Barium oxide
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Other names
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3D model (JSmol)
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| ChemSpider | |
| ECHA InfoCard | 100.013.753 |
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PubChem CID
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| UNII | |
| UN number | 1884 |
CompTox Dashboard (EPA)
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| Properties | |
| BaO | |
| 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) |
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| Solubility | soluble in ethanol, dilute mineral acids and alkalies; insoluble in acetone and liquid ammonia |
| -29.1·10−6 cm3/mol | |
| Structure | |
| cubic, cF8 | |
| Fm3m, No. 225 | |
| Octahedral | |
| Thermochemistry | |
Heat capacity (C)
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47.7 J/K mol |
Std molar
entropy (S⦵298) |
70 J·mol−1·K−1[1] |
Std enthalpy of
formation (ΔfH⦵298) |
−582 kJ·mol−1[1] |
| Hazards | |
| GHS labelling: | |
  
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| Danger | |
| 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) | |
| Flash point | Non-flammable |
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Other anions
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Other cations
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| 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).
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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]
Uses
[edit]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 most known for its use in the Brin process, named after its inventors, a reaction that was used as a large scale method to produce oxygen before air separation became the dominant method in the beginning of the 20th century, as BaO can be a source of pure oxygen through heat fluctuation.
BaO(s) + ½O2(g) ⇌ BaO2(s)
It oxidises to BaO2 by formation of a peroxide ion ([O−O]2−, or O2−2) — with the same charge of O2−, and therefore keeping the electrochemical balance with the most stable Ba2+. Using the Kröger-Vink notation,
½O2(g) + O2–
O ⇌ [O
2]2–
O
where J
O is the species J in the oxygen position within the rock-salt lattice. The complete peroxidation of BaO to BaO2 occurs at moderate temperatures by oxygen uptake within the BaO rock-salt lattice:
Calculations using Density Functional Theory (DFT) suggest that the oxygen incorporation reaction is exothermic, and that the most energetically favoured occupation site is indeed the peroxide ion at the oxide lattice — other than interstitial positions, for instance. However, the increased entropy of the system is what leads BaO2 to decompose to BaO and release O2 between 800 and 1100 K (520 and 820 °C).[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]
Preparation
[edit]Barium oxide from metalic barium readly forms from its exothermic oxidation with dioxygen in air:
2 Ba(s) + O2(g) → 2 BaO(s).
It's most commonly made by heating barium carbonate at temperatures of 1000–1450 °C.
BaCO3(s) → BaO(s) + CO2(g)
Likewise, it is often formed through the thermal decomposition of other barium salts,[6] like barium nitrate.[7]
Safety issues
[edit]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
[edit]- Barium – chemical element with symbol Ba and atomic number 56
References
[edit]- ^ a b Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. ISBN 978-0-618-94690-7.
- ^ "Barium Oxide (chemical compound)". Encyclopædia Britannica. 2007. Retrieved 2007-02-19.
- ^ Nield, Gerald; Washecheck, Paul; Yang, Kang (1980-07-01). "United States Patent 4210764". Retrieved 2007-02-20.
- ^ a b Middleburgh, Simon C.; Lagerlof, Karl Peter D.; Grimes, Robin W. (2013). "Accommodation of Excess Oxygen in Group II Monoxides". Journal of the American Ceramic Society. 96 (1): 308–311. doi:10.1111/j.1551-2916.2012.05452.x. ISSN 1551-2916.
- ^ 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.
- ^ "Compounds of barium: barium (II) oxide". Web Elements. The University of Sheffield. 2007-01-26. Retrieved 2007-02-22.
- ^ Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN 0-07-049439-8
- ^ "Barium Oxide (ICSC)". IPCS. October 1999. Archived from the original on 26 February 2007. Retrieved 2007-02-19.
External links
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