Iron forms many intensely colored compounds, and as a fairly reactive metal, it can displace hydrogen gas from acids and undergoes many exothermic redox reactions. Here are a few of the more common reactions involving iron and its compounds that meet your criteria.
Iron oxides. The best known reaction of iron is rusting, in which iron reacts with oxygen and water to form red hydrated Fe2O3. Wrapping a thermometer bulb in wet steel wool results in a small but noticeable temperature rise; rusting is exothermic. The color of dried Fe2O3 is intense, and it is used as a red pigment in pints, rubber, ceramics, and glass. Burning the dry steel wool in air results in magnetite, FeO·Fe2O3, a black addition compound that is weakly magnetic.
Iron sulfides. Heating iron filings with powdered sulfur results in a strongly exothermic reaction that produces a charcoal-grey substance. While the substance is commonly identified as iron(II) sulfide, FeS, it has a complex structure and is not a stoichiometric compound (see Cotton and Wilkinson for details).
FeS2 (pyrite) is a brassy, shiny mineral sometimes mistaken for gold. It is not iron(IV) sulfide; the iron is in a +2 oxidation state, combined with an S22- ion (the sulfur analog of the peroxide ion). Pyrite burns when heated to form sulfur dioxide and iron(III) oxide:
4 FeS2(s) + 11 O2 2 Fe2O3(s) + 8 SO2(g)
The reaction is sometimes used in industry as a source of sulfur dioxide for the manufacture of sulfuric acid.
Iron is a fairly active metal and can easily displace hydrogen from mineral acid solutions.
It reacts vigorously and exothermically with sulfuric acid to produce iron(II) sulfate:
Fe(s) + H2SO4(aq) FeSO4(aq) + H2(g)
Drying the solution produces green vitriol: blue-green crystals of FeSO4·7 H2O. Air oxidizes iron(II) salts to iron(III), and the crystals are soon crusted with brown iron(III) hydroxides and sulfates. Iron(II) sulfate is used to make writing inks and dyes by reaction with "tannic acid" (a complex mixture of organic acids extracted from tree bark), followed
by air oxidation to make intensely blue-black iron(III) tannates.
Iron displaces hydrogen from hydrochloric acid to form pale green iron(II) chloride:
Fe(s) + 2 HCl(aq) FeCl2(aq) + H2(g)
The chloride crystallizes as FeCl2·4 H2O. Exposure to air gradually oxidizes the iron(II) to FeCl3 and Fe2O3.
Oxidizing FeCl2 with Cl2 produces orange-yellow FeCl3, which has metal-nonmetal bonds with covalent character.
When iron(III) ions are added to a solution containing thiocyanate ions, a series of brilliantly red complexes form (Fe(SCN)3, Fe(SCN)63-, FeSCN+2...) Thiocyanate is a sensitive test for iron(III) only- iron(II) does not cause a color change.
This makes iron(II) thiocyanate useful as a quick test for the presence of peroxides and oxygen, which rapidly oxidize pale green
Fe(SCN)2·3H2O crystals to red iron(III) thiocyanate.
Iron complex cyanides
Both iron(II) and iron(III) ions form very stable complexes with the cyanide ion (in fact, iron is used to complex waste cyanides to render them less toxic.) The ferrocyanide (Fe(CN)64+) and ferricyanide ion (Fe(CN)63+) contain covalent iron-carbon bonds arranged octahedrally around the central iron.
An intensely colored pigment can be made by combining these ions with iron in a different oxidation state:
K4Fe(CN)6(aq) + Fe3+(aq) KFe[Fe(CN)6](s) + 3 K+(aq)
The same pigment can be obtained from the reaction of iron(II) with ferricyanide:
K3Fe(CN)6(aq) + Fe2+(aq) KFe[Fe(CN)6](s) + 2 K+(aq)
The product of the reaction is the blue color ingredient in many artist's pigments, printing inks, and dyes (including Berlin blue, Chinese blue, mineral blue, Paris blue, and Prussian blue). The reaction is also
used in blueprinting. The undeveloped paper is coated with iron(III), ferricyanide ion, and citrate.
When the paper is exposed to light, the citrate reduces the iron(III) to iron(II). Moistening the paper
forms the deep blue pigment.
- F. A. Cotton and G. Wilkinson in Advanced Inorganic Chemistry, Wiley Interscience, New York, 1988.
Author: Fred Senese firstname.lastname@example.org