Iron(II) sulfate

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Iron(II) sulfate
Skeletal formula of iron(II) sulfate
Structure of iron(II) sulfate heptahydrate
Sample of iron(II) sulfate heptahydrate
IUPAC name
Iron(II) sulfate
Other names
Ferrous sulfate, Green vitriol, Iron vitriol, Copperas, Melanterite, Szomolnokite
7720-78-7 (anhydrous) YesY
17375-41-6 (monohydrate) N
7782-63-0 (heptahydrate) N
3D model (Jmol) Interactive image
ChEMBL ChEMBL1200830 N
ChemSpider 22804 (anhydrous) YesY
56459 (monohydrate) N
22804 (heptahydrate) N
ECHA InfoCard 100.028.867
EC Number 231-753-5
PubChem 24393 (anhydrous)
62712 (monohydrate)
62662 (heptahydrate)
RTECS number NO8500000 (anhydrous)
NO8510000 (heptahydrate)
UNII 2IDP3X9OUD (anhydrous) N
RIB00980VW (monohydrate) N
G0Z5449449 (dihydrate) N
39R4TAN1VT (heptahydrate) N
UN number 3077
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 crystals (heptahydrate)
Odor Odorless
Density 3.65 g/cm3 (anhydrous)
3 g/cm3 (monohydrate)
2.15 g/cm3 (pentahydrate)[1]
1.934 g/cm3 (hexahydrate)[2]
1.895 g/cm3 (heptahydrate)[3]
Melting point 680 °C (1,256 °F; 953 K)
(anhydrous) decomposes[5]
300 °C (572 °F; 573 K)
(monohydrate) decomposes
60–64 °C (140–147 °F; 333–337 K)
(heptahydrate) decomposes[3][10]
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)[4]
Solubility Negligible in alcohol
Solubility in ethylene glycol 6.4 g/100 g (20 °C)[5]
Vapor pressure 1.95 kPa (heptahydrate)[6]
1.24·10−2 cm3/mol (anhydrous)
1.05·10−2 cm3/mol (monohydrate)
1.12·10−2 cm3/mol (heptahydrate)[3]
+10,200·10−6 cm3/mol
1.591 (monohydrate)[7]
1.526–1.528 (21 °C, tetrahydrate)[8]
1.513–1.515 (pentahydrate)[1]
1.468 (hexahydrate)[2]
1.471 (heptahydrate)[9]
Orthorhombic, oP24 (anhydrous)[11]
Monoclinic, mS36 (monohydrate)[7]
Monoclinic, mP72 (tetrahydrate)[8]
Triclinic, aP42 (pentahydrate)[1]
Monoclinic, mS192 (hexahydrate)[2]
Monoclinic, mP108 (heptahydrate)[3][9]
Pnma, No. 62 (anhydrous) [11]
C2/c, No. 15 (monohydrate, hexahydrate)[2][7]
P21/n, No. 14 (tetrahydrate)[8]
P1, No. 2 (pentahydrate)[1]
P21/c, No. 14 (heptahydrate)[9]
2/m 2/m 2/m (anhydrous)[11]
2/m (monohydrate, tetrahydrate, hexahydrate, heptahydrate)[2][7][8][9]
1 (pentahydrate)[1]
a = 8.704(2) Å, b = 6.801(3) Å, c = 4.786(8) Å (293 K, anhydrous)[11]
α = 90°, β = 90°, γ = 90°
Octahedral (Fe2+)
100.6 J/mol·K (anhydrous)[3]
394.5 J/mol·K (heptahydrate)[12]
107.5 J/mol·K (anhydrous)[3]
409.1 J/mol·K (heptahydrate)[12]
−928.4 kJ/mol (anhydrous)[3]
−3016 kJ/mol (heptahydrate)[12]
−820.8 kJ/mol (anhydrous)[3]
−2512 kJ/mol (heptahydrate)[12]
B03AA07 (WHO)
GHS pictograms The exclamation-mark pictogram in the Globally Harmonized System of Classification and Labelling of Chemicals (GHS)[6]
GHS signal word Warning
H302, H315, H319[6]
Irritant Xi Harmful Xn
R-phrases R22, R36/38
S-phrases (S2), S46
NFPA 704
Flammability code 1: Must be pre-heated before ignition can occur. Flash point over 93 °C (200 °F). E.g., canola oil Health code 2: Intense or continued but not chronic exposure could cause temporary incapacitation or possible residual injury. E.g., chloroform Reactivity code 1: Normally stable, but can become unstable at elevated temperatures and pressures. E.g., calcium Special hazards (white): no codeNFPA 704 four-colored diamond
Lethal dose or concentration (LD, LC):
237 mg/kg (rat, oral)[10]
US health exposure limits (NIOSH):
REL (Recommended)
TWA 1 mg/m3[13]
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 YesYN ?)
Infobox references

Iron(II) sulfate (British English: iron(II) sulphate) or ferrous sulfate are salts with the formula FeSO4.xH2O. These compounds exist most commonly as the heptahydrate (x = 7) but are known for several values of x. The hydrated form 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.

It is on the World Health Organization's List of Essential Medicines, the most important medications needed in a basic health system.[14]


Industrially, ferrous sulfate is mainly used as a precursor to other iron compounds. It is a reducing agent, mostly for the reduction of chromate in cement. The term copperas refers to the historical use of ferrous sulfate as a chemical in the textile industry from the 17th century when it was used as a textile dye fixative, as a mean to blacken leather and also as a constituent of ink.[15]

Medical use[edit]

Main article: Iron supplement

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.[16] 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.[17]

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

Other uses[edit]

In horticulture it is used for treating iron chlorosis.[18] 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.[19] It is also used as a lawn conditioner,[19] and moss killer.

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

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.[citation needed]

It is used as a traditional method of treating wood panelling[clarification needed] on houses, either alone, dissolved in water, or as a component of water-based paint.[citation needed]

Green vitriol is also a useful reagent in the identification of mushrooms.[20]


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

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.[4]

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.

Iron(II) sulfate tetrahydrate, Rozenite (FeSO4·4H2O)

Production and reactions[edit]

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.[21]

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:

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


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 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
Ferrous sulfate outside titanium dioxide factory in Kaanaa, Pori, Finland.

Upon exposure to air, it oxidizes to form a corrosive brown-yellow coating of basic ferric sulfate, which is an adduct of ferric oxide and ferric sulfate:

12 FeSO4 + 3 O2 → 4 Fe2(SO4)3 + 2 Fe2O3

See also[edit]


  1. ^ a b c d e f "Siderotil Mineral Data". Retrieved 2014-08-03. 
  2. ^ a b c d e f "Ferrohexahydrite Mineral Data". Retrieved 2014-08-03. 
  3. ^ a b c d e f g h Lide, David R., ed. (2009). CRC Handbook of Chemistry and Physics (90th ed.). Boca Raton, Florida: CRC Press. ISBN 978-1-4200-9084-0. 
  4. ^ a b Seidell, Atherton; Linke, William F. (1919). Solubilities of Inorganic and Organic Compounds (2nd ed.). New York: D. Van Nostrand Company. p. 343. 
  5. ^ a b Anatolievich, Kiper Ruslan. "iron(II) sulfate". Retrieved 2014-08-03. 
  6. ^ a b c d Sigma-Aldrich Co., Iron(II) sulfate heptahydrate. Retrieved on 2014-08-03.
  7. ^ a b c d e Ralph, Jolyon; Chautitle, Ida. "Szomolnokite". Retrieved 2014-08-03. 
  8. ^ a b c d e "Rozenite Mineral Data". Retrieved 2014-08-03. 
  9. ^ a b c d e "Melanterite Mineral Data". Retrieved 2014-08-03. 
  10. ^ a b c "MSDS of Ferrous sulfate heptahydrate". Fair Lawn, New Jersey: Fisher Scientific, Inc. Retrieved 2014-08-03. 
  11. ^ a b c d 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. 
  12. ^ a b c d Anatolievich, Kiper Ruslan. "iron(II) sulfate heptahydrate". Retrieved 2014-08-03. 
  13. ^ "NIOSH Pocket Guide to Chemical Hazards #0346". National Institute for Occupational Safety and Health (NIOSH). 
  14. ^ "WHO Model List of Essential Medicines (19th List)" (PDF). World Health Organization. April 2015. Retrieved 8 December 2016. 
  15. ^ British Archeology magazine.
  16. ^ Hyatt, The Interpreter's Bible, 1951, volume V, p. 1067
  17. ^ How To Stain Concrete with Iron Sulfate
  18. ^ Koenig, Rich and Kuhns, Mike: Control of Iron Chlorosis in Ornamental and Crop Plants. (Utah State University, Salt Lake City, August 1996) p.3
  19. ^ a b 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. 
  20. ^ Svrček, Mirko (1975). A color guide to familiar mushrooms. (2nd ed.). London: Octopus Books. p. 30. ISBN 0-7064-0448-3. 
  21. ^ 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.
  22. ^ Pryce, William (1778). Mineralogia Cornubiensis; a Treatise on Minerals, Mines and Mining. - London, Phillips 1778. Phillips. p. 33. 

External links[edit]