Iron(II) oxide

Iron(II) oxide
Names
IUPAC name
Iron(II) oxide
Other names
Ferrous oxide,iron monoxide
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.014.292
13590
UNII
Properties
FeO
Molar mass 71.844 g/mol
Appearance black crystals
Density 5.745 g/cm3
Melting point 1,377 °C (2,511 °F; 1,650 K)[1]
Boiling point 3,414 °C (6,177 °F; 3,687 K)
Insoluble
Solubility insoluble in alkali, alcohol
dissolves in acid
+7200·10−6 cm3/mol
2.23
Hazards
Main hazards can be pyrophoric
Safety data sheet ICSC 0793
NFPA 704
Flammability code 3: Liquids and solids that can be ignited under almost all ambient temperature conditions. Flash point between 23 and 38 °C (73 and 100 °F). E.g., gasolineHealth code 0: Exposure under fire conditions would offer no hazard beyond that of ordinary combustible material. E.g., sodium chlorideReactivity code 2: Undergoes violent chemical change at elevated temperatures and pressures, reacts violently with water, or may form explosive mixtures with water. E.g., phosphorusSpecial hazards (white): no codeNFPA 704 four-colored diamond
3
0
2
variable
Related compounds
Other anions
 iron(II) fluoride, iron(II) sulfide, iron(II) selenide, iron(II) telluride
Other cations
 manganese(II) oxide, cobalt(II) oxide
Related compounds
 Iron(III) oxide, Iron(II,III) oxide
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Infobox references

Iron(II) oxide or ferrous oxide is the inorganic compound with the formula FeO. Its mineral form is known as wüstite. One of several iron oxides, it is a black-colored powder that is sometimes confused with rust, the latter of which consists of hydrated iron(III) oxide (ferric oxide). Iron(II) oxide also refers to a family of related non-stoichiometric compounds, which are typically iron deficient with compositions ranging from Fe0.84O to Fe0.95O.[2]

Preparation

FeO can be prepared by the thermal decomposition of iron(II) oxalate.

FeC2O4 → FeO + CO2 + CO

The procedure is conducted under an inert atmosphere to avoid the formation of ferric oxide. A similar procedure can also be used for the synthesis of manganous oxide and stannous oxide.[3][4]

Stoichiometric FeO can be prepared by heating Fe0.95O with metallic iron at 770 °C and 36 kbar.[5]

Reactions

FeO is thermodynamically unstable below 575 °C, tending to disproportionate to metal and Fe3O4:[2]

4FeO → Fe + Fe3O4

Structure

Iron(II) oxide adopts the cubic, rock salt structure, where iron atoms are octahedrally coordinated by oxygen atoms and the oxygen atoms octahedrally coordinated by iron atoms. The non-stoichiometry occurs because of the ease of oxidation of FeII to FeIII effectively replacing a small portion of FeII with two thirds their number of FeIII, which take up tetrahedral positions in the close packed oxide lattice.[5]

Below 200 K there is a minor change to the structure which changes the symmetry to rhombohedral and samples become antiferromagnetic.[5]

Occurrence in nature

Iron(II) oxide makes up approximately 9% of the Earth's mantle. Within the mantle, it may be electrically conductive, which is a possible explanation for perturbations in Earth's rotation not accounted for by accepted models of the mantle's properties.[6]

Iron III hydroxide staining caused by oxidation of dissolved iron II and precipitation, Perth, Western Australia.

Iron dissolved in groundwater is in the reduced iron II form. If this groundwater comes in contact with oxygen at the surface, e.g. in natural springs, iron II is oxidised to iron III and forms insoluble hydroxides in water.[7]

Uses

Iron(II) oxide is used as a pigment. It is FDA-approved for use in cosmetics and it is used in some tattoo inks. It can also be used as a phosphate remover from home aquaria.

References

  1. ↑ Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN 0-07-049439-8
  2. 1 2 Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0-08-037941-9.
  3. ↑ H. Lux "Iron (II) Oxide" in Handbook of Preparative Inorganic Chemistry, 2nd Ed. Edited by G. Brauer, Academic Press, 1963, NY. Vol. 1. p. 1497.
  4. ↑ Practical Chemistry for Advanced Students, Arthur Sutcliffe, 1930 (1949 Ed.), John Murray - London
  5. 1 2 3 Wells A.F. (1984) Structural Inorganic Chemistry 5th edition Oxford University Press ISBN 0-19-855370-6
  6. ↑ Science Jan 2012 Archived January 24, 2012, at the Wayback Machine.
  7. ↑ lenntech.com
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