Iron(II) sulfate

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Iron(II) sulfate
Fe(H2O)6SO4.png
Iron(II) sulfate when dissolved in water
Iron(II)-sulfate-heptahydrate-3D-balls.tiff
Iron(II)-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
3D model (JSmol)
ChEBI
ChEMBL
ChemSpider
ECHA InfoCard 100.028.867 OOjs UI icon edit-ltr-progressive.svg
EC Number
  • anhydrous:231-753-5
PubChem CID
RTECS number
  • anhydrous:NO8500000 (anhydrous)
    NO8510000 (heptahydrate)
UNII
UN number 3077
  • InChI=1S/Fe.H2O4S/c;1-5(2,3)4/h;(H2,1,2,3,4)/q+2;/p-2 Yes check.svgY
    Key: BAUYGSIQEAFULO-UHFFFAOYSA-L Yes check.svgY
  • anhydrous:InChI=1/Fe.H2O4S/c;1-5(2,3)4/h;(H2,1,2,3,4)/q+2;/p-2
    Key: BAUYGSIQEAFULO-NUQVWONBAS
  • anhydrous:[O-]S(=O)(=O)[O-].[Fe+2]
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)
AppearanceWhite 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 [5]
300 °C (572 °F; 573 K)
(monohydrate) decomposes
60–64 °C (140–147 °F; 333–337 K)
(heptahydrate) decomposes [4] [6]
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) [7]
Solubility Negligible in alcohol
Solubility in ethylene glycol 6.38 g/100 g (20 °C) [5]
Vapor pressure 1.95 kPa (heptahydrate) [8]
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) [9]
1.526–1.528 (21 °C, tetrahydrate) [10]
1.513–1.515 (pentahydrate) [2]
1.468 (hexahydrate) [3]
1.471 (heptahydrate) [11]
Structure
Orthorhombic, oP24 (anhydrous) [12]
Monoclinic, mS36 (monohydrate) [9]
Monoclinic, mP72 (tetrahydrate) [10]
Triclinic, aP42 (pentahydrate) [2]
Monoclinic, mS192 (hexahydrate) [3]
Monoclinic, mP108 (heptahydrate) [4] [11]
Pnma, No. 62 (anhydrous) [12]
C2/c, No. 15 (monohydrate, hexahydrate) [3] [9]
P21/n, No. 14 (tetrahydrate) [10]
P1, No. 2 (pentahydrate) [2]
P21/c, No. 14 (heptahydrate) [11]
2/m 2/m 2/m (anhydrous) [12]
2/m (monohydrate, tetrahydrate, hexahydrate, heptahydrate) [3] [9] [10] [11]
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]
−928.4 kJ/mol (anhydrous) [4]
−3016 kJ/mol (heptahydrate) [13]
−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:
GHS-pictogram-exclam.svg [8]
Warning
H302, H315, H319 [8]
P305+P351+P338 [8]
NFPA 704 (fire diamond)
[16]
NFPA 704.svgHealth 1: Exposure would cause irritation but only minor residual injury. E.g. turpentineFlammability 0: Will not burn. E.g. waterInstability 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no code
1
0
0
Lethal dose or concentration (LD, LC):
237 mg/kg (rat, oral) [6]
NIOSH (US health exposure limits):
REL (Recommended)
TWA 1 mg/m3 [17]
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).
X mark.svgN  verify  (what is  Yes check.svgYX mark.svgN ?)

Iron(II) sulfate (British English: iron(II) sulphate) or ferrous sulfate denotes a range of salts with the formula Fe SO4·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]

Uses

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

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]

Hydrates

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 Siran zeleznaty.PNG
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. [7]

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

Reactions

Iron(II) sulfate outside a titanium dioxide factory in Kaanaa, Pori, Finland. Ferric sulphate, Kemira.jpg
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 FeSO4Fe2O3 + 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

Related Research Articles

<span class="mw-page-title-main">Sulfuric acid</span> Chemical compound (H₂SO₄)

Sulfuric acid or sulphuric acid, known in antiquity as oil of vitriol, is a mineral acid composed of the elements sulfur, oxygen, and hydrogen, with the molecular formula H2SO4. It is a colorless, odorless, and viscous liquid that is miscible with water.

In chemistry, a half reaction is either the oxidation or reduction reaction component of a redox reaction. A half reaction is obtained by considering the change in oxidation states of individual substances involved in the redox reaction. Often, the concept of half reactions is used to describe what occurs in an electrochemical cell, such as a Galvanic cell battery. Half reactions can be written to describe both the metal undergoing oxidation and the metal undergoing reduction.

<span class="mw-page-title-main">Ferrous</span> The element iron in its +2 oxidation state

In chemistry, iron(II) refers to the element iron in its +2 oxidation state. The adjective ferrous or the prefix ferro- is often used to specify such compounds, as in ferrous chloride for iron(II) chloride (FeCl2). The adjective ferric is used instead for iron(III) salts, containing the cation Fe3+. The word ferrous is derived from the Latin word ferrum, meaning "iron".

<span class="mw-page-title-main">Magnesium sulfate</span> Chemical compound with formula MgSO4

Magnesium sulfate or magnesium sulphate is a chemical compound, a salt with the formula MgSO4, consisting of magnesium cations Mg2+ (20.19% by mass) and sulfate anions SO2−4. It is a white crystalline solid, soluble in water but not in ethanol.

<span class="mw-page-title-main">Copper(II) sulfate</span> Chemical compound

Copper(II) sulfate, also known as copper sulphate, is an inorganic compound with the chemical formula CuSO4. It forms hydrates CuSO4·nH2O, where n can range from 1 to 7. The pentahydrate (n = 5), a bright blue crystal, is the most commonly encountered hydrate of copper(II) sulfate, while its anhydrous form is white. Older names for the pentahydrate include blue vitriol, bluestone, vitriol of copper, and Roman vitriol. It exothermically dissolves in water to give the aquo complex [Cu(H2O)6]2+, which has octahedral molecular geometry. The structure of the solid pentahydrate reveals a polymeric structure wherein copper is again octahedral but bound to four water ligands. The Cu(II)(H2O)4 centers are interconnected by sulfate anions to form chains.

<span class="mw-page-title-main">Zinc sulfate</span> Chemical compound

Zinc sulfate describes a family of inorganic compounds with the formula ZnSO4(H2O)x. All are colorless solids. The most common form includes water of crystallization as the heptahydrate, with the formula ZnSO4·7H2O. As early as the 16th century it was prepared on the large scale, and was historically known as "white vitriol" (the name was used, for example, in 1620s by the collective writing under the pseudonym of Basil Valentine). Zinc sulfate and its hydrates are colourless solids.

<span class="mw-page-title-main">Chalcanthite</span> Sulfate mineral

Chalcanthite (from Ancient Greek χάλκανθον (khálkanthon), from χαλκός (khalkós) 'copper', and ἄνθος (ánthos) 'flower, bloom') is a richly colored blue-green water-soluble sulfate mineral CuSO4·5H2O. It is commonly found in the late-stage oxidation zones of copper deposits. Due to its ready solubility, chalcanthite is more common in arid regions.

<span class="mw-page-title-main">Vitriol</span> Index of chemical compounds with the same name

Vitriol is the general chemical name encompassing a class of chemical compounds comprising sulfates of certain metals – originally, iron or copper. Those mineral substances were distinguished by their color, such as green vitriol for hydrated iron(II) sulfate and blue vitriol for hydrated copper(II) sulfate.

In chemistry, water(s) of crystallization or water(s) of hydration are water molecules that are present inside crystals. Water is often incorporated in the formation of crystals from aqueous solutions. In some contexts, water of crystallization is the total mass of water in a substance at a given temperature and is mostly present in a definite (stoichiometric) ratio. Classically, "water of crystallization" refers to water that is found in the crystalline framework of a metal complex or a salt, which is not directly bonded to the metal cation.

In chemistry, disproportionation, sometimes called dismutation, is a redox reaction in which one compound of intermediate oxidation state converts to two compounds, one of higher and one of lower oxidation states. The reverse of disproportionation, such as when a compound in an intermediate oxidation state is formed from precursors of lower and higher oxidation states, is called comproportionation, also known as synproportionation.

<span class="mw-page-title-main">Nickel(II) sulfate</span> Chemical compound

Nickel(II) sulfate, or just nickel sulfate, usually refers to the inorganic compound with the formula NiSO4(H2O)6. This highly soluble blue green coloured salt is a common source of the Ni2+ ion for electroplating.

<span class="mw-page-title-main">Manganese(II) sulfate</span> Chemical compound

Manganese(II) sulfate usually refers to the inorganic compound with the formula MnSO4·H2O. This pale pink deliquescent solid is a commercially significant manganese(II) salt. Approximately 260,000 tonnes of manganese(II) sulfate were produced worldwide in 2005. It is the precursor to manganese metal and many other chemical compounds. Manganese-deficient soil is remediated with this salt.

<span class="mw-page-title-main">Ammonium iron(II) sulfate</span> Chemical compound

Ammonium iron(II) sulfate, or Mohr's salt, is the inorganic compound with the formula (NH4)2Fe(SO4)2(H2O)6. Containing two different cations, Fe2+ and NH+4, it is classified as a double salt of ferrous sulfate and ammonium sulfate. It is a common laboratory reagent because it is readily crystallized, and crystals resist oxidation by air. Like the other ferrous sulfate salts, ferrous ammonium sulfate dissolves in water to give the aquo complex [Fe(H2O)6]2+, which has octahedral molecular geometry. Its mineral form is mohrite.

<span class="mw-page-title-main">Iron(III) sulfate</span> Chemical compound

Iron(III) sulfate (or ferric sulfate), is a family of inorganic compounds with the formula Fe2(SO4)3(H2O)n. A variety of hydrates are known, including the most commonly encountered form of "ferric sulfate". Solutions are used in dyeing as a mordant, and as a coagulant for industrial wastes. Solutions of ferric sulfate are also used in the processing of aluminum and steel.

Tutton's salts are a family of salts with the formula M2M'(SO4)2(H2O)6 (sulfates) or M2M'(SeO4)2(H2O)6 (selenates). These materials are double salts, which means that they contain two different cations, M+ and M'2+ crystallized in the same regular ionic lattice. The univalent cation can be potassium, rubidium, caesium, ammonium (NH4), deuterated ammonium (ND4) or thallium. Sodium or lithium ions are too small. The divalent cation can be magnesium, vanadium, chromium, manganese, iron, cobalt, nickel, copper, zinc or cadmium. In addition to sulfate and selenate, the divalent anion can be chromate (CrO42−), tetrafluoroberyllate (BeF42−), hydrogenphosphate (HPO42−) or monofluorophosphate (PO3F2−). Tutton's salts crystallize in the monoclinic space group P21/a. The robustness is the result of the complementary hydrogen-bonding between the tetrahedral anions and cations as well their interactions with the metal aquo complex [M(H2O)6]2+.

<span class="mw-page-title-main">Cobalt(II) sulfate</span> Inorganic compound

Cobalt(II) sulfate is any of the inorganic compounds with the formula CoSO4(H2O)x. Usually cobalt sulfate refers to the hexa- or heptahydrates CoSO4.6H2O or CoSO4.7H2O, respectively. The heptahydrate is a red solid that is soluble in water and methanol. Since cobalt(II) has an odd number of electrons, its salts are paramagnetic.

<span class="mw-page-title-main">Mohrite</span> Rare ammonium iron(II) sulfate mineral

Mohrite, (NH4)2Fe(SO4)2·6 H2O, is a rare ammonium iron(II) sulfate mineral originally found in the geothermal fields of Tuscany, Italy. This Fe-dominant analogue of boussingaultite is sometimes reported from burning coal dumps where it is a product of pyrite oxidation.

Chvaleticeite is a monoclinic hexahydrite manganese magnesium sulfate mineral with formula: (Mn2+, Mg)[SO4]·6(H2O). It occurs in the oxidized zone of manganese silicate deposits with pyrite and rhodochrosite that have undergone regional and contact metamorphism. It is defined as the manganese dominant member of the hexahydrite group.

Iron(II) selenate (ferrous selenate) is an inorganic compound with the formula FeSeO4. It has anhydrous and several hydrate forms. The pentahydrate has the structure, [Fe(H2O)4]SeO4•H2O, isomorphous to the corresponding iron(II) sulfate. Heptahydrate is also known, in form of unstable green crystalline solid.

References

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  3. 1 2 3 4 5 6 "Ferrohexahydrite Mineral Data" . Retrieved 3 August 2014.
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  10. 1 2 3 4 5 "Rozenite Mineral Data" . Retrieved 3 August 2014.
  11. 1 2 3 4 5 "Melanterite Mineral Data" . Retrieved 3 August 2014.
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  16. ↑ Safety Data Sheet
  17. ↑ NIOSH Pocket Guide to Chemical Hazards. "#0346". National Institute for Occupational Safety and Health (NIOSH).
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  20. ↑ "The Top 300 of 2021". ClinCalc. Archived from the original on 15 January 2024. Retrieved 14 January 2024.
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  22. ↑ British Archaeology magazine. http://www.archaeologyuk.org/ba/ba66/feat2.shtml (archive)
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  24. ↑ "Acid or alkaline soil: Modifying pH - Sunset Magazine". www.sunset.com. 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. 1 2 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.
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  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. 1 2 "Metal-sulfate Salts from Sulfide Mineral Oxidation". pubs.geoscienceworld.org. Retrieved 18 November 2022.
  38. ↑ "Ferrohexahydrite".
  39. ↑ "Melanterite".
  40. ↑ Peterson RC (2003). "THE RELATIONSHIP BETWEEN Cu CONTENT AND DISTORTION IN THE ATOMIC STRUCTURE OF MELANTERITE FROM THE RICHMOND MINE, IRON MOUNTAIN, CALIFORNIA" (PDF).
  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.
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