6SO4.png) | |
-sulfate-heptahydrate-sample.jpg) | |
| Names | |
|---|---|
| IUPAC name Iron(II) sulfate | |
| Other names Iron(II) sulphate; Ferrous sulfate, Green vitriol, Iron vitriol, Ferrous vitriol, Copperas, Melanterite, Szomolnokite, | |
| Identifiers | |
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3D model (JSmol) |
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| ChEBI |
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| ChEMBL |
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| ChemSpider | |
| ECHA InfoCard | 100.028.867 |
| EC Number |
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| RTECS number |
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| UNII |
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| UN number | 3077 |
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| Properties | |
| FeSO4 | |
| 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-greendeliquescent[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] |
| Monohydrate: 44.69 g/100 mL (77 °C) 35.97 g/100 mL (90.1 °C) Heptahydrate: 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 inalcohol |
| Solubility inethylene 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 | |
Refractive index (nD) | 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] |
| Structure | |
| 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+) | |
| Thermochemistry | |
| 100.6 J/mol·K (anhydrous)[4] 394.5 J/mol·K (heptahydrate)[13] | |
Std molar entropy(S⦵298) | 107.5 J/mol·K (anhydrous)[4] 409.1 J/mol·K (heptahydrate)[13] |
Std enthalpy of formation(ΔfH⦵298) | −928.4 kJ/mol (anhydrous)[4] −3016 kJ/mol (heptahydrate)[13] |
Gibbs free energy(ΔfG⦵) | −820.8 kJ/mol (anhydrous)[4] −2512 kJ/mol (heptahydrate)[13] |
| Pharmacology | |
| B03AA07 (WHO) | |
| none | |
| Pharmacokinetics: | |
| 4 days[14] | |
| 2-4 months with peak activity at 7-10 days[15] | |
| Legal status | |
| Hazards | |
| GHS labelling: | |
 [7] | |
| Warning | |
| H302,H315,H319[7] | |
| P305+P351+P338[7] | |
| NFPA 704 (fire diamond) | |
| Lethal dose or concentration (LD, LC): | |
LD50 (median dose) | 237 mg/kg (rat, oral)[11] |
| NIOSH (US health exposure limits): | |
REL (Recommended) | TWA 1 mg/m3[16] |
| Related compounds | |
Othercations | 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 theirstandard state (at 25 °C [77 °F], 100 kPa). | |
Iron(II) sulfate orferrous sulfate (British English:sulphate instead of sulfate) denotes a range ofsalts with the formulaFeSO4·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 preventiron deficiency, and also for industrial applications. Known since ancient times ascopperas and asgreen vitriol (vitriol is an archaic name forhydrated sulfate minerals), 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 sameaquo complex [Fe(H2O)6]2+, which hasoctahedral molecular geometry and isparamagnetic. 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 theWorld Health Organization's List of Essential Medicines.[19] In 2023, it was the 89th most commonly prescribed medication in the United States, with more than 7 million prescriptions.[20][21]
Industrially, ferrous sulfate is mainly used as a precursor to other iron compounds. It is areducing agent, and as such is useful for the reduction ofchromate incement to less toxic Cr(III) compounds. Historically, ferrous sulfate was used in the textile industry for centuries as adye fixative. It is used historically to blacken leather and as a constituent ofiron gall ink.[22] The preparation ofsulfuric acid ('oil of vitriol') by the distillation of green vitriol (iron(II) sulfate) has been known for at least 700 years.
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]
Inhorticulture, it is used for treating ironchlorosis.[25] Although not as rapid-acting asferric 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 alawn conditioner.[26] It can also be used to eliminatesilvery thread moss in golf course putting greens.[27]
Ferrous sulfate can be used to stain concrete and some limestones and sandstones a yellowish rust color.[28]
Woodworkers use ferrous sulfate solutions to colormaple wood a silvery hue.
Green vitriol is also a useful reagent in the identification of mushrooms.[29]
Ferrous sulfate was used in the manufacture ofinks, most notablyiron gall ink, which was used from theMiddle Ages until the end of the 18th century. Chemical tests made on theLachish 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 inwooldyeing as amordant.Harewood, a material used inmarquetry andparquetry since the 17th century, is also made using ferrous sulfate.
Two different methods for the direct application ofindigo dye were developed in England in the 18th century and remained in use well into the 19th century. One of these, known aschina blue, involved iron(II) sulfate. After printing an insoluble form of indigo onto the fabric, the indigo was reduced toleuco-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 forcollodion process images.[32]
Iron(II) sulfate can be found in various states ofhydration, and several of these forms exist in nature or were created synthetically.
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, withcopiapite being a common example.[41]
In the finishing ofsteel before plating or coating, the steel sheet or rod is passed throughpickling baths of sulfuric acid. This treatment produces large quantities of iron(II) sulfate as a by-product.[42]
Another source of large amounts results from the production oftitanium dioxide fromilmenite via the sulfate process.
Ferrous sulfate is also prepared commercially by oxidation ofpyrite:[43]
It can be produced by displacement of metals less reactive thaniron from solutions of their sulfate:
Upon dissolving in water, ferrous sulfates form themetal aquo complex [Fe(H2O)6]2+, which is an almost colorless,paramagnetic ion.
On heating, iron(II) sulfate first loses itswater of crystallization and the original green crystals are converted into a whiteanhydrous solid. When further heated, the anhydrous materialdecomposes intosulfur dioxide andsulfur trioxide, leaving a reddish-browniron(III) oxide.Thermolysis of iron(II) sulfate begins at about 680 °C (1,256 °F).
Like other iron(II) salts, iron(II) sulfate is a reducing agent. For example, it reducesnitric acid tonitrogen monoxide andchlorine tochloride:
Its mild reducing power is of value in organic synthesis.[44] It is used as the iron catalyst component ofFenton's reagent.
Ferrous sulfate can be detected by thecerimetric method, which is the official method of theIndian Pharmacopoeia. This method includes the use offerroin solution, showing a red to light green colour change during titration.[45]
