Iron(II) hydroxide

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Iron(II) hydroxide
IUPAC name

Iron(II) hydroxide

Other names

Ferrous hydroxide

Identifiers
CAS number 18624-44-7 YesY
PubChem 10129897
ChemSpider 8305416 YesY
UNII 7JIM5W32UU YesY
Jmol-3D images {{#if:[Fe+2].[OH-].[OH-]|Image 1
Properties
Molecular formula Fe(OH)2
Molar mass 89.86 g/mol
Appearance brown solid
Density 3.4 g/cm3 [1]
Solubility in water 0.00015 g/100 mL (20 °C)
Solubility product, Ksp 8.0 x 10-16[2]
Hazards
EU Index Not listed
Flash point Non-flammable
Related compounds
Related compounds Iron(II) oxide
Iron(III) hydroxide
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Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
Infobox references

Iron(II) hydroxide or ferrous hydroxide is a compound with the formula Fe(OH)2. It is produced when iron(II) ions, from a compound such as iron(II) sulfate, react with hydroxide ions. Iron(II) hydroxide itself is practically white, but even traces of oxygen impart a greenish tinge. If the solution was not deoxygenated and the iron reduced, the precipitate can vary in color starting from green to reddish brown depending on the iron(III) content. This precipitate is also known as "green rust" in the crystal lattice of which iron(II) ions are easily substituted by iron(III) ions produced by its progressive oxidation. In the presence of oxygen the color changes quickly.

Green rust is a powerful reducing agent and also a layer double hydroxide (LDH) capable to sorb anions because of the presence of positive electrical charges borne on its surface. The mineralogical form of green rust is a recently discovered fougerite. All forms of green rust (including fougerite) are more complex and variable than the ideal iron(II) hydroxide compound. The natural analogue of iron(II) hydroxide compound is a very rare mineral amakinite, (Fe,Mg)(OH)2.

Iron(II) hydroxide is poorly soluble (1.43 × 10−3 g/L). It precipitates from the reaction of iron(II) sulfate and hydroxide ions (from a soluble compound containing hydroxide ion):

FeSO4 + 2 OH− → Fe(OH)2 + SO42−

It is also easily formed as an undesirable by-product of other reactions, a.o., in the synthesis of siderite, an iron carbonate (FeCO3), if the crystal growth conditions are poorly controlled (reagent concentrations, addition rate, addition order, pH, pCO2, T, aging time, ...).

Aging and transformation into magnetite

The second tube from the left has almost pure whitish iron(II) hydroxide.

Under anaerobic conditions, the iron(II) hydroxide can be oxidized by the protons of water to form iron(II,III) oxide and molecular hydrogen. This process is described by the Schikorr reaction:

3 Fe(OH)2 → Fe3O4 + H2 + 2 H2O

The well crystallized iron(II,III) oxide (Fe3O4) is thermodynamically more stable than the iron(II) hydroxide.

Uses

This container has a precipitate of dark greenish iron(II) hydroxide in it.

Anions such as selenite and selenate can be easily adsorbed on the positively charged surface of iron(II) hydroxide where they are subsequently reduced by Fe2+. The resulting products are poorly soluble (Se0, FeSe, or FeSe2).

Iron(II) hydroxide has also been investigated as an agent for the removal of toxic selenate and selenite ions from water systems such as wetlands. The iron(II) hydroxide reduces these ions to elemental selenium, which is insoluble in water and precipitates out.[3]

Note: pKsp = 15.097 where p is the -log and Ksp is the Solubility Product Constant. This means it has a low tendency to dissolve, but is not entirely insoluble. An acidic solution would allow this to disassociate more because the H+ would react with the OH- in the compound.

In a basic solution iron(II) hydroxide is the electrochemically active material of the negative electrode of the nickel-iron battery.

See also

References

  1. ↑ Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN 0-07-049439-8
  2. ↑ http://www.gfredlee.com/SurfaceWQ/StummOxygenFerrous.pdf
  3. ↑ Zingaro, Ralph A.; et al. (1997). "Reduction of oxoselenium anions by iron(II) hydroxide". Environment International 23 (3): 299–304. doi:10.1016/S0160-4120(97)00032-9. 

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