Rubidium oxide

Rubidium oxide
Names
IUPAC name
Rubidium oxide
Other names
Rubidium(I) oxide
Dirubidium oxide
Identifiers
3D model (JSmol)
ChemSpider
ECHA InfoCard 100.038.161 Edit this at Wikidata
  • InChI=1S/O.2Rb/q-2;2*+1
    Key: YIONJVUULJNSMK-UHFFFAOYSA-N
  • [Rb+].[O-2].[Rb+]
Properties
Rb2O
Molar mass 186.94 g/moL
Appearance Yellow solid
Density 4 g/cm3
Melting point >500 °C
Reacts to give RbOH
+1527.0·10−6 cm3/mol
Structure
Antifluorite (cubic), cF12
Fm3m, No. 225
Tetrahedral (Rb+); cubic (O2−)
Hazards
Occupational safety and health (OHS/OSH):
Main hazards
 Corrosive, reacts violently with water
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasFlammability 0: Will not burn. E.g. waterInstability 1: Normally stable, but can become unstable at elevated temperatures and pressures. E.g. calciumSpecial hazard W: Reacts with water in an unusual or dangerous manner. E.g. sodium, sulfuric acid
3
0
1
W
Flash point Non-flammable
Related compounds
Other anions
 Rubidium sulfide
Rubidium selenide
Rubidium telluride
Rubidium polonide
Other cations
 Lithium oxide
Sodium oxide
Potassium oxide
Caesium oxide
Rubidium suboxide
Rubidium peroxide
Rubidium sesquioxide
Rubidium superoxide
Rubidium ozonide
Related compounds
 Rubidium hydroxide
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
checkY verify (what is checkYX markN ?)

Rubidium oxide is the chemical compound with the formula Rb2O. Rubidium oxide is highly reactive towards water, and therefore it would not be expected to occur naturally. The rubidium content in minerals is often calculated and quoted in terms of Rb2O. In reality, the rubidium is typically present as a component of (actually, an impurity in) silicate or aluminosilicate. A major source of rubidium is lepidolite, KLi2Al(Al,Si)3O10(F,OH)2, wherein Rb sometimes replaces K.

Rb2O is a yellow colored solid. The related species Na2O, K2O, and Cs2O are colorless, pale-yellow, and orange, respectively.

The alkali metal oxides M2O (M = Li, Na, K, Rb) crystallise in the antifluorite structure. In the antifluorite motif, the positions of the anions and cations are reversed relative to their positions in CaF2, with rubidium ions 4-coordinate (tetrahedral) and oxide ions 8-coordinate (cubic).

Properties

Like other alkali metal oxides, Rb2O is a strong base. Thus, Rb2O reacts exothermically with water to form rubidium hydroxide.

Rb2O + H2O → 2 RbOH

So reactive is Rb2O toward water that it is considered hygroscopic. Upon heating, Rb2O reacts with hydrogen to rubidium hydroxide and rubidium hydride:

Rb2O + H2 → RbOH + RbH

Synthesis

For laboratory use, RbOH, which is available commercially, is used in place of the oxide. The hydroxide is less reactive than the oxide.

As is the case for potassium oxide, Rb2O can be prepared by thermal reaction of the metal with a deficiency of oxygen:

4 Rb + O2 → 2 Rb2O

The Rb reaction with O2 is relatively colorful as it proceeds via bronze-colored Rb6O and copper-colored Rb9O2. The suboxides of rubidium that have been characterized by X-ray crystallography include Rb9O2 and Rb6O, as well as the mixed Cs-Rb suboxides Cs11O3Rbn (n = 1, 2, 3).

The product of oxygenation with excess oxygen is principally RbO2, rubidium superoxide:

Rb + O2 → RbO2

This superoxide can then be reduced to Rb2O using excess rubidium meta.

Alternatively rubidium metal reacts with mercuric oxide:

2 Rb + HgO → Rb2O + Hg

Typical for alkali metal hydroxides, RbOH cannot be dehydrated to the oxide.

References

  1. ↑ Wells, Alexander Frank (1984). Structural Inorganic Chemistry (5th ed.). Oxford: Clarendon Press. ISBN 978-0-19-855370-0.
  2. ↑ Nechamkin, Howard (1968). The chemistry of the elements. New York: McGraw-Hill. p. 34.
  3. ↑ Holleman, A.F.; Wiberg, E., eds. (2001). Inorganic Chemistry. San Diego: Academic Press. ISBN 978-0-12-352651-9.
  4. 1 2 E. Dönges (1963). "Alkali metals". In G. Brauer (ed.). Handbook of Preparative Inorganic Chemistry, 2nd Ed. Vol. 2. NY, NY: Academic Press. p. 977.
  5. ↑ Holleman, A.F.; Wiberg, E., eds. (2001). Inorganic Chemistry. San Diego: Academic Press. ISBN 978-0-12-352651-9.
  6. ↑ Simon, A. (1997). "Group 1 and 2 suboxides and subnitrides — Metals with atomic size holes and tunnels". Coordination Chemistry Reviews. 163: 253–270. doi:10.1016/S0010-8545(97)00013-1.

Further reading