Iron(II) sulfate

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Iron(II) sulfate
FeSO4.svg
Iron(II)-sulfate-heptahydrate-sample.jpg
Identifiers
PubChem 24393
ChemSpider 22804 YesY
UNII RIB00980VW YesY
EC number 231-753-5
ChEBI CHEBI:75832 N
ChEMBL CHEMBL1200830 N
RTECS number NO8500000
ATC code B03AA07
Jmol-3D images Image 1
Properties
Molecular formula FeSO4
Molar mass 151.908 g/mol (anhydrous)
169.92 g/mol (monohydrate)
278.05 g/mol (heptahydrate)
Appearance blue/green or white crystals
Odor odorless
Density 2.84 g/cm3 (anhydrous)
2.2 g/cm3 (pentahydrate)
1.898 g/cm3 (heptahydrate)
Melting point 70 °C (dehydration of heptahydrate)
400 °C (decomposes)
Solubility in water 25.6 g/100mL (anhydrous)
48.6 g/100 mL (heptahydrate) (50 °C)
Solubility negligible in alcohol
Refractive index (nD) 1.536 (pentahydrate)
1.478 (heptahydrate)
Hazards
EU Index 026-003-00-7 (anhydrous)
026-003-01-4 (heptahydrate)
EU classification Harmful (Xn)
Irritant (Xi)
R-phrases R22, R36/38
S-phrases (S2), S46
NFPA 704
Flammability code 0: Will not burn. E.g., water Health code 1: Exposure would cause irritation but only minor residual injury. E.g., turpentine Reactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g., liquid nitrogen Special hazards (white): no codeNFPA 704 four-colored diamond
Flash point Non-flammable
Thermochemistry
Std molar
entropy
So298
121 J·mol−1·K−1[1]
Std enthalpy of
formation
ΔfHo298
−929 kJ·mol−1[1]
Related compounds
Other cations Cobalt(II) sulfate
Copper(II) sulfate
Manganese(II) sulfate
Nickel(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)
 N (verify) (what is: YesY/N?)
Infobox references

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.

Hydrates[edit]

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)

At 90 °C, the heptahydrate loses water to form the colorless monohydrate. In its anhydrous, crystalline state, its standard enthalpy of formation is ΔfH°solid = -928.4 kJ·mol−1 and its standard molar entropy is S°solid = 107.5 J·K−1·mol−1. 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[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.[2]

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

Reactions[edit]

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 480 °C.

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.

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

Uses[edit]

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.

Nutritional supplement[edit]

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.

Colorant[edit]

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.[3] 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.[4]

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

Other uses[edit]

In horticulture it is used for treating iron chlorosis.[5] 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.[6] It is also used as a lawn conditioner,[6] and moss killer.

In the second half of the 19th century, ferrous sulfate was also 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 a turbine condenser. It forms a corrosion-resistant, protective coating on the inside of the tube.

It is used as a gold refining chemical to precipitate metallic gold from auric chloride solutions (gold that has been dissolved into solution with aqua regia).

It has been applied for 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 panel 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.[7]

See also[edit]

References[edit]

  1. ^ a b Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A22. ISBN 0-618-94690-X. 
  2. ^ 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.
  3. ^ Hyatt, The Interpreter's Bible, 1951, volume V, p. 1067
  4. ^ How To Stain Concrete with Iron Sulfate
  5. ^ Koenig, Rich and Kuhns, Mike: Control of Iron Chlorosis in Ornamental and Crop Plants. (Utah State University, Salt Lake City, August 1996) p.3
  6. ^ 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. 
  7. ^ Svrček, Mirko (1975). A color guide to familiar mushrooms. (2nd ed.). London: Octopus Books. p. 30. ISBN 0-7064-0448-3. 

External links[edit]