Barium oxide
Other names
  • Neutral barium oxide (1:1)
  • Barium protoxide
  • Calcined baryta
  • Baria
3D model (JSmol)
ECHA InfoCard 100.013.753 Edit this at Wikidata
EC Number
  • 215-127-9
RTECS number
  • CQ9800000
UN number 1884
  • InChI=1S/Ba.O checkY
  • InChI=1/Ba.O/rBaO/c1-2
  • [Ba]=O
Molar mass 153.326 g/mol
Appearance white solid
Density 5.72 g/cm3, solid
Melting point 1,923 °C (3,493 °F; 2,196 K)
Boiling point ~ 2,000 °C (3,630 °F; 2,270 K)
  • 3.48 g/100 mL (20 °C)
  • 90.8 g/100 mL (100 °C)
  • Reacts to form Ba(OH)2
Solubility soluble in ethanol, dilute mineral acids and alkalies; insoluble in acetone and liquid ammonia
-29.1·10−6 cm3/mol
cubic, cF8
Fm3m, No. 225
47.7 J/K mol
70 J·mol−1·K−1[1]
−582 kJ·mol−1[1]
GHS labelling:
GHS05: CorrosiveGHS06: ToxicGHS07: Exclamation mark
H301, H302, H314, H315, H332, H412
P210, P220, P221, P260, P261, P264, P270, P271, P273, P280, P283, P301+P310, P301+P312, P301+P330+P331, P302+P352, P303+P361+P353, P304+P312, P304+P340, P305+P351+P338, P306+P360, P310, P312, P321, P330, P332+P313, P362, P363, P370+P378, P371+P380+P375, P405, P501
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 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no code
Flash point Non-flammable
Related compounds
Other anions
Other cations
Supplementary data page
Barium oxide (data page)
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
☒N verify (what is checkY☒N ?)

Barium oxide, also known as baria, is a white hygroscopic non-flammable compound with the formula BaO. It has a cubic structure and is used in cathode ray tubes, crown glass, and catalysts. It is harmful to human skin and if swallowed in large quantity causes irritation. Excessive quantities of barium oxide may lead to death.

It is prepared by heating barium carbonate with coke, carbon black or tar or by thermal decomposition of barium nitrate.[citation needed]


Barium oxide is used as a coating for hot cathodes, for example, those in cathode ray tubes. It replaced lead(II) oxide in the production of certain kinds of glass such as optical crown glass. While lead oxide raised the refractive index, it also raised the dispersive power, which barium oxide does not alter.[2] Barium oxide also has use as an ethoxylation catalyst in the reaction of ethylene oxide and alcohols, which takes place between 150 and 200 °C.[3]

It is also a source of pure oxygen through heat fluctuation. It readily oxidises to BaO2 by formation of a peroxide ion. The complete peroxidation of BaO to BaO2 occurs at moderate temperatures but the increased entropy of the O2 molecule at high temperatures means that BaO2 decomposes to O2 and BaO at 1175K.[4] The reaction was used as a large scale method to produce oxygen before air separation became the dominant method in the beginning of the 20th century. The method was named the Brin process, after its inventors.[5]


Barium oxide is made by heating barium carbonate at temperatures of 1000–1450 °C. It may also be prepared by thermal decomposition of barium nitrate.[6] Likewise, it is often formed through the decomposition of other barium salts.[7]

2 Ba + O2 → 2 BaO
BaCO3 → BaO + CO2

Safety issues

Barium oxide is an irritant. If it contacts the skin or the eyes or is inhaled it causes pain and redness. However, it is more dangerous when ingested. It can cause nausea and diarrhea, muscle paralysis, cardiac arrhythmia, and can cause death. If ingested, medical attention should be sought immediately.

Barium oxide should not be released environmentally; it is harmful to aquatic organisms.[8]

See also


  1. ^ a b Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. ISBN 978-0-618-94690-7.
  2. ^ "Barium Oxide (chemical compound)". Encyclopædia Britannica. Encyclopædia Britannica. 2007. Retrieved 2007-02-19.
  3. ^ Nield, Gerald; Washecheck, Paul; Yang, Kang (1980-07-01). "United States Patent 4210764". Retrieved 2007-02-20.
  4. ^ S.C. Middleburgh; K.P.D. Lagerlof; R.W. Grimes (2012-09-29). "Accommodation of Excess Oxygen in Group II Oxides". Journal of the American Ceramic Society. Retrieved 2022-03-27.
  5. ^ Jensen, William B. (2009). "The Origin of the Brin Process for the Manufacture of Oxygen". Journal of Chemical Education. 86 (11): 1266. Bibcode:2009JChEd..86.1266J. doi:10.1021/ed086p1266.
  6. ^ Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN 0-07-049439-8
  7. ^ "Compounds of barium: barium (II) oxide". Web Elements. The University of Sheffield. 2007-01-26. Retrieved 2007-02-22.
  8. ^ "Barium Oxide (ICSC)". IPCS. October 1999. Archived from the original on 26 February 2007. Retrieved 2007-02-19.