| Bismuth(III) oxide | |
|---|---|
|
Bismuth trioxide |
|
|
Other names
Bismite, bismuth sesquioxide |
|
| Identifiers | |
| CAS number | 1304-76-3 |
| ChemSpider | 14093 |
| UNII | A6I4E79QF1 |
| Jmol-3D images | Image 1 |
|
|
|
|
| Properties | |
| Molecular formula | Bi2O3 |
| Molar mass | 465.96 g/mol |
| Appearance | yellow crystals or powder |
| Density | 8.90 g/cm3, solid |
| Melting point |
817 °C, 1090 K, 1503 °F[1] |
| Boiling point |
1890 °C, 2163 K, 3434 °F |
| Solubility in water | insoluble |
| Solubility | soluble in acids |
| Structure | |
| Crystal structure | monoclinic, mP20, Space group P21/c (No 14) |
| Coordination geometry |
pseudo-octahedral |
| Hazards | |
| MSDS | External MSDS |
| EU classification | not listed |
| NFPA 704 |
0
0
0
|
| Flash point | non-flammable |
| Related compounds | |
| Other anions | Bismuth trisulfide |
| Other cations | Arsenic trioxide Antimony trioxide |
| Supplementary data page | |
| Structure and properties |
n, εr, etc. |
| Thermodynamic data |
Phase behaviour Solid, liquid, gas |
| Spectral data | UV, IR, NMR, MS |
| (verify) (what is: /?) Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa) |
|
| Infobox references | |
Bismuth(III) oxide is perhaps 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. Bismuth trioxide is commonly used to produce the "Dragon's eggs" effect in fireworks, as a replacement of red lead.
Contents |
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 6s2 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 °C), 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.
Bismuth trioxide is commercially made from bismuth subnitrate. The latter is produced by dissolving bismuth in hot nitric acid. Addition of excess sodium hydroxide followed by continuous heating of the mixture precipitates bismuth(III) oxide as a heavy yellow powder. Also, the trioxide can be prepared by ignition of bismuth hydroxide.[1]
Oxidation with ammonium persulfate and dilute caustic soda gives bismuth tetroxide. The same product can be obtained by using other oxidizing agents such as potassium ferricyanide and concentrated caustic potash solution.
Electrolysis of bismuth(III) oxide in hot concentrated alkali solution gives a scarlet red precipitate of bismuth(V) oxide. Bismuth(III) oxide reacts with mineral acids to give the corresponding bismuth(III) salts.
Reaction with acetic anhydride and oleic acid gives bismuth trioleate.
|
|||||