Iron(II) sulfate
Skeletal formula of iron(II) sulfate
Iron(II) sulfate when dissolved in water
Structure of iron(II) sulfate heptahydrate
Sample of iron(II) sulfate heptahydrate
IUPAC name
Iron(II) sulfate
Other names
Iron(II) sulphate; Ferrous sulfate, Green vitriol, Iron vitriol, Ferrous vitriol, Copperas, Melanterite, Szomolnokite,
3D model (JSmol)
ECHA InfoCard 100.028.867 Edit this at Wikidata
EC Number
  • anhydrous: 231-753-5
RTECS number
  • anhydrous: NO8500000 (anhydrous)
    NO8510000 (heptahydrate)
UN number 3077
  • InChI=1S/Fe.H2O4S/c;1-5(2,3)4/h;(H2,1,2,3,4)/q+2;/p-2 checkY
  • anhydrous: InChI=1/Fe.H2O4S/c;1-5(2,3)4/h;(H2,1,2,3,4)/q+2;/p-2
  • anhydrous: [O-]S(=O)(=O)[O-].[Fe+2]
Molar mass 151.91 g/mol (anhydrous)
169.93 g/mol (monohydrate)
241.99 g/mol (pentahydrate)
260.00 g/mol (hexahydrate)
278.02 g/mol (heptahydrate)
Appearance White crystals (anhydrous)
White-yellow crystals (monohydrate)
Blue-green deliquescent[1] crystals (heptahydrate)
Odor Odorless
Density 3.65 g/cm3 (anhydrous)
3 g/cm3 (monohydrate)
2.15 g/cm3 (pentahydrate)[2]
1.934 g/cm3 (hexahydrate)[3]
1.895 g/cm3 (heptahydrate)[4]
Melting point 680 °C (1,256 °F; 953 K)
(anhydrous) decomposes[6]
300 °C (572 °F; 573 K)
(monohydrate) decomposes
60–64 °C (140–147 °F; 333–337 K)
(heptahydrate) decomposes[4][11]
44.69 g/100 mL (77 °C)
35.97 g/100 mL (90.1 °C)
15.65 g/100 mL (0 °C)
19.986 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)[5]
Solubility Negligible in alcohol
Solubility in ethylene glycol 6.38 g/100 g (20 °C)[6]
Vapor pressure 1.95 kPa (heptahydrate)[7]
1.24×10−2 cm3/mol (anhydrous)
1.05×10−2 cm3/mol (monohydrate)
1.12×10−2 cm3/mol (heptahydrate)[4]
+10200×10−6 cm3/mol
1.591 (monohydrate)[8]
1.526–1.528 (21 °C, tetrahydrate)[9]
1.513–1.515 (pentahydrate)[2]
1.468 (hexahydrate)[3]
1.471 (heptahydrate)[10]
Orthorhombic, oP24 (anhydrous)[12]
Monoclinic, mS36 (monohydrate)[8]
Monoclinic, mP72 (tetrahydrate)[9]
Triclinic, aP42 (pentahydrate)[2]
Monoclinic, mS192 (hexahydrate)[3]
Monoclinic, mP108 (heptahydrate)[4][10]
Pnma, No. 62 (anhydrous) [12]
C2/c, No. 15 (monohydrate, hexahydrate)[3][8]
P21/n, No. 14 (tetrahydrate)[9]
P1, No. 2 (pentahydrate)[2]
P21/c, No. 14 (heptahydrate)[10]
2/m 2/m 2/m (anhydrous)[12]
2/m (monohydrate, tetrahydrate, hexahydrate, heptahydrate)[3][8][9][10]
1 (pentahydrate)[2]
a = 8.704(2) Å, b = 6.801(3) Å, c = 4.786(8) Å (293 K, anhydrous)[12]
α = 90°, β = 90°, γ = 90°
Octahedral (Fe2+)
100.6 J/mol·K (anhydrous)[4]
394.5 J/mol·K (heptahydrate)[13]
107.5 J/mol·K (anhydrous)[4]
409.1 J/mol·K (heptahydrate)[13]
−928.4 kJ/mol (anhydrous)[4]
−3016 kJ/mol (heptahydrate)[13]
−820.8 kJ/mol (anhydrous)[4]
−2512 kJ/mol (heptahydrate)[13]
B03AA07 (WHO)
4 days [14]
2-4 months with peak activity at 7-10 days [15]
Legal status
GHS labelling:
GHS07: Exclamation mark[7]
H302, H315, H319[7]
NFPA 704 (fire diamond)
Lethal dose or concentration (LD, LC):
237 mg/kg (rat, oral)[11]
NIOSH (US health exposure limits):
REL (Recommended)
TWA 1 mg/m3[16]
Related compounds
Other cations
Cobalt(II) sulfate
Copper(II) sulfate
Manganese(II) sulfate
Nickel(II) sulfate
Related compounds
Iron(III) 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 ?)

Iron(II) sulfate (British English: iron(II) sulphate) or ferrous sulfate denotes a range of salts with the formula FeSO4·xH2O. These compounds exist most commonly as the heptahydrate (x = 7) but several values for x are known. The hydrated form is used medically to treat or prevent iron deficiency, and also for industrial applications. Known since ancient times as copperas and as green vitriol (vitriol is an archaic name for sulfate), the blue-green heptahydrate (hydrate with 7 molecules of water) is the most common form of this material. All the iron(II) sulfates dissolve in water to give the same aquo complex [Fe(H2O)6]2+, which has octahedral molecular geometry and is paramagnetic. The name copperas dates from times when the copper(II) sulfate was known as blue copperas, and perhaps in analogy, iron(II) and zinc sulfate were known respectively as green and white copperas.[18]

It is on the World Health Organization's List of Essential Medicines.[19] In 2021, it was the 105th most commonly prescribed medication in the United States, with more than 6 million prescriptions.[20][21]


Industrially, ferrous sulfate is mainly used as a precursor to other iron compounds. It is a reducing agent, and as such is useful for the reduction of chromate in cement to less toxic Cr(III) compounds. Historically ferrous sulfate was used in the textile industry for centuries as a dye fixative. It is used historically to blacken leather and as a constituent of iron gall ink.[22] The preparation of sulfuric acid ('oil of vitriol') by the distillation of green vitriol (iron(II) sulfate) has been known for at least 700 years.

Medical use

Main article: Iron supplement

Plant growth

Iron(II) sulfate is sold as ferrous sulfate, a soil amendment[23] for lowering the pH of a high alkaline soil so that plants can access the soil's nutrients.[24]

In horticulture it is used for treating iron chlorosis.[25] Although not as rapid-acting as ferric EDTA, its effects are longer-lasting. It can be mixed with compost and dug into the soil to create a store which can last for years.[26] Ferrous sulfate can be used as a lawn conditioner.[26] It can also be used to eliminate silvery thread moss in golf course putting greens.[27]

Pigment and craft

Ferrous sulfate can be used to stain concrete and some limestones and sandstones a yellowish rust color.[28]

Woodworkers use ferrous sulfate solutions to color maple wood a silvery hue.

Green vitriol is also a useful reagent in the identification of mushrooms.[29]

Historical uses

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 18th century. Chemical tests made on the Lachish letters (c. 588–586 BCE) showed the possible presence of iron.[30] It is thought that oak galls and copperas may have been used in making the ink on those letters.[31] 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 18th century and remained in use well into the 19th 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.

In the second half of the 1850s ferrous sulfate was used as a photographic developer for collodion process images.[32]


Iron(II) sulfate can be found in various states of hydration, and several of these forms exist in nature or were created synthetically.

Anhydrous iron(II) sulfate

The tetrahydrate is stabilized when the temperature of aqueous solutions reaches 56.6 °C (133.9 °F). At 64.8 °C (148.6 °F) these solutions form both the tetrahydrate and monohydrate.[5]

Mineral forms are found in oxidation zones of iron-bearing ore beds, e.g. pyrite, marcasite, chalcopyrite, etc. They are also found in related environments, like coal fire sites. Many rapidly dehydrate and sometimes oxidize. Numerous other, more complex (either basic, hydrated, and/or containing additional cations) Fe(II)-bearing sulfates exist in such environments, with copiapite being a common example.[41]

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.[42]

Fe + H2SO4 → FeSO4 + H2

Another source of large amounts results from the production of titanium dioxide from ilmenite via the sulfate process.

Ferrous sulfate is also prepared commercially by oxidation of pyrite:[43]

2 FeS2 + 7 O2 + 2 H2O → 2 FeSO4 + 2 H2SO4

It can be produced by displacement of metals less reactive than Iron from solutions of their sulfate:

CuSO4 + Fe → FeSO4 + Cu


Iron(II) sulfate outside a titanium dioxide factory in Kaanaa, Pori, Finland.

Upon dissolving in water, ferrous sulfates form the metal aquo complex [Fe(H2O)6]2+, which is an almost colorless, paramagnetic ion.

On heating, iron(II) sulfate first loses its water of crystallization and the original green crystals are converted into a white anhydrous solid. When further heated, the anhydrous material decomposes into sulfur dioxide and sulfur trioxide, leaving a reddish-brown iron(III) oxide. Thermolysis of iron(II) sulfate begins at about 680 °C (1,256 °F).

2 FeSO4 Fe2O3 + SO2 + SO3

Like other iron(II) salts, iron(II) sulfate is a reducing agent. For example, it reduces nitric acid to nitrogen monoxide 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

Its mild reducing power is of value in organic synthesis.[44] It is used as the iron catalyst component of Fenton's reagent.

Ferrous sulfate can be detected by the cerimetric method, which is the official method of the Indian Pharmacopoeia. This method includes the use of ferroin solution showing a red to light green colour change during titration.[45]

See also


  1. ^ Li R, Shi Y, Shi L, Alsaedi M, Wang P (1 May 2018). "Harvesting Water from Air: Using Anhydrous Salt with Sunlight". Environmental Science & Technology. 52 (9): 5398–5406. Bibcode:2018EnST...52.5398L. doi:10.1021/acs.est.7b06373. hdl:10754/627509. PMID 29608281.
  2. ^ a b c d e f "Siderotil Mineral Data". Retrieved 3 August 2014.
  3. ^ a b c d e f "Ferrohexahydrite Mineral Data". Retrieved 3 August 2014.
  4. ^ a b c d e f g h Lide DR, ed. (2009). CRC Handbook of Chemistry and Physics (90th ed.). Boca Raton, Florida: CRC Press. ISBN 978-1-4200-9084-0.
  5. ^ a b Seidell A, Linke WF (1919). Solubilities of Inorganic and Organic Compounds (2nd ed.). New York: D. Van Nostrand Company. p. 343.
  6. ^ a b Anatolievich KR. "iron(II) sulfate". Retrieved 3 August 2014.
  7. ^ a b c d Sigma-Aldrich Co., Iron(II) sulfate heptahydrate. Retrieved on 3 August 2014.
  8. ^ a b c d e Ralph J, Chautitle I. "Szomolnokite". Retrieved 3 August 2014.
  9. ^ a b c d e "Rozenite Mineral Data". Retrieved 3 August 2014.
  10. ^ a b c d e "Melanterite Mineral Data". Retrieved 3 August 2014.
  11. ^ a b "MSDS of Ferrous sulfate heptahydrate". Fair Lawn, New Jersey: Fisher Scientific, Inc. Retrieved 3 August 2014.
  12. ^ a b c d Weil M (2007). "The High-temperature β Modification of Iron(II) Sulfate". Acta Crystallographica Section E. 63 (12). International Union of Crystallography: i192. Bibcode:2007AcCrE..63I.192W. doi:10.1107/S160053680705475X. Retrieved 3 August 2014.
  13. ^ a b c d Anatolievich KR. "iron(II) sulfate heptahydrate". Retrieved 3 August 2014.
  14. ^ "Ferrous sulfate". Retrieved 11 December 2023.
  15. ^ "Ferrous sulfate". Retrieved 11 December 2023.
  16. ^ NIOSH Pocket Guide to Chemical Hazards. "#0346". National Institute for Occupational Safety and Health (NIOSH).
  17. ^ Safety Data Sheet
  18. ^ Brown, Lesley (1993). The New shorter Oxford English dictionary on historical principles. Oxford [Eng.]: Clarendon. ISBN 0-19-861271-0.
  19. ^ World Health Organization (2019). World Health Organization model list of essential medicines: 21st list 2019. Geneva: World Health Organization. hdl:10665/325771. WHO/MVP/EMP/IAU/2019.06. License: CC BY-NC-SA 3.0 IGO.
  20. ^ "The Top 300 of 2021". ClinCalc. Archived from the original on 15 January 2024. Retrieved 14 January 2024.
  21. ^ "Ferrous Sulfate - Drug Usage Statistics". ClinCalc. Retrieved 14 January 2024.
  22. ^ British Archaeology magazine. (archive)
  23. ^ "Why Use Ferrous Sulfate for Lawns?". Retrieved 14 April 2018.
  24. ^ "Acid or alkaline soil: Modifying pH - Sunset Magazine". 3 September 2004. Retrieved 14 April 2018.
  25. ^ Koenig, Rich and Kuhns, Mike: Control of Iron Chlorosis in Ornamental and Crop Plants. (Utah State University, Salt Lake City, August 1996) p.3
  26. ^ a b Handreck K (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.
  27. ^ Controlling moss in putting greens by Cook, Tom; McDonald, Brian; and Merrifield, Kathy.
  28. ^ How To Stain Concrete with Iron Sulfate
  29. ^ Svrček M (1975). A color guide to familiar mushrooms (2nd ed.). London: Octopus Books. p. 30. ISBN 0-7064-0448-3.
  30. ^ Torczyner, Lachish Letters, pp. 188–95
  31. ^ Hyatt, The Interpreter's Bible, 1951, volume V, p. 1067
  32. ^ Brothers A (1892). Photography: its history, processes. London: Griffin. p. 257. OCLC 558063884.
  33. ^ a b Meusburger J (September 2019). "Transformation mechanism of the pressure-induced C2/c-to-P transition in ferrous sulfate monohydrate single crystals". Journal of Solid State Chemistry. 277: 240–252. doi:10.1016/j.jssc.2019.06.004. S2CID 197070809.
  34. ^ "Rozenite".
  35. ^ Meusburger J (September 2022). "Low-temperature crystallography and vibrational properties of rozenite (FeSO4·4H2O), a candidate mineral component of the polyhydrated sulfate deposits on Mars" (PDF).
  36. ^ "Siderotil".
  37. ^ a b "Metal-sulfate Salts from Sulfide Mineral Oxidation". Retrieved 18 November 2022.
  38. ^ "Ferrohexahydrite".
  39. ^ "Melanterite".
  41. ^ "Copiapite".
  42. ^ Wildermuth E, Stark H, Friedrich G, Ebenhöch FL, Kühborth B, Silver J, et al. "Iron Compounds". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. ISBN 978-3527306732.
  43. ^ Lowson RT (1982). "Aqueous oxidation of pyrite by molecular oxygen". Chem. Rev. 82 (5): 461–497. doi:10.1021/cr00051a001.
  44. ^ Lee Irvin Smith, J. W. Opie (1948). "o-Aminobenzaldehyde". Org. Synth. 28: 11. doi:10.15227/orgsyn.028.0011.
  45. ^ Al-Obaidi AH. "ASSAY OF FERROUS SULPHATE" (PDF). Archived from the original (PDF) on 29 September 2023.
  46. ^ Pryce W (1778). Mineralogia Cornubiensis; a Treatise on Minerals, Mines and Mining. London: Phillips. p. 33.