life of 48.8 billion years—more than three times as long as the estimated age of the universe.

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Rubidium, ₃₇Rb

Rubidium
Pronunciation	/ruːˈbɪdiəm/ ​(roo-BID-ee-əm)
Appearance	grey white
Standard atomic weight Ar°(Rb)
	85.4678±0.0003[1] 85.468±0.001 (abridged)[2]
Rubidium in the periodic table
Hydrogen Helium Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson K ↑ Rb ↓ Cs krypton ← rubidium → strontium
Atomic number (Z)	37
Group	group 1: hydrogen and alkali metals
Period	period 5
Block	s-block
Electron configuration	[Kr] 5s1
Electrons per shell	2, 8, 18, 8, 1
Physical properties
Phase at STP	solid
Melting point	312.45 K ​(39.30 °C, ​102.74 °F)
Boiling point	961 K ​(688 °C, ​1270 °F)
Density (at 20° C)	1.534 g/cm3[3]
when liquid (at m.p.)	1.46 g/cm3
Triple point	312.41 K, ​? kPa[4]
Critical point	2093 K, 16 MPa (extrapolated)[4]
Heat of fusion	2.19 kJ/mol
Heat of vaporization	69 kJ/mol
Molar heat capacity	31.060 J/(mol·K)
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, +1 (a strongly basic oxide)
Electronegativity	Pauling scale: 0.82
Ionization energies	1st: 403 kJ/mol 2nd: 2632.1 kJ/mol 3rd: 3859.4 kJ/mol
Atomic radius	empirical: 248 pm
Covalent radius	220±9 pm
Van der Waals radius	303 pm
Spectral lines of rubidium
Other properties
Natural occurrence	primordial
Crystal structure	​body-centered cubic (bcc) (cI2)
Lattice constant	a = 569.9 pm (at 20 °C)[3]
Thermal expansion	85.6×10−6/K (at 20 °C)[3]
Thermal conductivity	58.2 W/(m⋅K)
Electrical resistivity	128 nΩ⋅m (at 20 °C)
Magnetic ordering	paramagnetic[5]
Molar magnetic susceptibility	+17.0×10−6 cm3/mol (303 K)[6]
Young's modulus	2.4 GPa
Bulk modulus	2.5 GPa
Speed of sound thin rod	1300 m/s (at 20 °C)
Mohs hardness	0.3
Brinell hardness	0.216 MPa
CAS Number	7440-17-7
History
Discovery	Robert Bunsen and Gustav Kirchhoff (1861)
First isolation	George de Hevesy
Isotopes of rubidium v e
Main isotopes[7] Decay abun­dance half-life (t1/2) mode pro­duct 82Rb synth 1.2575 m β+ 82Kr 83Rb synth 86.2 d ε 83Kr γ – 84Rb synth 32.9 d ε 84Kr β+ 84Kr γ – β− 84Sr 85Rb 72.2% stable 86Rb synth 18.7 d β− 86Sr γ – 87Rb 27.8% 4.923×1010 y β− 87Sr
Category: Rubidium view talk edit | references

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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 rubidus, 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.^[9] 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.

Rubidium is a very soft, ductile, silvery-white metal.^[10] It is the second most electropositive of the stable alkali metals and melts at a temperature of 39.3 °C (102.7 °F). Like other alkali metals, rubidium metal reacts violently with water. As with potassium (which is slightly less reactive) and caesium (which is slightly more reactive), this reaction is usually vigorous enough to ignite the hydrogen gas it produces. Rubidium has also been reported to ignite spontaneously in air.^[10] It forms amalgams with mercury and alloys with gold, iron, caesium, sodium, and potassium, but not lithium (even though rubidium and lithium are in the same group).^[11]

Rubidium has a very low ionization energy of only 406 kJ/mol.^[12] 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.^{[citation needed]}

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.^[23] Several methods are available for separating potassium, rubidium, and caesium. The fractional crystallization of a rubidium and caesium alum (Cs,Rb)Al(SO₄)₂·12H₂O yields after 30 subsequent steps pure rubidium alum. Two other methods are reported, the chlorostannate process and the ferrocyanide process.^[23][30]

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.^[31] Today the largest producers of caesium produce rubidium as a by-product from pollucite.^[23]

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 rubidus, meaning "deep red".^[32][33]

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 44,000 litres (12,000 US gal) of mineral water, which yielded 7.3 grams of caesium chloride and 9.2 grams of rubidium chloride.^[32][33] 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.^[34]

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).^[32] 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 (Rb
₂Cl); however, the product was probably a colloidal mixture of the metal and rubidium chloride.^[35] 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.^[36]

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.^[37][38]

Rubidium had minimal industrial value before the 1920s.^[39] 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,^[40] for which the discoverers, Eric Allin Cornell, Carl Edwin Wieman and Wolfgang Ketterle, won the 2001 Nobel Prize in Physics.^[41]

Rubidium compounds are sometimes used in fireworks to give them a purple color.^[42] 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.^[43] 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.^[44][45] For cold-atom applications requiring tunable interactions, ⁸⁵Rb is preferred for its rich Feshbach spectrum.^[46]

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.^[47] Such spin-polarized ³He cells are useful for neutron polarization measurements and for producing polarized neutron beams for other purposes.^[48]

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.^[49][50] Such rubidium standards are often mass-produced for the telecommunication industry.^[51]

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.^[52] 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.^[53] In particular, ⁸⁷Rb is used with other alkali metals in the development of spin-exchange relaxation-free (SERF) magnetometers.^[53]

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.^[54] 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.^[55]

Rubidium was tested for the influence on manic depression and depression.^[56][57] Dialysis patients suffering from depression show a depletion in rubidium, and therefore a supplementation may help during depression.^[58] In some tests the rubidium was administered as rubidium chloride with up to 720 mg per day for 60 days.^[59][60]

Quick Facts Hazards ...

Rubidium

Hazards
GHS labelling:
Pictograms
Signal word	Danger
Hazard statements	H260, H314
Precautionary statements	P223, P231+P232, P280, P305+P351+P338, P370+P378, P422[61]
NFPA 704 (fire diamond)	3 4 2 W

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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.^[62]

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).^[63] 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.^[64] The biological half-life of rubidium in humans measures 31–46 days.^[56] 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.^[65][66]