iron(II) sulfate

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.British Archaeology magazine. http://www.archaeologyuk.org/ba/ba66/feat2.shtml ([https://web.archive.org/web/20141017234401/http://www.archaeologyuk.org/ba/ba66/feat2.shtml archive]) 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.

{{Main|Iron supplement}}

Iron(II) sulfate is sold as ferrous sulfate, a soil amendment{{cite news|url=http://homeguides.sfgate.com/use-ferrous-sulfate-lawns-83484.html|title=Why Use Ferrous Sulfate for Lawns?|access-date=14 April 2018 }} for lowering the pH of a high alkaline soil so that plants can access the soil's nutrients.{{cite web|url=https://www.sunset.com/garden/garden-basics/acid-alkaline-soil-modifying-ph|title=Acid or alkaline soil: Modifying pH - Sunset Magazine|website=www.sunset.com|date=3 September 2004|access-date=14 April 2018}}

In horticulture it is used for treating iron chlorosis.Koenig, Rich and Kuhns, Mike: Control of Iron Chlorosis in Ornamental and Crop Plants. (Utah State University, Salt Lake City, August 1996) [http://extension.usu.edu/files/publications/publication/AG-SO-01.pdf p.3] 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.{{cite book|last=Handreck|first=Kevin|title=Gardening Down Under: A Guide to Healthier Soils and Plants|publisher=CSIRO Publishing|location=Collingwood, Victoria|year=2002|edition=2nd|pages=146–47|isbn=0-643-06677-2}} Ferrous sulfate can be used as a lawn conditioner. It can also be used to eliminate silvery thread moss in golf course putting greens.[https://www.gcsaa.org/docs/default-source/research-and-information/weeds/controlling-moss-in-putting-greens.pdf#page=2/ Controlling moss in putting greens by Cook, Tom; McDonald, Brian; and Merrifield, Kathy.]

Ferrous sulfate can be used to stain concrete and some limestones and sandstones a yellowish rust color.[http://www.stainedfloor.com/Iron_Recipes.html How To Stain Concrete with Iron Sulfate]

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

Green vitriol is also a useful reagent in the identification of mushrooms.{{cite book|last=Svrček|first=Mirko|title=A color guide to familiar mushrooms.|publisher=Octopus Books|location=London|year=1975|edition=2nd|page=[https://archive.org/details/colourguidetofam00svrc/page/30 30]|isbn=0-7064-0448-3|url=https://archive.org/details/colourguidetofam00svrc/page/30}}

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 ({{circa|588–586 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.Hyatt, The Interpreter's Bible, 1951, volume V, p. 1067 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.{{cite book |last1=Brothers |first1=Alfred |title=Photography: its history, processes |url=https://archive.org/details/b24886853 |date=1892 |publisher=Griffin |location=London |oclc=558063884 |page=[https://archive.org/details/b24886853/page/257 257]}}

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

• FeSO₄·H₂O (mineral: szomolnokite,{{cite web|first1= Jolyon|last1= Ralph|first2= Ida|last2= Chautitle|title= Szomolnokite|url= http://www.mindat.org/min-3859.html|publisher= Mindat.org|access-date= 3 August 2014}} relatively rare, monoclinic{{Cite journal |last=Meusburger |first=Johannes |date=September 2019 |title=Transformation mechanism of the pressure-induced C2/c-to-P transition in ferrous sulfate monohydrate single crystals |journal=Journal of Solid State Chemistry |volume=277 |pages=240–252 |doi=10.1016/j.jssc.2019.06.004 |s2cid=197070809 |url=https://www.sciencedirect.com/science/article/abs/pii/S0022459619302853|url-access=subscription }})
• FeSO₄·H₂O (synthetic compound stable at pressures exceeding 6.2 GPa, triclinic)
• FeSO₄·4H₂O (mineral: rozenite,{{cite web|title= Rozenite Mineral Data|url= http://www.webmineral.com/data/Rozenite.shtml|access-date= 3 August 2014}}{{Cite web|url=https://www.mindat.org/min-3469.html|title=Rozenite}} white, relatively common, may be dehydration product of melanterite, monoclinic{{Cite web |last=Meusburger |first=Johannes |date=September 2022 |title=Low-temperature crystallography and vibrational properties of rozenite (FeSO4·4H2O), a candidate mineral component of the polyhydrated sulfate deposits on Mars |url=http://www.minsocam.org/msa/ammin/AM_Preprints/8502ForbesPreprint.pdf}})
• FeSO₄·5H₂O (mineral: siderotil,{{cite web|title= Siderotil Mineral Data|url= http://www.webmineral.com/data/Siderotil.shtml|access-date= 3 August 2014}}{{Cite web|url=https://www.mindat.org/min-3643.html|title=Siderotil}} relatively rare, triclinic{{Cite journal |url=https://pubs.geoscienceworld.org/msa/rimg/article-abstract/40/1/303/140666/Metal-sulfate-Salts-from-Sulfide-Mineral-Oxidation?redirectedFrom=fulltext |access-date=18 November 2022|title= Metal-sulfate Salts from Sulfide Mineral Oxidation|journal=Reviews in Mineralogy and Geochemistry|date=2000 |doi=10.2138/rmg.2000.40.6 |volume=40 |issue=1 |pages=303–350 |bibcode=2000RvMG...40..303J | vauthors = Jambor JL, Nordstrom DK, Alpers CN |url-access=subscription }})
• FeSO₄·6H₂O (mineral: ferrohexahydrite,{{cite web|title= Ferrohexahydrite Mineral Data|url= http://www.webmineral.com/data/Ferrohexahydrite.shtml|access-date= 3 August 2014}}{{Cite web|url=https://www.mindat.org/min-1517.html|title=Ferrohexahydrite}} very rare, monoclinic)
• FeSO₄·7H₂O (mineral: melanterite,{{cite web|title= Melanterite Mineral Data|url= http://www.webmineral.com/data/Melanterite.shtml|access-date= 3 August 2014}}{{Cite web|url=https://www.mindat.org/min-2633.html|title=Melanterite}} blue-green, relatively common, monoclinic{{Cite web |last=Peterson |first=RC |date=2003 |title=THE RELATIONSHIP BETWEEN Cu CONTENT AND DISTORTION IN THE ATOMIC STRUCTURE OF MELANTERITE FROM THE RICHMOND MINE, IRON MOUNTAIN, CALIFORNIA |url=https://rruff.info/doclib/cm/vol41/CM41_937.pdf}})

File:Síran železnatý.PNG

The tetrahydrate is stabilized when the temperature of aqueous solutions reaches {{convert|56.6|C|F}}. At {{convert|64.8|C|F}} these solutions form both the tetrahydrate and monohydrate.{{cite book|last1= Seidell|first1= Atherton|last2= Linke|first2= William F.|year= 1919|title= Solubilities of Inorganic and Organic Compounds|url= https://archive.org/details/solubilitiesino01seidgoog|publisher= D. Van Nostrand Company|place= New York|edition= 2nd|page= [https://archive.org/details/solubilitiesino01seidgoog/page/n368 343]}}

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.{{Cite web|url=https://www.mindat.org/min-1124.html|title=Copiapite}}

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.{{ullmann|first1=Egon|last1=Wildermuth|first2=Hans|last2=Stark|first3=Gabriele|last3=Friedrich|first4=Franz Ludwig|last4=Ebenhöch|first5=Brigitte|last5=Kühborth|first6=Jack|last6=Silver|first7=Rafael|last7=Rituper|title=Iron Compounds}}

:{{chem2 | 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:{{cite journal |title= Aqueous oxidation of pyrite by molecular oxygen |first= Richard T. |last= Lowson |journal= Chem. Rev. |year= 1982 |volume= 82 |issue= 5 |pages= 461–497 |doi= 10.1021/cr00051a001 }}

:{{chem2 | 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:

:{{chem2 | CuSO4 + Fe -> FeSO4 + Cu }}

File:Ferric sulphate, Kemira.jpg 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 {{convert|680|C|F}}.

:{{chem2 | 2 FeSO4 }} {{overset|Δ|→}} {{chem2 | 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:

:{{chem2 | 6 FeSO4 + 3 H2SO4 + 2 HNO3 -> 3 Fe2(SO4)3 + 4 H2O + 2 NO }}

:{{chem2 | 6 FeSO4 + 3 Cl2 -> 2 Fe2(SO4)3 + 2 FeCl3 }}

Its mild reducing power is of value in organic synthesis.{{cite journal|journal=Org. Synth.|year=1948|volume=28|page=11|doi=10.15227/orgsyn.028.0011|title=o-Aminobenzaldehyde|author1=Lee Irvin Smith|author2=J. W. Opie}} 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.{{Cite web |url=https://cpha.tu.edu.iq/images/%D8%B9%D9%85%D8%B1_%D8%AD%D8%B3%D9%8A%D9%86/ASSAY_OF_FERROUS_SULPHATE__ali_hussein-%D9%85%D8%AD%D9%88%D9%84_1.pdf |archive-url=https://web.archive.org/web/20230929051907/https://cpha.tu.edu.iq/images/%D8%B9%D9%85%D8%B1_%D8%AD%D8%B3%D9%8A%D9%86/ASSAY_OF_FERROUS_SULPHATE__ali_hussein-%D9%85%D8%AD%D9%88%D9%84_1.pdf |archive-date=2023-09-29 |last=Al-Obaidi |first=Ali Hussein Mustafa|title=ASSAY OF FERROUS SULPHATE }}

• Iron supplement
• Iron(III) sulfate (ferric sulfate), the other common simple sulfate of iron
• Copper(II) sulfate
• Ammonium iron(II) sulfate, also known as Mohr's salt, the common double salt of ammonium sulfate with iron(II) sulfate
• Chalcanthum
• Ephraim Seehl, known as an early manufacturer of iron(II) sulfate, which he called 'green vitriol'{{cite book |last=Pryce |first=William |title=Mineralogia Cornubiensis; a Treatise on Minerals, Mines and Mining|location=London|url=https://archive.org/details/b30449170|year=1778|publisher=Phillips|page=[https://archive.org/details/b30449170/page/33 33]}}

{{Reflist}}

{{Commons category|Iron(II) sulfate}}

• {{cite web |url=http://www.chemicalland21.com/industrialchem/inorganic/FERROUS%20SULFATE%20HEPTAHYDRATE.htm |title=Product Information |publisher=Chemical Land21 |date=10 January 2007}}
• {{cite AmCyc |last=Hunt |first=T. Sterry |author-link=T. Sterry Hunt |wstitle=Copperas |short=x}}

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