Barium oxide

In this article, we will explore the fascinating life and legacy of Barium oxide, whose impact has transcended borders and generations. From his humble beginnings to his consecration as a leading figure in his field, Barium oxide has left an indelible mark on history. Throughout these pages, we will discover the most significant milestones in his career, his most notable achievements and the impact he has had on society. Through testimonies, analysis and reflections, we will immerse ourselves in the life and work of Barium oxide, recognizing his relevance and celebrating his invaluable contribution to humanity.

"BaO" redirects here. For other uses, see BAO (disambiguation) and Bao (disambiguation).
Barium oxide
Names
Other names
  • Neutral barium oxide (1:1)
  • Barium protoxide
  • Calcined baryta
  • Baria
Identifiers
3D model (JSmol)
ChemSpider
ECHA InfoCard 100.013.753 Edit this at Wikidata
EC Number
  • 215-127-9
RTECS number
  • CQ9800000
UNII
UN number 1884
  • InChI=1S/Ba.O checkY
    Key: QVQLCTNNEUAWMS-UHFFFAOYSA-N checkY
  • InChI=1/Ba.O/rBaO/c1-2
    Key: QVQLCTNNEUAWMS-FXUTYLCTAB
  • =O
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)
  • 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
Structure
cubic, cF8
Fm3m, No. 225
Octahedral
Thermochemistry
47.7 J/K mol
70 J·mol−1·K−1[1]
−582 kJ·mol−1[1]
Hazards
GHS labelling:
GHS05: CorrosiveGHS06: ToxicGHS07: Exclamation mark
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)
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
3
0
0
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 , 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]

Uses

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 ⇌ 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:

Barium oxide peroxidation from oxygen uptake, adapted from Middleburgh et al, 2012.[4]

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

Barium oxide from metallic barium readily 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

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

  • Barium – chemical element with symbol Ba and atomic number 56

References

  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. 2007. Retrieved 2007-02-19.
  3. ^ Nield, Gerald; Washecheck, Paul; Yang, Kang (1980-07-01). "United States Patent 4210764". Retrieved 2007-02-20.
  4. ^ 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.
  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. ^ "Compounds of barium: barium (II) oxide". Web Elements. The University of Sheffield. 2007-01-26. Retrieved 2007-02-22.
  7. ^ Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN 0-07-049439-8
  8. ^ "Barium Oxide (ICSC)". IPCS. October 1999. Archived from the original on 26 February 2007. Retrieved 2007-02-19.