Iodine

Hydriodic acid, HI, is formed by the direct union of its components in the presence of a catalytic agent; for this purpose platinum black is used, and the hydrogen and iodine vapour are passed over the heated substance. On shaking up iodine with a solution of sulphuretted hydrogen in water, a solution of hydriodic acid is obtained, sulphur being at the same time precipitated. The acid cannot be prepared by the action of concentrated sulphuric acid on an iodide on account of secondary reactions taking place, which result in the formation of free iodine and sulphur dioxide. The usual method is to make a mixture of amorphous phosphorus and a large excess of iodine and then to allow water to drop slowly upon it; the reaction starts readily, and the gas obtained can be freed from any admixed iodine vapour by passing it through a tube containing some amorphous phosphorus. It is a colourless sharp-smelling gas which fumes strongly on exposure to air. It readily liquefies at 0° C. under a pressure of four atmospheres, the liquefied acid boiling at −34.14° C. (730.4 mm.); it can also be obtained as a solid melting at −50.8° C. It is readily soluble in water, one volume of water at 10° C. dissolving 425 volumes of the acid. The saturated aqueous solution is colourless and fumes strongly on exposure to air; after a time it darkens in colour owing to liberation of iodine. The gas is readily decomposed by heat into its constituent elements. It is a powerful reducing agent, and is frequently employed for this purpose in organic chemistry; thus hydroxy acids are readily reduced on heating with the concentrated acid, and nitro compounds are reduced to amino compounds, &c. It is preferable to use the acid in the presence of amorphous phosphorus, for the iodine liberated during the reduction is then utilized in forming more hydriodic acid, and consequently the original amount of acid goes much further. It forms addition compounds with unsaturated compounds.

It has all the characteristics of an acid, dissolving many metals with evolution of hydrogen and formation of salts, called iodides. The iodides can be prepared either by direct union of iodine with a metal, from hydriodic acid and a metal, oxide, hydroxide or carbonate, or by action of iodine on some metallic hydroxides or carbonates (such as those of potassium, sodium, barium, &c.; other products, however, are formed at the same time). The iodides as a class resemble the chlorides and bromides, but are less fusible and volatile. Silver iodide, mercurous iodide, and mercuric iodide are insoluble in water; lead iodide is sparingly soluble, whilst most of the other metallic iodides are soluble. Strong heating decomposes the majority of the iodides. Nitrous acid and chlorine readily decompose them with liberation of iodine; the same effect being produced when they are heated with concentrated sulphuric acid and manganese dioxide. The soluble iodides, on the addition of silver nitrate to their nitric acid solution, give a yellow precipitate of silver iodide, which is insoluble in ammonia solution. Hydriodic acid and the iodides may be estimated by conversion into silver iodide.

Iodine combines with chlorine to form iodine monochloride, ICl, which may be obtained by passing dry chlorine over dry iodine until the iodine is completely liquefied, or according to R. Bunsen by boiling iodine with aqua regia and extracting with ether. It exists in two different crystalline forms, the more stable or α form melting at 27.2° C., and the less stable or β form melting at 13.9° C. It is readily decomposed by water. The trichloride, ICl3, results from the action of excess of chlorine on iodine, or from iodic acid and hydrochloric acid, or by heating iodine pentoxide with phosphorus pentachloride. It crystallizes in long yellow needles and decomposes readily on heating into the monochloride and chlorine. It is readily soluble in water, but excess of water decomposes it. (See W. Stortenbeker, Zeit. phys. Chem., 1889, 3, p. 11.) Iodine monochloride in glacial acetic acid solution was used by A. Michael and T. H. Norton (Ber., 1876, 9, p. 1752) for the preparation of paraiodo-acetanilide.

Iodine Pentoxide, I2O5, the best-known oxide, is obtained as a white crystalline solid by heating iodic acid to 170° C.; it is easily soluble in water, combining with the water to regenerate iodic acid; and when heated to 300° C. it breaks up into its constituent elements, (see M. Guichard, Compt. rend., 1909, 148, p. 925.) Iodine dioxide, I2O4, obtained by Millon, and reinvestigated by M. M. P. Muir (Jour. Chem. Soc., 1909, 95, p. 656), is a lemon-yellow solid obtained by acting on iodic acid with sulphuric acid, oxygen being evolved. By acting with ozone on a chloroform solution of iodine, F. Fichter and F. Rohner (Ber., 1909, 42, p. 4093) obtained a yellowish white oxide, of the formula I4O9, which they regard as an iodate of tervalent iodine, Millon’s oxide being considered a basic iodate.

Although hypoiodous acid is not known, it is extremely probable that on adding iodine or iodine monochloride to a dilute solution of a caustic alkali, hypoiodites are formed, the solution obtained having a characteristic smell of iodoform, and being of a pale yellow colour. It oxidizes arsenites, sulphites and thiosulphates immediately. The solution is readily decomposed on the addition of sodium or potassium bicarbonates, with liberation of iodine. The hypoiodite disappears gradually on standing, and rapidly on warming, being converted into iodate (see R. L. Taylor, Jour. Chem. Soc., 1897, 71, p. 725, and K. J. P. Orton, ibid. p. 830). The peculiar nature of the action between iodine and chlorine in aqueous solution has led to the suggestion that the product is a base, i.e. iodine hydroxide. Tri-iodine hydroxide, I3·OH, is obtained by oxidizing potassium iodide with sulphuric acid and potassium permanganate (A. Skrabal and F. Buchter, Chem. Zeit., 1909, 33, pp. 1184, 1193).

Iodic Acid, HIO3, can be prepared by dissolving iodine pentoxide in water; by boiling iodine with fuming nitric acid, 6I + 10HNO3 = 6HIO3 + 10NO + 2H2O; by decomposing barium iodate with the calculated quantity of sulphuric acid, previously diluted with water, or by suspending iodine in water and passing in chlorine, I2 + 5Cl2 + 6H2O = 2HIO3 + 10HCl. It is a white crystalline solid, easily soluble in water, the solution showing a strongly acid reaction with litmus; the colour, however, is ultimately discharged by the bleaching power of the compound. It is a most powerful oxidizing agent, phosphorus being readily oxidized to phosphoric acid, arsenic to arsenic acid, silicon at 250° C. to silica, and hydrochloric acid to chlorine and water. It is readily reduced, with separation of iodine, by sulphur dioxide, hydriodic acid or sulphuretted hydrogen, thus:—

HIO3 + 5HI = 3H2O + 3I2; 2HIO3+5SO2 + 4H2O = 5H2SO4 + I2;
2HIO3 + 5H2S = I2 + 5S + 6H2O.

The salts, known as the iodates, can be prepared by the action of the acid on a base, or sometimes by the oxidation of iodine in the presence of a base. They are mostly insoluble or only very slightly soluble in water. The iodates of the alkali metals are, however, readily soluble in water (except potassium iodate). They are more easily reduced than the corresponding chlorates; an aqueous solution of hydriodic acid giving free iodine and a metallic oxide, whilst aqueous hydrochloric acid gives iodine trichloride, chlorine, water and a chloride. They are decomposed on heating, with liberation of oxygen, in some cases leaving a residue of iodide and in others a residue of oxide of the metal, with liberation of iodine as well as of oxygen.

Periodic Acid, HIO4·2H2O, is only known in the hydrated form. It can be prepared by the action of iodine on perchloric acid, or by boiling normal silver periodate with water: 2AgIO4 + 4H2O = Ag2H3IO6 + HIO4·2H2O. It is a colourless, crystalline, deliquescent solid which melts at 135° C., and at 140° C. is completely decomposed into iodine pentoxide, water and oxygen. The periodates are a very complex class of salts, and may be divided into four classes, namely, meta-periodates derived from the acid HIO4; meso-periodates from HIO4·H2O, para-periodates from HIO4·2H2O and the diperiodates from 2HIO4·H2O (see C. Kimmins, Jour. Chem. Soc., 1887, 51, p. 356).

Iodine has extensive applications in volumetric analysis, being used more especially for the determination of copper.

The atomic weight of iodine was determined by J. S. Stas, from the analysis of pure silver iodate, and by C. Marignac from the determinations of the ratios of silver to iodine, and of silver iodide to iodine; the mean value obtained for the atomic weight being 126.53. G. P. Baxter (Jour. Amer. Chem. Soc., 1904, 26, p. 1577; 1905, 27, p. 876; 1909, 31, p. 201), using the method of Marignac, obtained the value 126.985 (O = 16). P. Köthner and E. Aeuer (Ber., 1904, 37, p. 2536; Ann., 1904, 337, p. 362), who converted pure ethyl iodide into hydriodic acid and subsequently into silver iodide, which they then analysed, obtained the value 126.026 (H = 1); a discussion of this and other values gave as a mean 126.97 (O = 16).

Small quantities of the iodate (KIO3) are a frequent impurity in iodide of potassium, and cause the congeries of symptoms known as iodism. These comprise dyspepsia, skin eruption and the manifestations which are usually identified with a “cold in the head.” In many cases, as in syphilis, aneurism, lead poisoning, &c., the life of the patient depends on the free and continued use of the iodide, and this is best to be accomplished by securing an absolutely pure supply of the salt. Another often successful method of preventing the onset of symptoms of poisoning is to administer small doses of ammonium carbonate with the drug, thereby neutralizing the iodic acid which is liberated in the stomach.