Ammonium Iron(II) Sulfate
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Ammonium Iron II Sulfate
Ammonium iron(II) sulfate
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IUPAC name
Ammonium iron(II) sulfate
Other names
Ferrous ammonium sulfate
Ammonium iron sulfate
Mohr's salt
3D model (JSmol)
ECHA InfoCard 100.030.125 Edit this at Wikidata
EC Number
  • 233-151-8
Fe(SO4)(NH4)2(SO4) (anhydrous)
Fe(SO4)(NH4)2(SO4)·6H2O (hexahydrate)
Molar mass 284.05 g mol-1 (anhydrous)
392.13 g mol-1 (hexahydrous)
Appearance Blue-green solid
Density 1.86 g/cm3
Melting point 100 to 110 °C (212 to 230 °F; 373 to 383 K)
Boiling point Not applicable
269 g/L (hexahydrate)
Safety data sheet Fisher MSDS
GHS pictograms GHS07: Harmful
GHS Signal word Warning
H315, H319, H335
P261, P264, P271, P280, P302+352, P304+340, P305+351+338, P312, P321, P332+313, P337+313, P362, P403+233, P405, P501
NFPA 704 (fire diamond)
Flammability code 0: Will not burn. E.g. waterHealth code 2: Intense or continued but not chronic exposure could cause temporary incapacitation or possible residual injury. E.g. chloroformReactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no codeNFPA 704 four-colored diamond
Related compounds
Related compounds
Ammonium iron(III) sulfate
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Infobox references

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 NH4+, 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.[1] Its mineral form is mohrite.


This compound is a member of a group of double sulfates called Schönites or Tutton's salts. Tutton's salts form monoclinic crystals and have formula M2N(SO4)2.6H2O (M = various monocations). With regards to the bonding, crystals consist of octahedra [Fe(H2O)6]2+ centers, which are hydrogen bonded to sulfate and ammonium.[2]

Structure of ferrous ammonium sulfate with hydrogen bonding network highlighted (N is violet, O is red; S is orange, Fe = large red).

Mohr's salt is named after the German chemist Karl Friedrich Mohr, who made many important advances in the methodology of titration in the 19th century.


In analytical chemistry, this salt is the preferred source of ferrous ions as the solid has a long shelf life, being resistant to oxidation. This stability extends somewhat to solutions reflecting the effect of pH on the ferrous/ferric redox couple. This oxidation occurs more readily at high pH. The ammonium ions make solutions of Mohr's salt slightly acidic, which slows this oxidation process.[1][3] Sulfuric acid is commonly added to solutions to reduce oxidation to ferric iron.

It is used in the Fricke's dosemeter to measure high doses of gamma rays.[4]


Mohr's salt is prepared by dissolving an equimolar mixture of hydrated ferrous sulfate and ammonium sulfate in water containing a little sulfuric acid, and then subjecting the resulting solution to crystallization. Ferrous ammonium sulfate forms light green crystals. This salt when heated ionises to give all cations and anions present in it.


Common impurities include magnesium, nickel, manganese , lead, and zinc, many of which form isomorphous salts.[5]


  1. ^ a b Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
  2. ^ Ephraim, Fritz (1926). Inorganic Chemistry. tr P. C. L. Thorne. London: Gurney and Jackson. pp. 484-485.
  3. ^ "Ammonium Ferrous Sulphate 100 g (Mohr's Salt)". 2012. Retrieved 2013.
  4. ^ Hickman, C.; Lorrain, S.; Barthe, J.R.; Portal, G. (1986). "Use of Mohr's Salt for High Level Gamma Dosimetry (Up to 108 Gy)". Radiation Protection Dosimetry. Oxford Journals. 17 (1-4): 255-257. doi:10.1093/oxfordjournals.rpd.a079818.
  5. ^ Vogel, Arthur I. (1961). A Text-book of Quantitative Inorganic Analysis Including Elementary Instrumental Analysis (3 ed.). Longmans. pp. 281-282.

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