Rubidium

Rubidium
Atomic number	37
Symbol	Rb
Group	1 (alkali metals)
Boiling point	688 °C
Electron configuration	[Kr] 5s1
Density	1.53 g/cm^3
Main isotopes	85Rb, 87Rb
Melting point	39.3 °C
Block	s
Phase STP	Solid
Oxidation states	+1
Wikidata	Q895

Rubidium (Rb, atomic number 37) is a soft, silvery-white alkali metal in Group 1 of the periodic table. Like other alkali metals (lithium, sodium, potassium, cesium), it has a single valence electron and is extremely electropositive. Rubidium metal is solid at ordinary temperatures (its melting point is only about 312 K, 39 °C) but tarnishes and oxidizes immediately in air. It ignites in air and reacts explosively with water to form rubidium hydroxide (RbOH) and hydrogen gas. Because of its high reactivity it must be stored under oil or inert gas. In chemical compounds Rb almost always has oxidation state +1, forming ionic salts with large Rb⁺ cations. Rubidium has some high-tech uses – for example Rb-87 is used in atomic clocks and Bose–Einstein condensate experiments – and its salts produce a characteristic reddish-violet flame color (used in spectrochemical tests). Discovered spectroscopically in 1861 by Robert Bunsen and Gustav Kirchhoff (named from Latin rubidus “deep red”), rubidium remains a relatively rare element with only niche applications, and it must be handled with care due to its vigorous reactivity and caustic products.

Overview and Key Facts

Rubidium (Rb) belongs to the alkali metals (Group 1, period 5, s-block). Its atomic mass is about 85.47 g/mol, reflecting the natural mix of isotopes. The element is a soft, silvery-white solid that is highly reactive: it ignites in air and even reacts explosively with water. Its only common oxidation state is +1 (forming Rb⁺); a rubidium anion (Rb⁻) is known only in special chemical contexts. At standard conditions rubidium metal crystallizes in a body-centered cubic lattice. It floats on water (density ≈1.53 g/cm³ at 20 °C) and has one of the lowest melting points of any metal (≈312 K or 39 °C). Its single valence electron (electron configuration [Kr] 5s¹) makes rubidium strongly reducing. Rubidium’s electronegativity (Pauling) is about 0.82 and its first ionization energy is very low (~403 kJ/mol), trends shared with the other alkali metals but more extreme. In fact, rubidium is more electropositive and reactive than potassium and sodium, and is surpassed only by cesium (and hypothetical francium) in reactivity. The element was named for the deep red lines in its emission spectrum and was first isolated in 1861 by Bunsen and Kirchhoff.

Atomic Structure and Electron Configuration

A rubidium atom has 37 electrons with shell structure 2–8–18–8–1. The electron configuration is [Kr] 5s¹, meaning there is one loosely held electron outside a krypton-like core. This lone 5s electron resides in the fifth principal shell (n=5) and is readily lost, which explains rubidium’s +1 chemistry. After that electron is removed, the Rb⁺ ion has a noble-gas (krypton) core and is relatively inert; this is reflected in the large jumps in ionization energy beyond the first. The first ionization energy is about 403 kJ/mol, whereas the second ionization energy jumps above 2600 kJ/mol (since Rb⁺ is already closed-shell). Rubidium’s atomic radius is large (on the order of a few hundred picometers) and its Rb⁺ ionic radius is about 152 pm (Pauling). These large sizes and low effective nuclear charge mean rubidium has an exceptionally low electronegativity (0.82) and ionization potential within the group. In general, alkali metals become larger and more reactive down the group, so rubidium (below potassium and above cesium) follows this trend: it is larger in radius and more reactive than K, but slightly less so than Cs.

Isotopes and Nuclear Properties

Naturally occurring rubidium is a mixture of two isotopes. The stable isotope ^85Rb makes up about 72% of natural rubidium, and ^87Rb about 28%. Rubidium-87 is technically radioactive: it undergoes beta-minus decay to stable ^87Sr with an extremely long half-life (about 4.9×10^10 years, tens of billions of years), so it persists in nature as if stable. Because of this, normal rubidium (even as metal) is slightly radioactive – enough that a piece can fog photographic film after weeks of proximity. Rubidium-87’s long half-life is also exploited in geochronology. In the rubidium–strontium dating method, the decay of ^87Rb to ^87Sr is used to date rocks and meteorites over geological timescales. The isotope ^87Rb has nuclear spin I = 3/2, whereas ^85Rb (abundance 72%) has spin I = 5/2. These spins lead to hyperfine structure in the atomic ground state. In fact, the 6.834 GHz ground-state hyperfine splitting of ^87Rb (F=1 and F=2 levels) is used as the basis for rubidium atomic clocks (see below).

Rubidium has no other long-lived isotopes in nature, and dozens of radioisotopes have been synthesized. Notably, ^82Rb (half-life 1.3 min) is used as a positron-emitting tracer in medical PET scans for myocardial perfusion. It is generated from a ^82Sr/^82Rb generator system and then injected so its gamma rays can image blood flow in the heart. Shorter-lived Rb radioisotopes (such as ^83Rb, ^84Rb, ^86Rb at tens of days half-life) are mainly of research interest. Other basic nuclear properties include the fact that ^85Rb has 37 protons and 48 neutrons, ^87Rb has 50 neutrons, and their natural abundances yield a weighted atomic mass of ~85.47. There are no nuclear spin 0 rubidium isotopes naturally; both have half-integer spin.

Allotropes and Typical Compounds

Rubidium has no molecular allotropes (unlike carbon or phosphorus); it forms one solid metallic phase (bcc) at ambient conditions. The metal is extremely soft (soft enough to be cut by a knife, similar to sodium) and can become liquid on a warm day. In chemical compounds, rubidium behaves like a typical alkali metal: it forms predominantly ionic species with Rb⁺ cations. Characteristic bonding involves the large Rb⁺ ion pairing with a variety of anions. For example, oxides of rubidium include the simple oxide Rb₂O, the peroxide Rb₂O₂, and especially the superoxide RbO₂ (yellow) when rubidium metal is exposed to excess oxygen. These oxygen compounds resemble those of potassium or cesium (which also form superoxides). Rubidium hydroxide (RbOH) is a strong, water-soluble base analogous to KOH. Other common Rb compounds are halides: rubidium fluoride (RbF), chloride (RbCl), bromide (RbBr) and iodide (RbI), all forming white ionic salts. Rubidium forms hydride (RbH) on reaction with hydrogen, and nitride (Rb₃N) or amide (RbNH₂) when melted with nitrogen or ammonia, though these are typically prepared only under controlled conditions. Rubidium also dissolves readily in liquid ammonia, producing the deep blue “alkali metal electride” solutions (solvated electrons) characteristic of alkali metals.

In chemistry, rubidium’s Rb⁺ is a large, weakly polarizing cation. It forms soluble salts with common anions: rubidium carbonate (Rb₂CO₃), nitrate (RbNO₃), sulfate (Rb₂SO₄), and perchlorate (RbClO₄) are all water-soluble ionic solids. In solid-state chemistry, rubidium can substitute for potassium in various minerals (lepidolite, feldspars, etc.) up to a few percent. It can also form alloys with other alkali metals (a continuous solid solution with cesium exists, yielding a eutectic melting point near 9 °C) and even forms amalgams with mercury and alloys with gold. Finally, rubidium’s flame test is distinctive: its salts color a flame deep reddish-violet (pink-purple), which is sometimes used analytically to detect the element.

Physical Properties

Rubidium is a metal and thus conducts electricity and heat well (its thermal conductivity is about 58 W/(m·K)). It has a shiny metallic luster that tarnishes rapidly in air. Its solid density is 1.532 g/cm³ at 20 °C – lower than water (so the metal floats) and higher than potassium (0.86 g/cm³) but lower than cesium (1.93 g/cm³). The metal’s crystal structure is body-centered cubic (bcc) under ambient conditions, like the other alkali metals. (No other solid forms or allotropes are known; at very low temperature it remains bcc until it melts near 312 K.) Rubidium’s bulk melting point is 312.4 K (39.3 °C) and its boiling point is 961 K (688 °C). Owing to its large atomic volume, the thermal expansion of rubidium is relatively high.

Optically, rubidium metal has characteristic atomic spectral lines. Rubidium atoms have “D” spectral lines analogous to sodium’s famous 589 nm resonance lines. The principal rubidium transitions are at about 780.0 nm (5s → 5p, the D₂ line) and 794.8 nm (the D₁ line) in the near-infrared (these are the strong red lines seen by Bunsen and Kirchhoff). These lines are routinely used in optical and laser physics – for example, diode lasers at 780–795 nm (originally common in CD/DVD technology) can excite rubidium atoms, which is why rubidium was chosen for early laser cooling and Bose–Einstein condensation.

Other measurable properties include a standard atomic weight 85.47, specific heat about 0.363 J/(g·K), and moderate melting/vaporization enthalpies (fusion ≈2.19 kJ/mol, vaporization ≈72 kJ/mol). In the infrared and visible, rubidium metal reflects light as a typical metal. No unusual magnetic ordering is seen; in bulk it is paramagnetic. Its superconducting or other exotic phases are not common knowledge. Rubidium’s electrical resistivity at 20 °C is around 0.194 μΩ·m, making it a reasonably good conductor (albeit poorer than typical transition metals).

Chemical Reactivity and Trends

Rubidium is highly chemically reactive, reflecting its position as an alkali metal. It reacts with oxygen immediately on exposure: the surface will form mixtures of oxides, peroxides (Rb₂O₂), and the superoxide RbO₂ (the latter often a yellowish solid). If sufficient oxygen is present or the metal is finely dispersed, Rb will ignite. With water, rubidium’s reaction is violently exothermic:

<code>2 Rb(s) + 2 H₂O(l) → 2 RbOH(aq) + H₂(g)↑  (reaction is very rapid and heat-generating)
</code>

The hydrogen gas often ignites or explodes from the heat of the reaction. The product is aqueous rubidium hydroxide, a very strong base (essentially fully dissociated Rb⁺ + OH⁻). With halogens, rubidium reacts readily to form halide salts. For example, Rb + ½Cl₂ → RbCl, releasing heat. It also will react spontaneously with bromine or iodine vapors to form RbBr, RbI, and even with fluorine to form RbF. With acids, rubidium metal also reacts vigorously: e.g. Rb + HCl → RbCl + ½ H₂.

Because it reacts so quickly, rubidium metal cannot exist in air or moisture for long. It even attacks glass and many container materials; pure Rb is kept under mineral oil or in sealed vessels. The metal forms alloys (“amalgams”) with mercury and mixes with other alkali metals. It will also react with nitrogen on heating to give Rb₃N and with hydrogen to give RbH (rubidium hydride), though these require controlled conditions. Rubidium metal does not passivate; it cannot be made nonreactive by coatings under normal conditions.

In liquid ammonia, like other alkalis, rubidium dissolves giving the deep blue solution of solvated electrons; this indicates that Rb is a very strong reducing agent. On the electrochemical series, Rb⁺/Rb has a standard reduction potential around –2.99 V, making rubidium among the strongest elemental reductants (stronger than Na or K, second only to Cs/Fr in theory).

As for chemical trends, rubidium behaves as expected for group 1. It is more reactive than potassium (lower ionization energy, larger size) but less reactive than cesium. Its chemistry otherwise parallels potassium’s: rubidium hydroxide is one of the strongest bases, Rb₂CO₃ is a stable carbonate, and RbNO₃ is a stable nitrate, etc. The Rb⁺ ion is too large and weakly polarizing to form covalent bonds of its own, so rubidium has virtually no chemistry beyond ionic compounds. Very rarely, rubidium can exhibit a formal –1 oxidation state in exotic compounds (for example, in certain alkali metal suboxides or alkaline-earth complexes), but in practice rubidium is always found as Rb⁺ or neutral metal.

Rubidium’s flame test is often noted: heating a rubidium salt in a flame yields a characteristic red-violet or pinkish-purple flame. This color comes from electronic transitions in the excited atoms and is used as a qualitative test for the element.

Occurrence and Production

Rubidium does not occur as a free element in nature. Its average abundance in Earth’s crust is roughly tens of parts per million (about 35–75 ppm by weight, similar to or a bit less than lithium or strontium). It is more abundant than zinc or copper, and far more plentiful than silver or gold.

Rubidium is typically found associated with potassium and lithium minerals. Important minerals containing rubidium (often substituting for K or Na) include lepidolite and zinnwaldite (lithium-bearing micas), pollucite (a cesium aluminosilicate), carnallite (a potassium–magnesium chloride), and leucite (a feldspathoid). Lepidolite, for example, may contain a few percent rubidium. Some brines and mineral springs also carry dissolved rubidium (for example, certain salt lake brines may contain a few mg/L of Rb). Rubidium is also present in sea water as Rb⁺ at very low concentration (~0.1 mg/L), but recovery from seawater is not economic.

Nearly all rubidium production is as a by-product of processing other minerals. The principal commercial source is potassium (potash) and lithium ores. For instance, in lithium extraction from lepidolite, the residual potassium-rich carbonates can contain ~20–25% Rb₂CO₃ along with K₂CO₃. During refining, rubidium is separated by chemical or ion-exchange methods. Because rubidium and cesium behave very similarly, specialized processes (such as crown-ether extraction or column chromatography) are used to isolate rubidium from cesium.

Once rubidium salts (usually RbCl or Rb₂CO₃) are obtained, the metal is produced chemically. A common method is reduction of rubidium chloride with metallic calcium or sodium:

<code>RbCl + Ca → 2 Rb + CaCl₂   (at high temperature)
</code>

Alternatively, direct electrolysis of molten rubidium salts (or fused cyanides) can yield the pure metal. Pure rubidium must be handled under oil or in inert gas to prevent immediate reaction. Because rubidium is produced only in small quantities (on the order of a few hundred to a few thousand kilograms per year worldwide), it is expensive and not a major industrial commodity. Major sources in the past have included Germany, Russia, and the United States, but production is limited and often integrated into existing potash/Li-ore processing. As a result, rubidium metal typically sells for hundreds of dollars per kilogram in recent years.

Applications and Technology

Rubidium’s high reactivity and unique isotopes give it a few specialized uses, despite its scarcity. One of the most important applications is in atomic clocks. Rubidium vapor cells are used in precision frequency standards: the hyperfine transition of ^87Rb at 6.834 GHz provides a stable microwave reference. Commercial rubidium oscillators are small and relatively low-cost compared to cesium clocks, so they are used in telecommunications, GPS ground stations, and network timekeeping where very good but not ultimate precision is required. (Rubidium clocks have lower accuracy than cesium standards, but are rugged and compact.)

In atomic and quantum physics research, rubidium is widely used as well. The isotope Rb-87 was one of the first atoms in which laser cooling and Bose–Einstein condensation (BEC) were achieved. In 1995, Cornell and Wieman (JILA) created a BEC in a gas of ultracold ^87Rb atoms. Rubidium is favored in such experiments because its atomic transitions match common laser wavelengths (780 nm) and it has suitable collisional properties. Many Bose–Einstein condensates and cold-atom experiments since have used rubidium or rubidium-potassium mixtures. Rb vapor cells are also used in atomic magnetometers (exploiting Rb ground-state coherence) and in other quantum-optical devices.

Rubidium metal and salts also have a few niche uses in electronics and glass. For example, rubidium is sometimes used in photoelectric devices: photocathodes made of cesium-antimony can be alloyed or dissolved with rubidium to tune work functions and response. Rubidium is used as a getter in vacuum tubes and vacuum devices – a small piece of rubidium will react with residual gases, helping to maintain high vacuum. In lighting and display technology, rubidium salts have been used in discharge lamps. Rubidium compounds have been added to certain optical glasses and ceramics to alter refractive index and dispersion (similar to how potassium or lanthanum oxides are used), although these uses are limited.

In pyrotechnics, rubidium salts are occasionally used to produce a purple flame color. A small amount of rubidium chloride or rubidium nitrate in a burner flares gives a deep red-violet flame. This is more a scientific demonstration than a large-scale application, since rubidium is expensive compared to other colorants.

From an industrial standpoint, rubidium has also been investigated as a propellant. Alkali metals like rubidium can serve in ion thrusters or magnetohydrodynamic power cycles because they easily ionize. Research has considered rubidium for electric propulsion in space or as a working fluid, but these have not become mainstream. One important niche application is medical imaging: as mentioned above, the radioisotope ^82Rb is used in positron-emission tomography (PET) to assess cardiac blood flow (a short-lived tracer that follows potassium pathways in the body).

Overall, except for these specialized roles, rubidium has very few large-scale commercial uses. Its high price and chemical similarity to cheaper alkali metals (potassium, sodium) means it is generally only used when its unique properties are essential.

Biology, Environment, and Safety

Biological role: Rubidium has no known essential biological function. However, chemically it is similar to potassium, and organisms do take up rubidium ions in place of potassium. The human body contains only trace amounts (~5 ppm by weight), mostly from dietary sources. Rubidium follows potassium pathways in cells; for example, it is transported by cells in the gut and placenta in much the same way as K⁺. In scientific research, radioactive ^86Rb has been used as a tracer for K⁺ transport in cells. At normal dietary levels, rubidium is not particularly toxic, but there is no metabolic requirement for it. Ingesting large amounts of rubidium salts can be harmful because the body cannot distinguish it fully from potassium; high rubidium can disrupt nerve or heart function similarly to high potassium (hyperkalemia-like effects). Chronic high intake of rubidium compounds has been associated with gastrointestinal upset, skin changes, and neurological symptoms in some animal studies. (Historically, rubidium chloride was once investigated as an antidepressant, but it has no current medical role.)

Environmental cycling: In the environment, rubidium behaves much like potassium and is not considered a pollutant. It is naturally present in soils and waters at low levels. Plants can absorb rubidium in trace amounts, often substituting for K⁺ without apparent harm. There are no specific ecological hazards associated with rubidium at typical concentrations. Its radioactivity from ^87Rb decay is so weak that it is of no environmental concern over human timescales. Rubidium is not bioaccumulative or persistent in the toxic-pollutant sense.

Safety: Chemically, rubidium is quite hazardous in elemental or compound form due to its reactivity. Rubidium metal is highly dangerous if it contacts water, air, or moisture – it can ignite or explode on contact. If rubidium metal burns, it forms Rb₂O₂ or RbO₂ and releases intense heat. Even small pieces of rubidium can cause fires. Contact of rubidium hydroxide (RbOH) or other soluble rubidium salts with skin or eyes causes severe chemical burns (they are strongly alkaline). Inhalation of rubidium compounds (like dust or aerosol) should be avoided; they can irritate respiratory tissues similarly to other alkalis.

Standard lab precautions apply: rubidium metal is handled under inert atmosphere (e.g. argon glovebox) or oil, and flammable metal should never be wetted. If rubidium metal contacts water inadvertently, use dry sand or Class D fire extinguishers (never water). Rubidium salts (e.g. RbCl) are of moderate toxicity if swallowed – on the order of “somewhat toxic” – and should be treated like any strong alkali (use gloves, eye protection). As a practical matter, rubidium is most often encountered in small quantities in laboratory settings or sealed devices (like vacuum tube getters), so accidental exposure is unlikely for the public. Regulatory exposure limits specifically for rubidium are generally not established, but safe handling mimics that of sodium and potassium compounds: avoid ingestion, inhalation, or skin contact with concentrated material.

History and Etymology

Rubidium was discovered in 1861 by Robert Bunsen and Gustav Kirchhoff in Heidelberg, Germany. Using Bunsen’s newly invented spectroscope, they analyzed mineral water samples and observed two bright red lines in the spectrum, unlike any known element. The name “rubidium” comes from the Latin rubidus, meaning “deep red,” referencing these spectral lines. (Bunsen and Kirchhoff simultaneously discovered cesium (blue lines) by the same method.) They initially isolated rubidium through chemical precipitation of its compounds. The pure metal was later obtained by electrolysis of rubidium salts and by reduction methods.

Unlike some alkalis known from antiquity (e.g. sodium), rubidium was unknown until modern spectroscopy. The element found its place in science mainly through spectroscopy and later advances. For much of its history rubidium had few applications, hindered by its reactivity and high cost. In the 20th century it became notable mainly in physics research. In physics lore, rubidium-87 was the first atom to achieve Bose–Einstein condensation (in 1995, work honored by the 2001 Nobel Prize to Cornell, Wieman, and Ketterle). Rubidium atomic clocks were developed around the 1950s–60s as a practical frequency standard (first commercial rubidium vaporers appeared in the 1960s).

In summary, rubidium’s story spans from 19th-century spectral discovery to modern quantum technology. Its deep-red spectral signature gave it a poetic name, and it now quietly serves in a few niche high-tech roles – from timekeeping to atomic research – all the while reminding chemists of the extreme reactivity of the alkali metals.

Data Table

Property	Value
Symbol	Rb
Atomic number (Z)	37
Standard atomic weight	85.4678
Phase at STP	Solid (metal)
Group / Period / Block	1 / 5 / s
Electron configuration	[Kr] 5s¹
Electronegativity (Pauling)	0.82
First ionization energy	403 kJ/mol
Atomic radius (covalent)	≈220 pm
Ionic radius (Rb⁺, CN=6)	152 pm
Density (20 °C)	1.53 g·cm⁻³
Melting point	312 K (39.2 °C)
Boiling point	961 K (688 °C)
Crystal structure (solid)	Body-centered cubic (bcc)
Common oxidation states	+1 (rarely −1)
Stable isotopes	^85Rb (72.2%), ^87Rb (27.8%)
CAS Registry Number	7440-17-7