What are some common reactions of iron and its compounds?

What are some common reactions involving iron compounds which may produce either the release of heat, formation of a gas, or a color change?
Thomas

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.

Common compounds and reactions of iron. Formulas in italics are actually nonstoichiometric compounds!

Compound	Appearance	Reactions
Fe2O3 (hematite, rouge)	rust red to brown	4 Fe(s) + 3 O2(g) 2 Fe2O3(s) (exothermic)
FeO (wüstite)	jet black powder
FeO·Fe2O3 (magnetite)	black to brown	3 Fe(s) + 2 O2(g) FeO·Fe2O3(s) (exothermic)
FeS (iron sulfuret)	gray to gray brown	Fe(s) + S(s) FeS(s) (exothermic) FeS(s) + 2 HCl(aq) H2S(g) + FeCl2(aq)
FeS2 (pyrite, fool's gold)	shiny gold or brassy crystals	4 FeS2(s) + 11 O2 8 SO2(g) + 2 Fe2O3(s)
FeSO4·7H2O (green vitriol)	blue-green crystals	Fe(s) + H2SO4(aq) FeSO4(aq) + H2(g)
Fe2(SO4)3	gray powder	Fe2O3(s) + 3 H2SO4(aq) Fe2(SO4)3(aq) + 3 H2O() Fe2(SO4)3(s) Fe2O3(s) + 3 SO3(g)
FeCl2·4H2O	blue-green crystals	Fe(s) + 2 HCl(aq) FeCl2(aq) + H2(g)
FeCl3·6H2O	orange-yellow crystals	2 FeCl2(aq) + Cl2(g) 2 FeCl3(aq)

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 S₂^2- 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 sulfates. 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 halides. 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.

Iron thiocyanates. When iron(III) ions are added to a solution containing thiocyanate ions, a series of brilliantly red complexes form (Fe(SCN)₃, Fe(SCN)₆^3-, 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.

References

F. A. Cotton and G. Wilkinson in Advanced Inorganic Chemistry, Wiley Interscience, New York, 1988.

Author: Fred Senese senese@antoine.frostburg.edu

Copyright © 1997-2010 by Fred Senese
Comments & questions to fsenese@frostburg.edu
Last Revised 02/23/18.URL: http://antoine.frostburg.edu/chem/senese/101/inorganic/faq/iron-reactions.shtml