Rubidium

Rubidium is a chemical element; it has symbol Rb

and atomic number 37. It is a very soft, whitish-grey
Rubidium, 37Rb
solid in the alkali metal group, similar to potassium
and caesium.[7] 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 85 Rb, and 28% is slightly
radioactive 87 Rb, with a half-life of 48.8 billion years
—more than three times as long as the estimated age
of the universe.

German chemists Robert Bunsen and Gustav Rubidium

Kirchhoff discovered rubidium in 1861 by the newly Pronunciation /ruːˈbɪdiəm/ ​
developed technique, flame spectroscopy. The name (roo-BID-ee-əm)
comes from the Latin word rubidus, meaning deep
red, the color of its emission spectrum. Rubidium's Appearance grey white
compounds have various chemical and electronic
Standard atomic weight Ar°(Rb)
applications. Rubidium metal is easily vaporized and
has a convenient spectral absorption range, making it 85.4678 ± 0.0003
a frequent target for laser manipulation of atoms. 85.468 ± 0.001 (abridged)[1]
Rubidium is not a known nutrient for any living
organisms. However, rubidium ions have similar Rubidium in the periodic table
properties and the same charge as potassium ions, K
and are actively taken up and treated by animal cells ↑
Rb
in similar ways. ↓
Cs
krypton ← rubidium → strontium
Characteristics
Atomic number (Z) 37
Group group 1: hydrogen and
alkali metals
Period period 5
Block s-block
Electron [Kr] 5s1
configuration
Electrons per 2, 8, 18, 8, 1
Partially molten rubidium metal in an shell
ampoule
Physical properties
Phase at STP solid
Rubidium is a very soft, ductile, silvery-white Melting point 312.45 K ​(39.30 °C, ​
metal.[8] It is the second most electropositive of the 102.74 °F)
stable alkali metals and melts at a temperature of
Boiling point 961 K ​(688 °C, ​1270 °F)
39.3 °C (102.7 °F). Like other alkali metals,
rubidium metal reacts violently with water. As with Density (near r.t.) 1.532 g/cm3
potassium (which is slightly less reactive) and when liquid 1.46 g/cm3
caesium (which is slightly more reactive), this (at m.p.)
reaction is usually vigorous enough to ignite the
hydrogen gas it produces. Rubidium has also been Triple point 312.41 K, ​? kPa[2]
reported to ignite spontaneously in air.[8] It forms Critical point 2093 K, 16 MPa
amalgams with mercury and alloys with gold, iron, (extrapolated)[2]
caesium, sodium, and potassium, but not lithium
Heat of fusion 2.19 kJ/mol
(even though rubidium and lithium are in the same
group).[9] Heat of 69 kJ/mol
vaporization
Molar heat 31.060 J/(mol·K)
capacity
Vapor pressure
P (Pa) 1 10 100 1k 10 k 100 k

at T (K) 434 486 552 641 769 958

Atomic properties
Rubidium crystals (silvery) compared
to caesium crystals (golden)
Oxidation states −1, +1 (a strongly basic
oxide)

Rubidium has a very low ionization energy of only Electronegativity Pauling scale: 0.82
406 kJ/mol.[10] Rubidium and potassium show a Ionization 1st: 403 kJ/mol
very similar purple color in the flame test, and energies 2nd: 2632.1 kJ/mol
distinguishing the two elements requires more
sophisticated analysis, such as spectroscopy. 3rd: 3859.4 kJ/mol

Atomic radius empirical: 248 pm

Compounds Covalent radius 220±9 pm
Van der Waals 303 pm
Rubidium chloride (RbCl) is radius
probably the most used
rubidium compound: among
several other chlorides, it is
used to induce living cells to Spectral lines of rubidium
take up DNA; it is also used as Other properties
a biomarker, because in nature,
Natural primordial
it is found only in small
Rb9O2 cluster quantities in living organisms occurrence
and when present, replaces Crystal structure ​body-centered cubic (bcc)
potassium. Other common
rubidium compounds are the corrosive rubidium
hydroxide (RbOH), the starting material for most
rubidium-based chemical processes; rubidium
Speed of sound 1300 m/s (at 20 °C)
carbonate (Rb2 CO3 ), used in some optical glasses,
thin rod
and rubidium copper sulfate, Rb2 SO4 ·CuSO4 ·6H2 O.
Rubidium silver iodide (RbAg4 I5 ) has the highest Thermal 90 µm/(m⋅K)[3] (at r.t.)
room temperature conductivity of any known ionic expansion
crystal, a property exploited in thin film batteries and
Thermal 58.2 W/(m⋅K)
other applications.[11][12]
conductivity
Rubidium forms a number of oxides when exposed Electrical 128 nΩ⋅m (at 20 °C)
to air, including rubidium monoxide (Rb2 O), Rb6 O, resistivity
and Rb9 O2 ; rubidium in excess oxygen gives the Magnetic ordering paramagnetic[4]
superoxide RbO2 . Rubidium forms salts with
Molar magnetic +17.0 × 10−6 cm3/mol
halogens, producing rubidium fluoride, rubidium
susceptibility (303 K)[5]
chloride, rubidium bromide, and rubidium iodide.[13]
Young's modulus 2.4 GPa
Bulk modulus 2.5 GPa
Isotopes
Mohs hardness 0.3
Although rubidium is monoisotopic, rubidium in the Brinell hardness 0.216 MPa
Earth's crust is composed of two isotopes: the stable CAS Number 7440-17-7
85 Rb (72.2%) and the radioactive 87 Rb (27.8%).[14]
Natural rubidium is radioactive, with specific activity History
of about 670 Bq/g, enough to significantly expose a Discovery Robert Bunsen and Gustav
photographic film in 110 days.[15][16] Thirty Kirchhoff (1861)
additional rubidium isotopes have been synthesized
First isolation George de Hevesy
with half-lives of less than 3 months; most are highly
radioactive and have few uses.[17] Isotopes of rubidium

Rubidium-87 has a half-life of 48.8 × 109 years, Main isotopes[6] Decay

which is more than three times the age of the abun­dance half-life (t1/2) mode pro­duct
universe of (13.799 ± 0.021) × 109 years,[18] making
it a primordial nuclide. It readily substitutes for
82
Rb synth 1.2575 m β+ 82
Kr
potassium in minerals, and is therefore fairly 83 83
Rb synth 86.2 d ε Kr
widespread. Rb has been used extensively in dating
rocks; 87 Rb beta decays to stable 87 Sr. During γ –
fractional crystallization, Sr tends to concentrate in
84 84
plagioclase, leaving Rb in the liquid phase. Hence, Rb synth 32.9 d ε Kr
the Rb/Sr ratio in residual magma may increase over
β+ 84
Kr
time, and the progressing differentiation results in
rocks with elevated Rb/Sr ratios. The highest ratios γ –
(10 or more) occur in pegmatites. If the initial amount
of Sr is known or can be extrapolated, then the age β− 84
Sr
can be determined by measurement of the Rb and Sr 85
Rb 72.2% stable
concentrations and of the 87 Sr/86 Sr ratio. The dates
indicate the true age of the minerals only if the rocks 86
Rb synth 18.7 d β− 86
Sr
have not been subsequently altered (see rubidium–
strontium dating).[19][20] γ –

87
Rb 27.8% 4.97 × 1010 y β− 87
Sr
Rubidium-82, one of the element's non-natural isotopes, is produced by electron-capture decay of
strontium-82 with a half-life of 25.36 days. With a half-life of 76 seconds, rubidium-82 decays by positron
emission to stable krypton-82.[14]

Occurrence

Rubidium is the twenty-third most abundant element in the Earth's crust, roughly as abundant as zinc and
rather more common than copper.[21] It occurs naturally in the minerals leucite, pollucite, carnallite, and
zinnwaldite, which contain as much as 1% rubidium oxide. Lepidolite contains between 0.3% and 3.5%
rubidium, and is the commercial source of the element.[22] Some potassium minerals and potassium
chlorides also contain the element in commercially significant quantities.[23]

Seawater contains an average of 125 µg/L of rubidium compared to the much higher value for potassium of
408 mg/L and the much lower value of 0.3 µg/L for caesium.[24] Rubidium is the 18th most abundant
element in seawater.[25]

Because of its large ionic radius, rubidium is one of the "incompatible elements".[26] During magma
crystallization, rubidium is concentrated together with its heavier analogue caesium in the liquid phase and
crystallizes last. Therefore, the largest deposits of rubidium and caesium are zone pegmatite ore bodies
formed by this enrichment process. Because rubidium substitutes for potassium in the crystallization of
magma, the enrichment is far less effective than that of caesium. Zone pegmatite ore bodies containing
mineable quantities of caesium as pollucite or the lithium minerals lepidolite are also a source for rubidium
as a by-product.[21]

Two notable sources of rubidium are the rich deposits of pollucite at Bernic Lake, Manitoba, Canada, and
the rubicline ((Rb,K)AlSi3 O8 ) found as impurities in pollucite on the Italian island of Elba, with a rubidium
content of 17.5%.[27] Both of those deposits are also sources of caesium.

Production
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.[21] Several methods are available
for separating potassium, rubidium, and caesium. The fractional crystallization
of a rubidium and caesium alum (Cs,Rb)Al(SO4 )2 ·12H2 O yields after 30
subsequent steps pure rubidium alum. Two other methods are reported, the
chlorostannate process and the ferrocyanide process.[21][28]

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.[29] Today the Flame test for rubidium
largest producers of caesium produce rubidium as a by-product from
pollucite.[21]

History
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".[30][31]
 Rubidium is a minor component in lepidolite. Kirchhoff and Bunsen
processed 150 kg of a lepidolite containing only 0.24% rubidium monoxide
(Rb2 O). 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
(Rb2 PtCl6 ) 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.[30][31]
Rubidium was the second element, shortly after caesium, to be discovered
Gustav Kirchhoff (left) and
by spectroscopy, just one year after the invention of the spectroscope by
Robert Bunsen (center) Bunsen and Kirchhoff.[32]
discovered rubidium by
spectroscopy. (Henry
The two scientists used the rubidium chloride to estimate that the atomic
Enfield Roscoe is on the
weight of the new element was 85.36 (the currently accepted value is
right.) 85.47).[30] 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
(Rb2 Cl); however, the product was probably a colloidal mixture of the metal and rubidium chloride.[33] 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.[34]

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 87 Rb to stable 87 Sr through beta decay was still under discussion in
the late 1940s.[35][36]

Rubidium had minimal industrial value before the 1920s.[37] 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,[38] for which the discoverers, Eric Allin
Cornell, Carl Edwin Wieman and Wolfgang Ketterle, won the 2001 Nobel Prize in Physics.[39]

Applications
Rubidium compounds are sometimes used in fireworks to give them a purple color.[40] 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.[41] These conduct electricity and act like an
armature of a generator, thereby generating an electric current. Rubidium, particularly vaporized 87 Rb, 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.[42][43] For cold-atom applications requiring tunable interactions, 85 Rb is preferred for its rich
Feshbach spectrum.[44]
 Rubidium has been used for polarizing 3 He, producing volumes of
magnetized 3 He gas, with the nuclear spins aligned rather than random.
Rubidium vapor is optically pumped by a laser, and the polarized Rb
polarizes 3 He through the hyperfine interaction.[45] Such spin-polarized
3 He cells are useful for neutron polarization measurements and for

producing polarized neutron beams for other purposes.[46]

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
A rubidium fountain atomic
transmitting, networking, and test equipment. These rubidium standards are
clock at the United States
often used with GPS to produce a "primary frequency standard" that has
Naval Observatory
greater accuracy and is less expensive than caesium standards.[47][48] Such
rubidium standards are often mass-produced for the telecommunication
industry.[49]

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.[50] 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.[51] In particular, 87 Rb is used with other alkali metals in the
development of spin-exchange relaxation-free (SERF) magnetometers.[51]

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.[52] 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.[53]

Rubidium was tested for the influence on manic depression and depression.[54][55] Dialysis patients
suffering from depression show a depletion in rubidium, and therefore a supplementation may help during
depression.[56] In some tests the rubidium was administered as rubidium chloride with up to 720 mg per
day for 60 days.[57][58]

Precautions and biological effects Rubidium

Hazards
Rubidium reacts violently with water and can cause fires. To
GHS labelling:
ensure safety and purity, this metal is usually kept under dry
mineral oil or sealed in glass ampoules in an inert atmosphere. Pictograms
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.[60] Signal word Danger
Hazard H260, H314
Rubidium, like sodium and potassium, almost always has +1 statements
oxidation state when dissolved in water, even in biological Precautionary P223, P231+P232,
contexts. The human body tends to treat Rb+ ions as if they were statements P280,
potassium ions, and therefore concentrates rubidium in the body's
P305+P351+P338,
intracellular fluid (i.e., inside cells).[61] The ions are not
P370+P378,
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 P422[59]
show negative effects in test persons.[62] The biological half-life of NFPA 704
rubidium in humans measures 31–46 days.[54] Although a partial (fire diamond) 4
substitution of potassium by rubidium is possible, when more than 3 2
50% of the potassium in the muscle tissue of rats was replaced W
with rubidium, the rats died.[63][64]

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Further reading
Meites, Louis (1963). Handbook of Analytical Chemistry (New York: McGraw-Hill Book
Company, 1963)
Steck, Daniel A. "Rubidium-87 D Line Data" (https://web.archive.org/web/20131102072437/
http://george.ph.utexas.edu/~dsteck/alkalidata/rubidium87numbers.pdf) (PDF). Los Alamos
National Laboratory (technical report LA-UR-03-8638). Archived from the original (http://geor
ge.ph.utexas.edu/~dsteck/alkalidata/rubidium87numbers.pdf) (PDF) on 2013-11-02.
Retrieved 2008-02-09.

External links
"Rubidium" (https://en.wikisource.org/wiki/1911_Encyclop%C3%A6dia_Britannica/Rubidiu
m). Encyclopædia Britannica. Vol. 23 (11th ed.). 1911. p. 809.
Rubidium (http://www.periodicvideos.com/videos/037.htm) at The Periodic Table of Videos
(University of Nottingham)

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