Iron(II) Compounds
Iron(II) salts are formed when iron is dissolved in dilute acids except for nitric acid. The most important of them is iron(II) sulfate or green vitriol FeSO4 ·7H2O forming light green crystals well soluble in water. Iron(II) sulfate gradually darkens in air and is simultaneously oxidized from its surface, transforming into the yellow-brown basic salt of iron(III).
Iron(II) sulfate is produced by dissolving steel shavings in 20—30% sulfuric acid:
Fe + H2SO4 = FeSO4 + H2↑.
When green vitriol is heated, water is released and a white mass of the anhydrous salt FeSO4 is obtained. At temperatures above 480 °C, the anhydrous salt decomposes with the liberation of sulfur dioxide and trioxide; the latter forms a heavy white vapour of sulfuric acid in humid air:
2FeSO4 = Fe2O3 +SO2↑ + SO3↑.
When a solution of an iron(II) salt reacts with an alkali, a white precipitate of iron(II) hydroxide Fe(OH)2 is formed, which in the air owing to oxidation rapidly acquires a greenish, and then a brown colour, transforming into iron(III) hydroxide Fe(OH)3:
4Fe(OH)2 +O2 + 2H2O = 4Fe(OH)3.
Anhydrous iron(II) oxide FeO can be prepared as a black readily oxidizable powder by reducing iron(III) oxide with carbon monoxide at 500 °C:
Fe2O3 +CO = 2FeO + CO2.
Alkali metal carbonates precipitate white iron(II) carbonate FeCO3 from solutions of iron(II) salts. In water containing CO2, iron(II) carbonate, like calcium carbonate, partly transforms into the more soluble acid salt Fe(HCO3)2. Iron is contained in natural ferruginous mineral waters in the form of this salt.
Iron(II) salts can readily be transformed into iron(III) salts during their reaction with various oxidizing agents—nitric acid, potassium permanganate, or chlorine, for instance:
6FeSO4 + 2HNO3 + 3H2SO4 = 3Fe2(SO4)3 + 2NO↑ + 4H2O.
10FeSO4 + 2KMnO4 + 8H2SO4 = 5Fe2(SO4)3 + K2SO4 + 2MnSO4 + 8H2O.
Owing to their ability of being readily oxidized, iron(II) salts are often used as reducing agents.
Double salts of general formula M2SO4.FeSO4.6H2O (M = alkali metal or ammonium) can be obtained by crystallisation of solutions containing the appropriate proportions of the two simple salts: an acid solution of the salt with M = NH4 (Mohr's salt, ferrous ammonium sulfate) is considerably less quickly oxidised by the air than is the simple iron(II) sulfate solution, and hence is used in volumetric analysis.
Iron(II)sulfide,
FeS,
may be prepared by heating the elements together, or by precipitation
from an iron(II) solution by sulfide ion; it is a black solid which
is non-stoichiometric, like the oxide. The yellow sulfide FeS2
(made up essentially of Fe2
+
and 
ions) occurs naturally as pyrites.
Cobalt (II) compounds
Salts. In some respects these salts resemble those of iron; the aquo-cation [Co(H2O)6]2+ (pink) occurs in solution and in some solid salts, for example CoSO4.7H2O. However, this aquo cation is less strongly reduced than [Fe(H2O)6]2+ and the water ligands are more readily replaced by other ligands than for iron(II) (see below). [Co(H2O)6]2+ is only slightly acid and a normal, hydrated carbonate CoCO3. 6H2O can be precipitated by the addition of carbonate ion to a simple cobalt(II) salt provided that an atmosphere of carbon dioxide is maintained over the solution.
Cobalt(II) halides can be obtained through direct combination of the elements, or through dehydration of their hydrates. Anhydrous cobalt(II) chloride is blue, and the solid contains octahedrally-coordinated cobalt; the hydrated salt CoCl2. 6H2O is pink, with each cobalt surrounded by four water molecules and two chloride ions in a distorted octahedron.
Cobalt(II) hydroxide is obtained as a precipitate when hydroxide ion is added to a solution containing cobalt(II) ions. The precipitate is often blue, but becomes pink on standing; it dissolves in excess alkali giving the blue [Co(OH)4]2– ion, and in slightly alkaline solution is easily oxidized by air to a brown solid of composition CoO(OH).
Cobalt(II) sulfide is precipitated as a black solid due to addition of sulfide ion to a solution of a cobalt(II) salt in alkaline solution.
M(II) Redox properties
The oxidation state +2 is an intermediate state between characteristic 0 and +3, that is why E(II)compounds can be reducing and oxidizing agents simultaneously.
Oxidition oj M behavior is not typical expert the action of strong oxidants:
NiO + C = Ni + CO,
Ni2+ + 2e = Ni.
Reducing properties in the series Fe(ІІ), Co(ІІ) and Ni(ІІ) decrease significantly due to reducing stability of +3 oxidation state in the series. For instance, Fe(OH)2 moistened with water reacts with air oxygen rapidly:
4Fe(OH)2 + O2 + 2H2O = Fe(OH)3.
Reaction of Со(ОН)2 proceeds very slowly. Even with pure O2 and under increased pressure it is completed for several days. Ni(OH)2 has almost no reaction with O2.
Red-Ox potentials are increased from Fe tо Ni in alkaline medium:
Half-reaction |
Fe(OH)3 + e = Fe(OH)2 + OH- |
Co(OH)3 + e = Co(OH)2 + OH- |
Ni(OH)3 + e = Ni(OH)2 + OH- |
Е0, V |
-1.76 |
-1.43 |
-1.57 |
Halogens (Cl2, Br2) and hypohalites actively oxidize all Е(ОН)2 in alkaline medium:
2Co(OH)2 + NaBrO + H2O = 2Co(OH)3 + NaBr;
2Ni(OH)2 + Cl2 + 2NaOH = 2Ni(OH)3 + 2NaCl.
In acid medium, Ео (Fe3+/Fe2+) = +0.77 V, thus many oxidants can oxidize Fe(ІІ) salts:
10FeSO4 + 2KMnO4 + 8H2SO4 = 2MnSO4 + 5Fe2(SO4)3 + K2SO4 + 8H2O;
6FeSO4 + K2Cr2O7 + 7H2SO4 = Cr2(SO4)3 + 3Fe2(SO4)3 + K2SO4 + 7H2O.
These reactions are used to determine quantitively Fe2+-ions content in aqueous solutions.
Fe(ІІ) salts(as well as Fe(OH)2) are also oxidized by air oxygen:
4FeSO4 + O2 + 2H2O = 4(FeOH)SO4.
But, unlike Fe(OH)2 oxidation, this reaction proceeds moderately.
Со2+-ions are less capable to oxidize in acid medium: Е Со3+/Со2+ = +1,92 V. Thus such oxidants as KMnO4, K2Cr2O7, O2, Cl2, Br2 don’t oxidize Со2+. To realize oxidation, F2, O3 or anode process are applied.
Even the strongest oxidants will not oxidise Ni2+-ions in acid medium.
Compounds + 3