|CAS number||(anhydrous) , (monohydrate) , (heptahydrate)|
|PubChem||(anhydrous) , (monohydrate) , (heptahydrate)|
|ChemSpider||(anhydrous) , (monohydrate) , (heptahydrate)|
|UNII||(anhydrous) , (monohydrate) , (dihydrate) , (heptahydrate)|
|RTECS number||NO8500000 (anhydrous)
|Jmol-3D images||Image 1|
|Molar mass||151.91 g mol−1|
|Appearance||White crystals (anhydrous)
White-yellow crystals (monohydrate)
Blue-green crystals (heptahydrate)
|Density||3.65 g/cm3 (anhydrous)
3 g/cm3 (monohydrate)
2.15 g/cm3 (pentahydrate)
1.934 g/cm3 (hexahydrate)
1.895 g/cm3 (heptahydrate)
|Melting point||680 °C (1,256 °F; 953 K)
300 °C (572 °F; 573 K)
60–64 °C (140–147 °F; 333–337 K)
|Solubility in water||Monohydrate:
44.69 g/100 mL (77 °C)
35.97 g/100 mL (90.1 °C)
15.65 g/100 mL (0 °C)
20.5 g/100 mL (10 °C)
29.51 g/100 mL (25 °C)
39.89 g/100 mL (40.1 °C)
51.35 g/100 mL (54 °C)
|Solubility||Negligible in alcohol|
|Solubility in ethylene glycol||6.4 g/100 g (20 °C)|
|Vapor pressure||1.95 kPa (heptahydrate)|
|Magnetic susceptibility||1.24·10−2 cm3/mol (anhydrous)
1.05·10−2 cm3/mol (monohydrate)
1.12·10−2 cm3/mol (heptahydrate)
|Refractive index (nD)||1.591 (monohydrate)
1.526–1.528 (21 °C, tetrahydrate)
|Crystal structure||Orthorhombic, oP24 (anhydrous)
Monoclinic, mS36 (monohydrate)
Monoclinic, mP72 (tetrahydrate)
Triclinic, aP42 (pentahydrate)
Monoclinic, mS192 (hexahydrate)
Monoclinic, mP108 (heptahydrate)
|Space group||Pnma, No. 62 (anhydrous) 
C2/c, No. 15 (monohydrate, hexahydrate)
P21/n, No. 14 (tetrahydrate)
P1, No. 2 (pentahydrate)
P21/c, No. 14 (heptahydrate)
|Point group||2/m 2/m 2/m (anhydrous)
2/m (monohydrate, tetrahydrate, hexahydrate, heptahydrate)
|Lattice constant||a = 8.704(2) Å, b = 6.801(3) Å, c = 4.786(8) Å (293 K, anhydrous)|
|Lattice constant||α = 90°, β = 90°, γ = 90°|
heat capacity C
|100.6 J/mol·K (anhydrous)
394.5 J/mol·K (heptahydrate)
|107.5 J/mol·K (anhydrous)
409.1 J/mol·K (heptahydrate)
|Std enthalpy of
|−928.4 kJ/mol (anhydrous)
−3016 kJ/mol (heptahydrate)
|Gibbs free energy ΔG||−820.8 kJ/mol (anhydrous)
−2512 kJ/mol (heptahydrate)
|GHS signal word||Warning|
|GHS hazard statements||H302, H315, H319|
|GHS precautionary statements||P305+351+338|
|EU Index||026-003-00-7 (anhydrous)
|EU classification||Xi Xn|
|LD50||237 mg/kg (rat, oral)|
|Other cations||Cobalt(II) sulfate
|Related compounds||Iron(III) sulfate|
|Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)|
|(what is: / ?)|
Iron(II) sulfate (Br.E. iron(II) sulphate) or ferrous sulfate is the chemical compound with the formula FeSO4. It is used medically to treat iron deficiency, and also for industrial applications. Known since ancient times as copperas and as green vitriol, the blue-green heptahydrate is the most common form of this material. All iron sulfates dissolve in water to give the same aquo complex [Fe(H2O)6]2+, which has octahedral molecular geometry and is paramagnetic.
Iron(II) sulfate can be found in various states of hydration, and several of these forms exist in nature.
- FeSO4·H2O (mineral: Szomolnokite, relatively rare)
- FeSO4·4H2O (mineral: Rozenite, white, relatively common, may be dehydratation product of melanterite)
- FeSO4·5H2O (mineral: Siderotil, relatively rare)
- FeSO4·6H2O (mineral: Ferrohexahydrite, relatively rare)
- FeSO4·7H2O (mineral: Melanterite, blue-green, relatively common)
The heptahydrate in solution (water as solvent) transforms to both heptahydrate and tetrahydrate when the temperature reaches 56.6 °C (133.9 °F). Then at 64.8 °C (148.6 °F) they form both tetrahydrate and monohydrate.
All mentioned mineral forms are connected with oxidation zones of Fe-bearing ore beds (pyrite, marcasite, chalcopyrite, etc.) and related environments (like coal fire sites). Many undergo rapid dehydration and sometimes oxidation.
Production and reactions
In the finishing of steel prior to plating or coating, the steel sheet or rod is passed through pickling baths of sulfuric acid. This treatment produces large quantities of iron(II) sulfate as a by-product.
- Fe + H2SO4 → FeSO4 + H2
Ferrous sulfate is also prepared commercially by oxidation of pyrite:
- 2 FeS2 + 7 O2 + 2 H2O → 2 FeSO4 + 2 H2SO4
On heating, iron(II) sulfate first loses its water of crystallization and the original green crystals are converted into a brown colored anhydrous solid. When further heated, the anhydrous material releases sulfur dioxide and white fumes of sulfur trioxide, leaving a reddish-brown iron(III) oxide. Decomposition of iron(II) sulfate begins at about 680 °C (1,256 °F).
- 2 FeSO4 → Fe2O3 + SO2 + SO3
Like all iron(II) salts, iron(II) sulfate is a reducing agent. For example, it reduces nitric acid to nitrogen oxide and chlorine to chloride:
- 6 FeSO4 + 3 H2SO4 + 2 HNO3 → 3 Fe2(SO4)3 + 4 H2O + 2 NO
- 6 FeSO4 + 3 Cl2 → 2 Fe2(SO4)3 + 2 FeCl3
- 12 FeSO4 + 3 O2 → 4 Fe2(SO4)3 + 2 Fe2O3
Together with other iron compounds, ferrous sulfate is used to fortify foods and to treat iron-deficiency anemia. Constipation is a frequent and uncomfortable side effect associated with the administration of oral iron supplements. Stool softeners often are prescribed to prevent constipation.
Ferrous sulfate was used in the manufacture of inks, most notably iron gall ink, which was used from the middle ages until the end of the eighteenth century. Chemical tests made on the Lachish letters [circa 588/6 BCE] showed the possible presence ... of iron (Torczyner, Lachish Letters, pp. 188–95). It is thought that oak galls and copperas may have been used in making the ink on those letters. It also finds use in wool dyeing as a mordant. Harewood, a material used in marquetry and parquetry since the 17th century, is also made using ferrous sulfate.
Two different methods for the direct application of indigo dye were developed in England in the eighteenth century and remained in use well into the nineteenth century. One of these, known as china blue, involved iron(II) sulfate. After printing an insoluble form of indigo onto the fabric, the indigo was reduced to leuco-indigo in a sequence of baths of ferrous sulfate (with reoxidation to indigo in air between immersions). The china blue process could make sharp designs, but it could not produce the dark hues of other methods. Sometimes, it is included in canned black olives as an artificial colorant.
Ferrous sulfate can also be used to stain concrete and some limestones and sandstones a yellowish rust color.
In horticulture it is used for treating iron chlorosis. Although not as rapid-acting as iron chelate, its effects are longer-lasting. It can be mixed with compost and dug into to the soil to create a store which can last for years. It is also used as a lawn conditioner, and moss killer.
Ferrous sulfate is sometimes added to the cooling water flowing through the brass tubes of turbine condensers to form a corrosion-resistant protective coating.
It is used in gold refining to precipitate metallic gold from auric chloride solutions (gold dissolved in solution with aqua regia).
It has been used in the purification of water by flocculation and for phosphate removal in municipal and industrial sewage treatment plants to prevent eutrophication of surface water bodies.
It is used as a traditional method of treating wood panelling on houses, either alone, dissolved in water, or as a component of water-based paint.
Green vitriol is also a useful reagent in the identification of mushrooms.
- Iron(III) sulfate, the other common simple sulfate of iron.
- Copper(II) sulfate,
- Mohr's salt (ammonium iron(II) sulfate), a common double salt of ammonium sulfate with iron(II) sulfate.
- "Siderotil Mineral Data". http://www.webmineral.com. Retrieved 2014-08-03.
- "Ferrohexahydrite Mineral Data". http://www.webmineral.com. Retrieved 2014-08-03.
- Lide, David R., ed. (2009). CRC Handbook of Chemistry and Physics (90th ed.). Boca Raton, Florida: CRC Press. ISBN 978-1-4200-9084-0.
- Seidell, Atherton; Linke, William F. (1919). Solubilities of Inorganic and Organic Compounds (2nd ed.). New York: D. Van Nostrand Company. p. 343.
- Anatolievich, Kiper Ruslan. "iron(II) sulfate". http://chemister.ru. Retrieved 2014-08-03.
- Sigma-Aldrich Co., Iron(II) sulfate heptahydrate. Retrieved on 2014-08-03.
- Ralph, Jolyon; Chautitle, Ida. "Szomolnokite". http://www.mindat.org. Mindat.org. Retrieved 2014-08-03.
- "Rozenite Mineral Data". http://www.webmineral.com. Retrieved 2014-08-03.
- "Melanterite Mineral Data". http://www.webmineral.com. Retrieved 2014-08-03.
- "MSDS of Ferrous sulfate heptahydrate". https://www.fishersci.ca. Fair Lawn, New Jersey: Fisher Scientific, Inc. Retrieved 2014-08-03.
- Weil, Matthias (2007). "The High-temperature β Modification of Iron(II) Sulfate". Acta Crystallographica Section E (International Union of Crystallography) 63 (12): i192. doi:10.1107/S160053680705475X. Retrieved 2014-08-03.
- Anatolievich, Kiper Ruslan. "iron(II) sulfate heptahydrate". http://chemister.ru. Retrieved 2014-08-03.
- Egon Wildermuth, Hans Stark, Gabriele Friedrich, Franz Ludwig Ebenhöch, Brigitte Kühborth, Jack Silver, Rafael Rituper “Iron Compounds” in Ullmann’s Encyclopedia of Industrial Chemistry Wiley-VCH, Wienheim, 2005.
- Hyatt, The Interpreter's Bible, 1951, volume V, p. 1067
- How To Stain Concrete with Iron Sulfate
- Koenig, Rich and Kuhns, Mike: Control of Iron Chlorosis in Ornamental and Crop Plants. (Utah State University, Salt Lake City, August 1996) p.3
- Handreck, Kevin (2002). Gardening Down Under: A Guide to Healthier Soils and Plants (2nd ed.). Collingwood, Victoria: CSIRO Publishing. pp. 146–47. ISBN 0-643-06677-2.
- Svrček, Mirko (1975). A color guide to familiar mushrooms. (2nd ed.). London: Octopus Books. p. 30. ISBN 0-7064-0448-3.
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|Salts and the ester of the Sulfate ion|