Iron(II) oxalate

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Iron(II) oxalate
Names
IUPAC name
Iron(II) oxalate
Other names
Iron oxalate
Ferrous oxalate
Identifiers
3D model (JSmol)
ECHA InfoCard 100.007.472 Edit this at Wikidata
EC Number
  • 208-217-4
UNII
  • InChI=1S/3C2H2O4.2Fe/c3*3-1(4)2(5)6;;/h3*(H,3,4)(H,5,6);;/q;;;2*+3/p-6 ☒N
    Key: VEPSWGHMGZQCIN-UHFFFAOYSA-H ☒N
  • [Fe+2].O=C([O-])-C([O-])=O
Properties
FeC2O4 (anhydrous)
FeC2O4 · 2 H2O (dihydrate)
Molar mass 143.86 g/mol (anhydrous)
179.89 g/mol (dihydrate)
Appearance yellow powder
Odor odorless
Density 2.28 g/cm3
Melting point dihydrate: 150–160 °C (302–320 °F; 423–433 K)
(decomposes)
dihydrate:
0.097 g/100ml (25 °C)[1]
Hazards
GHS labelling:
GHS07: Exclamation mark[2]
Warning
H302, H312[2]
P280[2]
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
☒N verify (what is checkY☒N ?)

Ferrous oxalate (iron(II) oxalate) are inorganic compound with the formula FeC2O4(H2O)x where x is 0 or 2. These are orange compounds, poorly soluble in water.

Structure and reactions

Like other iron oxalates, ferrous oxalates feature octahedral Fe centers. The dihydrate FeC2O4(H2O)x is a coordination polymer, consisting of chains of oxalate-bridged ferrous centers, each with two aquo ligands.[3]
Ball-and-stick model of a chain in the crystal structure of iron(II) oxalate dihydrate

When heated to 120 °C, the dihydrate dehydrates, and the anhydrous ferrous oxalate decomposes near 190 °C.[4] The products of thermal decomposition is a mixture of iron oxides and pyrophoric iron metal, as well as released carbon dioxide, carbon monoxide, and water.[5]

Ferrous oxalates are precursors to iron phosphates, which are of value in batteries.[6]

Natural occurrence

Anhydrous iron(II) oxalate is unknown among minerals as of 2020. However, the dihydrate is known as humboldtine.[7][8] A related, though much more complex mineral is stepanovite,
Na[Mg(H2O)6] [Fe3+(C2O4)3]·3H2O - an example of trioxalatoferrate(III).[9][8]

See also

References

  1. ^ "Iron(II) oxalate dihydrate".
  2. ^ a b c Sigma-Aldrich Co., Iron(II) oxalate dihydrate. Retrieved on 2014-05-03.
  3. ^ Echigo, Takuya; Kimata, Mitsuyoshi (2008). "Single-crystal X-ray diffraction and spectroscopic studies on humboldtine and lindbergite: weak Jahn–Teller effect of Fe2+ ion". Physics and Chemistry of Minerals. 35 (8): 467–475. Bibcode:2008PCM....35..467E. doi:10.1007/s00269-008-0241-7. S2CID 98739882.
  4. ^ Mu, Jacob; Perlmutter, D.D. (1981). "Thermal decomposition of carbonates, carboxylates, oxalates, acetates, formates, and hydroxides". Thermochimica Acta. 49 (2–3): 207–218. doi:10.1016/0040-6031(81)80175-x.
  5. ^ Hermanek, Martin; Zboril, Radek; Mashlan, Miroslav; Machala, Libor; Schneeweiss, Oldrich (2006). "Thermal Behaviour of Iron(II) Oxalate Dihydrate in the Atmosphere of Its Conversion Gases". J. Mater. Chem. 16 (13): 1273–1280. doi:10.1039/b514565a.
  6. ^ Ellis, B. L.; Makahnouk, W. R. M.; Makimura, Y.; Toghill, K.; Nazar, L. F. (2007). "A multifunctional 3.5 V iron-based phosphate cathode for rechargeable batteries". Nature Materials. 6 (10): 749–753. Bibcode:2007NatMa...6..749E. doi:10.1038/nmat2007. PMID 17828278.
  7. ^ "Humboldtine".
  8. ^ a b "List of Minerals". 21 March 2011.
  9. ^ "Stepanovite".
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