Bismuth(III) oxide

Chembox new
ImageFile = Bismuth(III) oxide.jpg
ImageName = Bismuth trioxide
IUPACName = Bismuth trioxide
Bismuth(III) oxide
Bismite (mineral)
OtherNames = Bismite
Section1 = Chembox Identifiers
CASNo = 1304-76-3
Section2 = Chembox Properties
Formula = Bi₂O₃
MolarMass = 465.96 g/mol

MolarMass_notes = "Bi=89.70%, O=10.30%"
Appearance = yellow crystals or powder
Density = 8.9 g/cm³, solid
Solvent = other solvents
SolubleOther = Insoluble
MeltingPt = 824°C
BoilingPt = 1890°C
Section3 = Chembox Structure
Coordination = pseudo-octahedral
CrystalStruct = monoclinic
Section7 = Chembox Hazards
ExternalMSDS =
EUClass = not listed
NFPA-H = 1
NFPA-F =
NFPA-R =
FlashPt = non-flammable
Section8 = Chembox Related
OtherAnions = Bismuth trisulfide
OtherCations = Arsenic trioxide
Antimony trioxide

Bismuth(III) oxide is the most industrially important compound of bismuth. It is also a common starting point for bismuth chemistry. It is found naturally as the mineral bismite (monoclinic) and sphaerobismoite (tetragonal, much more rare), but it is usually obtained as a by-product of the smelting of copper and lead ores. It may also be prepared by burning bismuth metal in air. Bismuth trioxide is commonly used to produce the "Dragon's eggs" effect in fireworks, as a replacement of red lead.

As a material for fuel cell electrolytes

Bismuth oxide has seen interest as a material for solid oxide fuel cells or SOFCs since it is an ionic conductor, i.e. oxygen atoms readily move through it. Pure bismuth oxide, Bi₂O₃ has four crystallographic polymorphs. It has a monoclinic crystal structure, designated α- Bi₂O₃, at room temperature. This transforms to the cubic fluorite-type crystal structure, δ-Bi₂O₃, when heated above 727°C, which remains the structure until the melting point, 824°C, is reached. The behaviour of Bi₂O₃ on cooling from the δ-phase is more complex, with the possible formation of two intermediate metastable phases; the tetragonal β-phase or the body-centred cubic γ-phase. The γ-phase can exist at room temperature with very slow cooling rates, but α- Bi₂O₃ always forms on cooling the β-phase.

δ- Bi₂O₃ has the highest reported conductivity. At 750°C the conductivity of δ- Bi₂O₃ is typically about 1 Scm⁻¹, about three orders of magnitude greater than the intermediate phases and four orders greater than the monoclinic phase. The conductivity in the β, γ and δ-phases is predominantly ionic with oxide ions being the main charge carrier. The α-phase exhibits p-type electronic conductivity (the charge is carried by positive holes) at room temperature which transforms to n-type conductivity (charge is carried by electrons) between 550°C and 650°C, depending on the oxygen partial pressure. It is therefore unsuitable for electrolyte applications. δ- Bi₂O₃ has a defective fluorite-type crystal structure in which two of the eight oxygen sites in the unit cell are vacant. These intrinsic vacancies are highly mobile due to the high polarisability of the cation sub-lattice with the 6"s"² lone pair electrons of Bi³⁺. The Bi-O bonds have covalent bond character and are therefore weaker than purely ionic bonds, so the oxygen ions can jump into vacancies more freely.

The arrangement of oxygen atoms within the unit cell of δ- Bi₂O₃ has been the subject of much debate in the past. Three different models have been proposed. Sillen (1937) used powder X-ray diffraction on quenched samples and reported the structure of Bi₂O₃ was a simple cubic phase with oxygen vacancies ordered along<111>, i.e. along the cube body diagonal (Figure 2a). Gattow and Schroder (1962) rejected this model, preferring to describe each oxygen site (8c site) in the unit cell as having 75% occupancy. In other words, the six oxygen atoms are randomly distributed over the eight possible oxygen sites in the unit cell. Currently, most experts seem to favour the latter description as a completely disordered oxygen sub-lattice accounts for the high conductivity in a better way.

Willis (1965) used neutron diffraction to study the fluorite (CaF₂) system. He determined that it could not be described by the ideal fluorite crystal structure, rather, the fluorine atoms were displaced from regular 8c positions towards the centres of the interstitial positions (Figure 2c). Shuk et al. (1996) and Sammes et al. (1999) suggest that because of the high degree of disorder in δ- Bi₂O₃, the Willis model could also be used to describe its structure.

In addition to electrical properties, thermal expansion properties are very important when considering possible applications for solid electrolytes. High thermal expansion coefficients represent large dimensional variations under heating and cooling which would limit the performance of an electrolyte. The transition from the high-temperature δ- Bi₂O₃ to the intermediate β- Bi₂O₃ is accompanied by a large volume change and consequently, a deterioration of the mechanical properties of the material. This, combined with the very narrow stability range of the δ-phase (727-824^oC), has led to studies on its stabilization to room temperature.

Bi₂O₃ easily forms solid solutions with many other metal oxides. These doped systems exhibit a complex array of structures and properties dependent on the type of dopant, the dopant concentration and the thermal history of the sample. The most widely studied systems are those involving rare earth metal oxides, Ln₂O₃, including yttria, Y₂O₃. Rare earth metal cations are generally very stable, have similar chemical properties to one another and are similar in size to Bi³⁺, which has a radius of 1.03 Å, making them all excellent dopants. Furthermore, their ionic radii decrease fairly uniformly from La³+ (1.032 Å), through Nd³⁺, (0.983 Å), Gd³⁺, (0.938 Å), Dy³⁺, (0.912 Å) and Er³⁺, (0.89 Å), to Lu³⁺, (0.861 Å) (known as the ‘lanthanide contraction’), making them useful to study the effect of dopant size on the stability of the Bi₂O₃ phases.

References

*Shannon, R. D., 1976 Acta Crystallographia A32:751
*Sammes, N. M., Tompsett, G. A., Cai, Z. H., 1999, Solid State Ionics 121:1-4
*Mairesse, G., Abraham, F., Nowogrocki, G., 1993, Journal of Solid State Chemistry 103:2
*Shuk, P., Wiemhofer, H.D., Guth, U., Gopel, W., Greenblatt, M., 1996 Solid State Ionics 89:3-4
*Willis, B. T. M., 1965 Acta Crystallographia 18:75
*Gattow, G., Schroder, H., 1962 Zeitschrift Für Anorganische Und Allgemeine Chemie 318:197
*Sillen, L. G. 1937 Ark. Kemi. Mineral. Geol. 12A:1
*Harwig, H. A., Gerards, A. G., 1978, Journal of Solid State Chemistry 26:265
*Harwig, H. A., 1978 Z. Anorg. Allg. Chem 444