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• Uses

  Iron(II) sulfide occurs in nature as the minerals magnetkies, troillite and pyrrhotine. The most important application of this compound is in Kipp’s apparatus as a source for laboratory preparation of hydrogen sulfide. It also is used in paints, pigments, and ceramics and lubricant coatings. ;

• Preparation

  Iron(II) sulfide may be synthesized from the elements but the product is contaminated with iron. The reaction is exothermic and the heat of reaction melts iron. Pure sulfide may be obtained by using a slight excess of sulfur: the excess then is distilled off.
  The compound also may be precipitated by treating an aqueous solution of an alkali metal sulfide with that of iron(II) chloride or any iron(II) salt solution:
  S2– (aq) + Fe2+ (aq) → FeS(s)
  Another method of preparation involves passing a mixture of hydrogen sulfide and hydrogen over iron(III) oxide at about 1,000°C:
  Fe2O3 + 2H2S + H2 → 2FeS + 3H2O ;

• Reactions

  Iron(II) sulfide reacts with acids evolving hydrogen sulfide:
  FeS + 2HCl → H2S + FeCl2
  The compound is readily oxidized under moist condition by action of air, forming triiron tetroxide and elemental sulfur:
  3FeS + 2O2 → Fe3O4 + 3S
  The above reaction is exothermic.
  Iron(II) sulfide decomposes to its elements when heated above 1,100°C:
  FeS → Fe + S
  When heated with boiled water, it generates sulfuric acid and hydrogen:
  4FeS + 8H2O + 7O2 → 4H2SO4 + 4H2 + Fe2O3 ;

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