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Incoming links: Iron(II) carbonate
Potassium iodide is prepared by converting iodine into iodide, then obtaining a pure crystalline potassium salt. One chemical route uses potassium hydroxide; an alternative first forms iron(II) iodide. The account here describes the chemical operations and their balances, without assigning a later pharmacopoeial formula to 1866.
In the alkaline route, iodine reacts with hot potassium hydroxide to give a mixture of potassium iodide and potassium iodate:This is disproportionation: iodine goes from oxidation state to both and . Merely evaporating the solution would leave the unwanted iodate with the iodide. Heating the dried salts with charcoal reduces the iodate to iodide. A limiting balance with carbon monoxide as the carbon oxide product isThe carbon oxides can depend on the reaction conditions; the indispensable chemical point is removal of the iodate's oxygen without losing its iodine. Extraction with water, filtration of insoluble residue, concentration and crystallization then recover potassium iodide. A modern USDA technical evaluation explicitly describes this same reaction sequence; it is corroboration of the chemistry, not evidence that its modern document was the 1866 official recipe.[ref]
Potassium Iodide technical evaluation by United States Department of Agriculture. 2011 report, printed pp. 5–6, Evaluation Question 1, method (a)
In the alternative iron route,followed byThe insoluble iron-containing material is separated and the potassium iodide recovered from the liquid. Oxidation of the ferrous precipitate during working can alter its final form, so the displayed chemical equation is the ideal metathesis balance rather than a guarantee that every collected particle remains pure iron(II) carbonate.
Pure potassium iodide is a colourless or white crystalline ionic solid, commonly in cubic crystals, readily soluble in water and less soluble in ethanol. It has no intrinsic brown iodine colour. Its aqueous solution contains potassium ions and iodide ions. Free iodine dissolves in the presence of iodide through triiodide formation,With silver nitrate, iodide gives yellow silver iodide, sparingly soluble in water and resistant to ordinary aqueous Ammonia:A little chlorine liberates iodine, which gives the characteristic colour with starch:By contrast, iodate contamination can liberate iodine simply on acidification in the presence of iodide:That distinction explains why removal of the iodate is part of preparing a well-defined product, rather than an optional cosmetic step. These are chemical properties and recognition principles, not advice to administer the salt.
The main distinction is between metallic iron, iron(II) preparations, iron(III) preparations, and mixed iron oxides. The old prefixes “proto-” and “per-” ordinarily distinguish the lower and higher oxidation states here. “Peroxide of iron” means the historical iron(III) oxide preparation, not a modern peroxide containing an oxygen–oxygen bond.
The principal medicinal preparations can be grouped by the chemical form supplied:
- Iron(II) sulfate: Ferri sulphas, its granulated form, and its dried form supply iron(II). The ordinary and granulated salts are represented as ; the 1867 dried preparation is represented as . Drying principally changes the water of crystallization, rather than converting iron(II) into iron(III).
- Iron(II) carbonate: Ferri carbonas saccharata, the carbonate pill, and the compound iron mixture contain freshly formed iron(II) carbonate, . Exposure to air introduces higher iron oxides; sucrose helps preserve the lower oxidation state. Thus the saccharated material is a protected preparation rather than a pure, permanently stable salt.
- Iron(II) iodide: Ferri iodidum, its syrup, and its pill supply mainly . The historical solid includes water of crystallization and some iron oxide; the syrup again employs sucrose to retard deterioration.
- Iron(II) phosphate: Ferri phosphas and the corresponding syrup supply the lower iron phosphate, conventionally . The official precipitated material is partly oxidized, so describing the actual powder as chemically pure iron(II) phosphate would be too strong.
- Iron(III) chloride: the solution and tincture of “perchloride of iron” contain the higher chloride, , in an acidic liquid. This historical “perchloride” is unrelated to a perchlorate.
- Iron(III) nitrate: solution of “pernitrate of iron” contains the higher nitrate, , again in an acidic liquid. The solution of “persulphate” used to manufacture several other preparations supplies iron(III) sulfate, ; it is useful to distinguish this manufacturing reagent from the widely administered iron(II) sulfate.
- Iron(III) with organic acids: citrate of iron and ammonium, and tartarated iron, contain the higher iron in citrate or tartrate combinations. Citrate of iron and quinine is prepared from a higher iron compound, but should not be assumed to contain exclusively iron(III): the 1867 monograph reports blue reactions with both potassium ferrocyanide and potassium ferricyanide, consistent with an iron(II) component as well. The tartarated preparation also contains potassium; citrate of iron and quinine is made with Ammonia. They should be described as soluble pharmaceutical combinations, without pretending that every historical batch had a single simple molecular formula. The acetate tincture similarly contains the higher iron combined with acetic acid.
- Iron oxides: Ferri peroxidum hydratum (dried hydrated “peroxide”) and Ferri peroxidum humidum (the moist material) contain iron(III) as iron(III) oxide or hydrated iron(III) oxide. The magnetic oxide is a mixed iron(II)–iron(III) oxide preparation; expresses the ideal mixed oxide, rather than a claim that the historical hydrated powder was pure crystalline magnetite.
- Metallic iron: Ferrum redactum, used also in lozenges, contains finely divided iron in oxidation state zero together with magnetic oxide. The official preparation is therefore not all elemental iron. Iron wine made by contacting wire with wine obtains a dissolved tartrate-containing iron preparation from reaction with the wine's acids; the wire placed in the vessel should not be confused with the chemical form in the liquid ultimately dispensed.
These identifications follow the historical monographs, rather than the names of present commercial supplements.[ref]
British Pharmacopoeia by the General Medical Council (1867). “Ferri carbonas saccharata” through “Ferrum tartaratum”, pp. 130–141; the iron solutions, pp. 186–189; “Vinum ferri”, p. 367
A green iron(II) sulfate solution gives the characteristic blue reaction with potassium ferricyanide, whereas the higher iron solutions give blue with potassium ferrocyanide. Those diagnostic reactions explain why the official tests distinguish the two oxidation states; the historical colour test alone does not establish every detail of a complex preparation's composition.