Iron(III) sulfate

Iron(III) sulfate (or ferric sulfate), is a family of inorganic compounds with the formula Fe2(SO4)3(H2O)n. A variety of hydrates are known, including the most commonly encountered form of "ferric sulfate". Solutions are used in dyeing as a mordant, and as a coagulant for industrial wastes. Solutions of ferric sulfate are also used in the processing of aluminum and steel.[2][3]

Iron(III) sulfate
Iron(III) sulfate
Names
IUPAC name
Iron(III) sulfate
Other names
Ferric sulfate
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.030.054
RTECS number
  • NO8505000
UNII
  • InChI=1S/2Fe.3H2O4S/c;;3*1-5(2,3)4/h;;3*(H2,1,2,3,4)/q2*+3;;;/p-6 checkY
    Key: RUTXIHLAWFEWGM-UHFFFAOYSA-H checkY
  • InChI=1/2Fe.3H2O4S/c;;3*1-5(2,3)4/h;;3*(H2,1,2,3,4)/q2*+3;;;/p-6
    Key: RUTXIHLAWFEWGM-CYFPFDDLAR
  • [Fe+3].[Fe+3].[O-]S(=O)(=O)[O-].[O-]S([O-])(=O)=O.[O-]S([O-])(=O)=O
Properties
Fe2(SO4)3
Molar mass 399.88 g/mol (anhydrous)
489.96 g/mol (pentahydrate)
562.00 g/mol (nonahydrate)
Appearance grayish-white crystals
Density 3.097 g/cm3 (anhydrous)
1.898 g/cm3 (pentahydrate)
Melting point 480 °C (896 °F; 753 K) (anhydrous)(decomposes)
175 °C (347 °F) (nonahydrate)
256g/L (monohydrate, 293K)
Solubility sparingly soluble in alcohol
negligible in acetone, ethyl acetate
insoluble in sulfuric acid, ammonia
1.814 (anhydrous)
1.552 (nonahydrate)
Hazards
NFPA 704 (fire diamond)
NFPA 704 four-colored diamond
1
0
0
Lethal dose or concentration (LD, LC):
500 mg/kg (oral, rat)
NIOSH (US health exposure limits):
REL (Recommended)
 TWA 1 mg/m3[1]
Related compounds
Other anions
 Iron(III) chloride
Iron(III) nitrate
Related compounds
 Iron(II) sulfate
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 ?)
Infobox references

Speciation

The various crystalline forms of Fe2(SO4)3(H2O)n are well-defined, often by X-ray crystallography. The nature of the aqueous solutions is often less certain, but aquo-hydroxo complexes such as [Fe(H2O)6]3+ and [Fe(H2O)5(OH)]2+ are often assumed.[4] Regardless, all such solids and solutions feature ferric ions, each with five unpaired electrons. By virtue of this high spin d5 electronic configuration, these ions are paramagnetic and are weak chromophores.

Production

Ferric sulfate solutions are usually generated from iron wastes. The actual identity of the iron species is often vague, but many applications do not demand high purity materials. It is produced on a large scale by treating sulfuric acid, a hot solution of ferrous sulfate, and an oxidizing agent. Typical oxidizing agents include chlorine, nitric acid, and hydrogen peroxide.[5]

2 FeSO4 + H2SO4 + H2O2 → Fe2(SO4)3 + 2 H2O

Natural occurrences

Iron sulfates occur as a variety of rare (commercially unimportant) minerals. Mikasaite, a mixed iron-aluminium sulfate of chemical formula (Fe3+, Al3+)2(SO4)3[6] is the name of mineralogical form of iron(III) sulfate. This anhydrous form occurs very rarely and is connected with coal fires. The hydrates are more common, with coquimbite[7] (nonahydrate) as probably the most often met among them. Paracoquimbite is the other, rarely encountered natural nonahydrate. Kornelite (heptahydrate) and quenstedtite (decahydrate) are rarely found. Andradite garnet is a yellow-green example found in Italy.[8] Lausenite (hexa- or pentahydrate) is a doubtful species. All the mentioned natural hydrates are unstable connected with the weathering (aerobic oxidation) of Fe-bearing primary minerals (mainly pyrite and marcasite).

Coquimbite crystal structure

See also

References
  1. NIOSH Pocket Guide to Chemical Hazards. "#0346". National Institute for Occupational Safety and Health (NIOSH).
  2. Ferric sulfate. The Columbia Encyclopedia, Sixth Edition. Retrieved November, 2007.
  3. Wildermuth, Egon; Stark, Hans; Friedrich, Gabriele; Ebenhöch, Franz Ludwig; Kühborth, Brigitte; Silver, Jack; Rituper, Rafael (2000). "Iron Compounds". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a14_591.
  4. Grant, M.; Jordan, R. B. (1981). "Kinetics of Solvent Water Exchange on Iron(III)". Inorganic Chemistry. 20: 55–60. doi:10.1021/ic50215a014.
  5. Iron compounds. Encyclopædia Britannica Article. Retrieved November, 2007
  6. Mikasaite
  7. "Minerals Colored by Metal Ions". minerals.gps.caltech.edu. Retrieved 2023-03-01.
  8. "Minerals Colored by Metal Ions". minerals.gps.caltech.edu. Retrieved 2023-03-01.
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