Iron(II) carbonate
Names | |
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Other names
ferrous carbonate
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Identifiers | |
3D model (JSmol)
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ChemSpider | |
PubChem CID
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UNII | |
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Properties | |
FeCO3 | |
Molar mass | 115.854 g/mol |
Appearance | white powder or crystals |
Density | 3.9 g/cm3[1] |
Melting point | decomposes |
0.0067 g/L;[2] Ksp = 1.28 × 10−11 [3] | |
Solubility product (Ksp)
|
3.13×10−11[4] |
+11,300·10−6 cm3/mol | |
Structure | |
Hexagonal scalenohedral / Trigonal (32/m) Space group: R 3c, a = 4.6916 Å, c = 15.3796 Å | |
6 | |
Related compounds | |
Other anions
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iron(II) sulfate |
Other cations
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copper(II) carbonate, zinc carbonate |
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Iron(II) carbonate, or ferrous carbonate, is a chemical compound with formula FeCO
3, that occurs naturally as the mineral siderite. At ordinary ambient temperatures, it is a green-brown ionic solid consisting of iron(II) cations Fe2+
and carbonate anions CO2−
3.[5]
Preparation
Ferrous carbonate can be prepared by reacting solution of the two ions, such as iron(II) chloride and sodium carbonate:[5]
- FeCl
2 + Na
2CO
3 → FeCO
3 + 2NaCl
Ferrous carbonate can be prepared also from solutions of an iron(II) salt, such as iron(II) perchlorate, with sodium bicarbonate, releasing carbon dioxide:[6]
- Fe(ClO
4)2 + 2NaHCO
3 → FeCO
3 + 2NaClO
4 + CO
2 + H
2O
Sel and others used this reaction (but with FeCl
2 instead of Fe(ClO
4)2) at 0.2 M to prepare amorphous FeCO
3.[7]
Care must be taken to exclude oxygen O
2 from the solutions, because the Fe2+
ion is easily oxidized to Fe3+
, especially at pH above 6.0.[6]
Ferrous carbonate also forms directly on steel or iron surfaces exposed to solutions of carbon dioxide, forming an "iron carbonate" scale:[3]
- Fe + CO
2 + H
2O → FeCO
3 + H
2
Properties
The dependency of the solubility in water with temperature was determined by Wei Sun and others to be
- Failed to parse (SVG (MathML can be enabled via browser plugin): Invalid response ("Math extension cannot connect to Restbase.") from server "https://wikimedia.org/api/rest_v1/":): {\displaystyle \log K_{\mathit{sp}} = -59.3498 - 0.041377 T - 2.1963/T + 24.5724 \log T + 2.518 \sqrt{I} - 0.657 I, }
where T is the absolute temperature in kelvins, and I is the ionic strength of the liquid.[3]
Iron carbonate decomposes at about 500–600 °C (773–873 K).[8]
Uses
Ferrous carbonate has been used as an iron dietary supplement to treat anemia.[9] It is noted to have very poor bioavailability in cats and dogs.[10]
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).[11]
References
- ^ D R. Lide, ed.(2000): "CRC Handbook of Chemistry and Physics". 81st Edition. Pages 4-65.
- ^ Patty, F., ed. (1963): "Industrial Hygiene and Toxicology"; volume II: 'Toxicology". 2nd ed. Interscience. Page 1053.
- ^ a b c Wei Sun (2009): "Kinetics of iron carbonate and iron sulfide scale formation in CO2/H2S corrosion". PhD Thesis, Ohio University.
- ^ John Rumble (June 18, 2018). CRC Handbook of Chemistry and Physics (99 ed.). CRC Press. pp. 5–188. ISBN 1138561630.
- ^ a b (1995): "Kirk-Othmer Encyclopedia of Chemical Technology". 4th ed. Volume 1.
- ^ a b Philip C. Singer and Werner Stumm (1970): "The solubility of ferrous iron in carbonate-bearing waters". Journal of the American Water Works Association, volume 62, issue 3, pages 198-202. https://www.jstor.org/stable/41266171
- ^ 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. doi:10.1016/j.gca.2012.03.011
- ^ "Kinetics of Thermal Decomposition of Iron Carbonate". Egyptian Journal of Chemistry. 53 (6): 871–884. 2010-12-31. doi:10.21608/ejchem.2010.1268. ISSN 2357-0245.
- ^ A .Osol and J. E. Hoover and others, eds. (1975): "Remington's Pharmaceutical Sciences". 15th ed. Mack Publishing. Page 775
- ^ "AAFCO methods for substantiating nutritional adequacy of dog and cat foods (proposed for 2014 publication)" (PDF). AAFCO. 2013.
- ^ 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
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