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Iron(III) sulfate
Names
	IUPAC name Iron(III) sulfate
	Other names Ferric sulfate
Identifiers
CAS Number	10028-22-5 ;
3D model (JSmol)	Interactive image;
ChEBI	CHEBI:53438 ;
ChemSpider	23211 ;
ECHA InfoCard	100.030.054
PubChem CID	24826;
RTECS number	NO8505000;
UNII	4YKQ1X5E5Y ;
CompTox Dashboard (EPA)	DTXSID5029712 ;
	InChI InChI=1S/2Fe.3H2O4S/c;;31-5(2,3)4/h;;3(H2,1,2,3,4)/q2+3;;;/p-6 Key: RUTXIHLAWFEWGM-UHFFFAOYSA-H ; InChI=1/2Fe.3H2O4S/c;;31-5(2,3)4/h;;3(H2,1,2,3,4)/q2+3;;;/p-6 Key: RUTXIHLAWFEWGM-CYFPFDDLAR;
	SMILES [Fe+3].[Fe+3].[O-]S(=O)(=O)[O-].[O-]S([O-])(=O)=O.[O-]S([O-])(=O)=O;
Properties
Chemical formula	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)
Solubility in water	256g/L (monohydrate, 293K)
Solubility	sparingly soluble in alcohol ; negligible in acetone, ethyl acetate ; insoluble in sulfuric acid, ammonia
Refractive index (nD)	1.814 (anhydrous) ; 1.552 (nonahydrate)
Hazards
NFPA 704 (fire diamond)	1 0 0
	Lethal dose or concentration (LD, LC):
LD50 (median dose)	500 mg/kg (oral, rat)
	NIOSH (US health exposure limits):
REL (Recommended)	TWA 1 mg/m3
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). verify (what is ?) Infobox references

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]

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).

References

^ NIOSH Pocket Guide to Chemical Hazards. "#0346". National Institute for Occupational Safety and Health (NIOSH).
^ Ferric sulfate. The Columbia Encyclopedia, Sixth Edition. Retrieved November, 2007.
^ 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. ISBN 978-3527306732.
^ Grant, M.; Jordan, R. B. (1981). "Kinetics of Solvent Water Exchange on Iron(III)". Inorganic Chemistry. 20: 55–60. doi:10.1021/ic50215a014.
^ Iron compounds. Encyclopædia Britannica Article. Retrieved November, 2007
^ Mikasaite
^ "Minerals Colored by Metal Ions". minerals.gps.caltech.edu. Retrieved 2023-03-01.
^ "Minerals Colored by Metal Ions". minerals.gps.caltech.edu. Retrieved 2023-03-01.