Iron(II) carbonate, or ferrous carbonate, is a chemical compound with formula , that occurs naturally as the mineral siderite. At ordinary ambient temperatures, it is a green-brown ionic solid consisting of iron(II) cations and carbonate anions . The compound crystallizes in the same motif as calcium carbonate. In this motif, the carbonate dianion is nearly planar. Its three oxygen atoms each bind to two Fe(II) centers, such that the Fe has an octahedral coordination geometry.
Properties The dependency of the solubility in water with temperature was determined by Wei Sun and others to be
where T is the absolute temperature in kelvins, and I is the ionic strength of the liquid.
Iron carbonate decomposes at about .
Preparation Ferrous carbonate can be prepared by reacting solution of the two ions, such as iron(II) chloride and sodium carbonate:
+ → + 2
Ferrous carbonate can be prepared also from solutions of an iron(II) salt, such as iron(II) perchlorate, with sodium bicarbonate, releasing carbon dioxide:
()2 + 2 → + 2 + +
Sel and others used this reaction (but with instead of ()2) at 0.2 M to prepare amorphous .
Care must be taken to exclude oxygen from the solutions, because the ion is easily oxidized to , especially at pH above 6.0.
Ferrous carbonate also forms directly on steel or iron surfaces exposed to solutions of carbon dioxide, forming an "iron carbonate" scale:
+ + → +
Uses Ferrous carbonate has been used as an iron dietary supplement to treat anemia. It is noted to have very poor bioavailability in cats and dogs.
Toxicity Ferrous carbonate is slightly toxic; the probable oral lethal dose is between 0.5 and 5 g/kg (between 35 and 350 g for a 70 kg person).
Iron(III) carbonate Unlike iron(II) carbonate, iron(III) carbonate is rarely encountered. Attempts to produce iron(III) carbonate by the reaction of aqueous ferric ions and carbonate ions result in the production of iron(III) oxide with the release of carbon dioxide or bicarbonate. It is instead produced by the heating of iron(III) oxide at 2300 °C under 33 GPa of carbon dioxide. Iron(III) carbonate is unstable at atmospheric pressure, decomposing back into iron(III) oxide and carbon dioxide.
References (1995): "Kirk-Othmer Encyclopedia of Chemical Technology". 4th ed. Volume 1. Patty, F., ed. (1963): "Industrial Hygiene and Toxicology"; volume II: 'Toxicology". 2nd ed. Interscience. Page 1053. D R. Lide, ed.(2000): "CRC Handbook of Chemistry and Physics". 81st Edition. Pages 4-65. A .Osol and J. E. Hoover and others, eds. (1975): "Remington's Pharmaceutical Sciences". 15th ed. Mack Publishing. Page 775
Gosselin, R.E., H.C. Hodge, R.P. Smith, and M.N. Gleason. Clinical Toxicology of Commercial Products. 4th ed. Baltimore: Williams and Wilkins, 1976., p. II-97 Wei Sun (2009): "Kinetics of iron carbonate and iron sulfide scale formation in CO2/H2S corrosion". PhD Thesis, Ohio University. Ozlem Sel, A.V. Radha, Knud Dideriksen, and Alexandra Navrotsky (2012): "Amorphous iron (II) carbonate: Crystallization energetics and comparison to other carbonate minerals related to CO2 sequestration". Geochimica et Cosmochimica Acta, volume 87, issue 15, pages 61–68.