Jump to content

Iron(II) sulfate: Difference between revisions

From Wikipedia, the free encyclopedia
Browse history interactively
← Previous editNext edit →
Content deleted Content added
VisualWikitext
m Reverted edits by 70.89.216.157 (talk) to last version by DePiep
m Chembox: unknown parameters: fix spelling, replace or remove. See also full parameter list (via AWB script)
Line 54: Line 54:
| ChEMBL_Ref = {{ebicite|changed|EBI}}
| ChEMBL_Ref = {{ebicite|changed|EBI}}
| RTECS = NO8500000 (anhydrous)<br> NO8510000 (heptahydrate)
| RTECS = NO8500000 (anhydrous)<br> NO8510000 (heptahydrate)
| EC number = 231-753-5
| EC_number = 231-753-5
| ATCCode_prefix = B03
| ATCCode_prefix = B03
| ATCCode_suffix = AA07
| ATCCode_suffix = AA07
| UNNumber = 3077
| UNNumber = 3077
| SMILES = [O-]S(=O)(=O)[O-].[Fe+2]
| SMILES = [O-]S(=O)(=O)[O-].[Fe+2]}}
| SMILES_Ref = {{smilescite}}
| InChI = 1/Fe.H2O4S/c;1-5(2,3)4/h;(H2,1,2,3,4)/q+2;/p-2
| InChI = 1/Fe.H2O4S/c;1-5(2,3)4/h;(H2,1,2,3,4)/q+2;/p-2
| InChIKey = BAUYGSIQEAFULO-NUQVWONBAS
| InChIKey = BAUYGSIQEAFULO-NUQVWONBAS
Line 68: Line 67:
}}
}}
|Section2={{Chembox Properties
|Section2={{Chembox Properties
| Fe=1 | S=1 | O=4
| Fe=1 | S=1 | O=4
| Appearance = White crystals (anhydrous)<br> White-yellow crystals (monohydrate)<br> Blue-green crystals (heptahydrate)
| Appearance = White crystals (anhydrous)<br> White-yellow crystals (monohydrate)<br> Blue-green crystals (heptahydrate)
| Odor = Odorless
| Odor = Odorless
Line 104: Line 103:
| HPhrases = {{H-phrases|302|315|319}}<ref name="sigma" />
| HPhrases = {{H-phrases|302|315|319}}<ref name="sigma" />
| PPhrases = {{P-phrases|305+351+338}}<ref name="sigma" />
| PPhrases = {{P-phrases|305+351+338}}<ref name="sigma" />
| EUIndex = 026-003-00-7 (anhydrous)<br> 026-003-01-4 (heptahydrate)
| EUClass = {{Hazchem Xi}} {{Hazchem Xn}}
| EUClass = {{Hazchem Xi}} {{Hazchem Xn}}
| RPhrases = {{R22}}, {{R36/38}}
| RPhrases = {{R22}}, {{R36/38}}
Line 111: Line 109:
| NFPA-F = 1
| NFPA-F = 1
| NFPA-R = 1
| NFPA-R = 1
| NFPA_Ref = <ref name=fca>{{cite web|title = MSDS of Ferrous sulfate heptahydrate|url = https://www.fishersci.ca/viewmsds.do?catNo=I1463|website = https://www.fishersci.ca|publisher = [[Fisher Scientific|Fisher Scientific, Inc]]|place = [[Fair Lawn, New Jersey]]|accessdate = 2014-08-03}}</ref>
| NFPA_ref = <ref name=fca>{{cite web|title = MSDS of Ferrous sulfate heptahydrate|url = https://www.fishersci.ca/viewmsds.do?catNo=I1463|website = https://www.fishersci.ca|publisher = [[Fisher Scientific|Fisher Scientific, Inc]]|place = [[Fair Lawn, New Jersey]]|accessdate = 2014-08-03}}</ref>
| LD50 = 237 mg/kg (rat, oral)<ref name=fca />
| LD50 = 237 mg/kg (rat, oral)<ref name=fca />
| REL = TWA 1 mg/m<sup>3</sup><ref>{{PGCH|0346}}</ref>
| REL = TWA 1 mg/m<sup>3</sup><ref>{{PGCH|0346}}</ref>

Revision as of 23:35, 1 November 2015

Iron(II) sulfate
Skeletal formula of iron(II) sulfate
Sample of iron(II) sulfate anhydrous
Names
IUPAC name
Iron(II) sulfate
Other names
Ferrous sulfate, Green vitriol, Iron vitriol, Copperas, Melanterite, Szomolnokite
Identifiers
3D model (JSmol)
ChEBI
ChEMBL
ChemSpider
ECHA InfoCard 100.028.867 Edit this at Wikidata
EC Number
  • 231-753-5
RTECS number
  • NO8500000 (anhydrous)
    NO8510000 (heptahydrate)
UNII
UN number 3077
  • [O-]S(=O)(=O)[O-].[Fe+2]
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
☒N verify (what is checkY☒N ?)

|Section2=! colspan=2 style="background: #f8eaba; text-align: center;" |Properties

|-

|

| FeO4S

|- | Molar mass

| 151.90 g·mol−1

|- | 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]

|-


|

| Monohydrate:
44.69 g/100 mL (77 °C)
35.97 g/100 mL (90.1 °C)
Heptahydrate:
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] |-

|

| 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] |- |Section3=! colspan=2 style="background: #f8eaba; text-align: center;" |Structure

|-

|

| 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+) |- |Section4=! colspan=2 style="background: #f8eaba; text-align: center;" |Thermochemistry

|-

|

| 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] |- |Section7=! colspan=2 style="background: #f8eaba; text-align: center;" |Hazards

|-

| colspan=2 style="text-align:left; background-color:#eaeaea;" | GHS labelling: |-

|-

| style="padding-left:1em;" |

| GHS07: Exclamation mark[6]

|-

|-

| style="padding-left:1em;" |

| Warning

|-

|-

| style="padding-left:1em;" |

| H302, H315, H319[6]

|-

|-

| style="padding-left:1em;" |

| P305+P351+P338[6]

|- | NFPA 704 (fire diamond)

|

[10]
NFPA 704 four-colored diamondHealth 2: Intense or continued but not chronic exposure could cause temporary incapacitation or possible residual injury. E.g. chloroformFlammability 1: Must be pre-heated before ignition can occur. Flash point over 93 °C (200 °F). E.g. canola oilInstability 1: Normally stable, but can become unstable at elevated temperatures and pressures. E.g. calciumSpecial hazards (white): no code
2
1
1

|-



| colspan=2 style="text-align:left; background-color:#eaeaea;" | Lethal dose or concentration (LD, LC): |-

|-

| style="padding-left:1em;" |

| 237 mg/kg (rat, oral)[10]

|-

|- | colspan=2 style="text-align:left; background-color:#eaeaea;" | NIOSH (US health exposure limits): |-

|-

| style="padding-left:1em;" |

REL (Recommended)

| TWA 1 mg/m3[13]

|-

|- |Section8=! colspan=2 style="background: #f8eaba; text-align: center;" |Related compounds

|-

|

Other cations

| Cobalt(II) sulfate
Copper(II) sulfate
Manganese(II) sulfate
Nickel(II) sulfate |-

|

Related compounds

| Iron(III) sulfate |- }} Iron(II) sulfate (British English: iron(II) sulphate) or ferrous sulfate is a salt 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

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

Iron(II) sulfate heptahydrate, Melanterite (FeSO4·7H2O)

The heptahydrate in solution (water as solvent) transforms to both heptahydrate and tetrahydrate when the temperature reaches 56.6 °C (133.9 °F). Then at 64.8 °C (148.6 °F) they form both 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

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.[14]

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

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.

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

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

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

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.[15] 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.[16]

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

Other uses

In horticulture it is used for treating iron chlorosis.[17] 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.[18] It is also used as a lawn conditioner,[18] 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 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.[19]

See also

References

  1. ^ a b c d e f "Siderotil Mineral Data". http://www.webmineral.com. Retrieved 2014-08-03. {{cite web}}: External link in |website= (help)
  2. ^ a b c d e f "Ferrohexahydrite Mineral Data". http://www.webmineral.com. Retrieved 2014-08-03. {{cite web}}: External link in |website= (help)
  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". http://chemister.ru. Retrieved 2014-08-03. {{cite web}}: External link in |website= (help)
  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". http://www.mindat.org. Mindat.org. Retrieved 2014-08-03. {{cite web}}: External link in |website= (help)
  8. ^ a b c d e "Rozenite Mineral Data". http://www.webmineral.com. Retrieved 2014-08-03. {{cite web}}: External link in |website= (help)
  9. ^ a b c d e "Melanterite Mineral Data". http://www.webmineral.com. Retrieved 2014-08-03. {{cite web}}: External link in |website= (help)
  10. ^ a b c "MSDS of Ferrous sulfate heptahydrate". https://www.fishersci.ca. Fair Lawn, New Jersey: Fisher Scientific, Inc. Retrieved 2014-08-03. {{cite web}}: External link in |website= (help)
  11. ^ a b c d Weil, Matthias (2007). "The High-temperature β Modification of Iron(II) Sulfate". Acta Crystallographica Section E. 63 (12). International Union of Crystallography: i192. doi:10.1107/S160053680705475X. Retrieved 2014-08-03.
  12. ^ a b c d Anatolievich, Kiper Ruslan. "iron(II) sulfate heptahydrate". http://chemister.ru. Retrieved 2014-08-03. {{cite web}}: External link in |website= (help)
  13. ^ NIOSH Pocket Guide to Chemical Hazards. "#0346". National Institute for Occupational Safety and Health (NIOSH).
  14. ^ 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.
  15. ^ Hyatt, The Interpreter's Bible, 1951, volume V, p. 1067
  16. ^ How To Stain Concrete with Iron Sulfate
  17. ^ Koenig, Rich and Kuhns, Mike: Control of Iron Chlorosis in Ornamental and Crop Plants. (Utah State University, Salt Lake City, August 1996) p.3
  18. ^ 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.
  19. ^ Svrček, Mirko (1975). A color guide to familiar mushrooms (2nd ed.). London: Octopus Books. p. 30. ISBN 0-7064-0448-3.
Wikimedia Commons has media related to Iron(II) sulfate.

Template:B03, B05, B06

Retrieved from "https://en.wikipedia.org/w/index.php?title=Iron(II)_sulfate&oldid=688602711"