Iron(III) sulfate
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| Names | |
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| IUPAC name
Iron(III) sulfate | |
| Other names
Ferric sulfate | |
| Identifiers | |
3D model (JSmol) |
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| ChEBI | |
| ChemSpider | |
| ECHA InfoCard | 100.030.054 |
PubChem CID |
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| RTECS number |
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| UNII | |
CompTox Dashboard (EPA) |
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| 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) |
| 256 g/L (monohydrate, 293 K) | |
| 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 | |
| GHS labelling:[1] | |
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| Danger | |
| H290, H302, H315, H317, H318, H319, H335 | |
| P234, P261, P264, P264+P265, P270, P271, P272, P280, P301+P317, P302+P352, P304+P340, P305+P351+P338, P305+P354+P338, P317, P319, P321, P330, P333+P317, P337+P317, P362+P364, P390, P403+P233, P405, P501 | |
| NFPA 704 (fire diamond) | |
| 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[2] |
| Related compounds | |
Other anions |
Iron(III) chloride Iron(III) nitrate |
Related compounds |
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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(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.[3][4]
Speciation
[edit]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.[5] 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
[edit]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.[6]
- 2 FeSO4 + H2SO4 + H2O2 → Fe2(SO4)3 + 2 H2O
Natural occurrences
[edit]Iron sulfates occur as a variety of rare commercially unimportant minerals. The mineralogical form of iron(III) sulfate, mikasaite, is a mixed iron-aluminium sulfate with the chemical formula (Fe3+, Al3+)2(SO4)3.[7] This anhydrous form occurs very rarely and is connected with coal fires. The hydrates are more common, with coquimbite[8] (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
[edit]- ↑ PubChem. "Ferric Sulfate". pubchem.ncbi.nlm.nih.gov. Retrieved 2026-06-08.
- ↑ 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 978-3-527-30673-2.
- ↑ 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 Archived 2012-03-22 at the Wayback Machine.
- 1 2 "Minerals Colored by Metal Ions". minerals.gps.caltech.edu. Retrieved 2023-03-01.
