Iodine

Iodine (symbol I, atomic weight 126.92), a chemical element, belonging to the halogen group. Its name is derived from Gr. ἰοειδής (violet-coloured), in allusion to the colour of its vapour. It was discovered in 1812 by B. Courtois when investigating the products obtained from the mother-liquors prepared by lixiviating kelp or burnt seaweed, and in 1815 L. J. Gay-Lussac showed that it was an element. Iodine does not occur in nature in the uncombined condition, but is found very widely but sparingly distributed in the form of iodides and iodates, chiefly of sodium and potassium. It is also found in small quantities in sea-water, in some seaweeds, and in various mineral and medicinal springs. Deep-sea weeds as a rule contain more iodine than those which are found in the shallow waters.

Iodine is obtained either from kelp (the ashes of burnt seaweed) or from the mother-liquors obtained in the purification of Chile saltpetre. In the former case the seaweed is burnt in large heaps, care being taken that too high a temperature is not reached, for if the ash be allowed to fuse much iodine is lost by volatilization. The product obtained after burning is known either as kelp or varec. Another method of obtaining kelp is to heat the seaweed in large retorts, whereby tarry and ammoniacal liquors pass over and a very porous residue of kelp remains. A later method consists in boiling the weed with sodium carbonate; the liquid is filtered and hydrochloric acid added to the filtrate, when alginic acid is precipitated; this is also filtered off, the filtrate neutralized by caustic soda, and the whole evaporated to dryness and carbonized, the residue obtained being known as kelp substitute. The kelp obtained by any of these methods is then lixiviated with water, which extracts the soluble salts, and the liquid is concentrated, when the less soluble salts, which are chiefly alkaline chlorides, sulphates and carbonates, crystallize out and are removed. Sulphuric acid is now added to the liquid, and any alkaline sulphides and sulphites present are decomposed, while iodides and bromides are converted into sulphates, and hydriodic and hydrobromic acids are liberated and remain dissolved in the solution. The liquid is run into the iodine still and gently warmed, manganese dioxide in small quantities being added from time to time, when the iodine distils over and is collected. In the second method it is found that the mother-liquors obtained from Chile saltpetre contain small quantities of sodium iodate NaIO3; this liquor is mixed with the calculated quantity of sodium bisulphite in large vats, and iodine is precipitated:—

2NaIO3 \+ 5NaHSO3 = 3NaHSO4 \+ 2Na2SO4 \+ H2O + I2.

The precipitate is washed and then distilled from iron retorts. Iodine may also be prepared by the decomposition of an iodide with chlorine, or by heating a mixture of an iodide and manganese dioxide with concentrated sulphuric acid. Commercial iodine may be purified by mixing it with a little potassium iodide and then subliming the mixture; in this way any traces of bromine or chlorine are removed. J. S. Stas recommends solution of the iodine in potassium iodide and subsequent precipitation by the addition of a large excess of water, the precipitate being washed, distilled in steam, and dried in vacuo over solid calcium nitrate, and then over solid caustic baryta.

Iodine is a greyish-black shining solid, possessing a metallic lustre and having somewhat the appearance of graphite. Its specific gravity is 4.948 (17°/4°). It melts at 114.2° C. and boils at 184.35° C. under atmospheric pressure (W. Ramsay and S. Young). The specific heat of solid iodine is 0.0541 (H. Kopp). Its latent heat of fusion is 11.7 calories, and its latent heat of vaporization is 23.95 calories (P. A. Favre and J. T. Silbermann). The specific heat of iodine vapour at constant pressure is 0.03489, and at constant volume 0.02697. It volatilizes slowly at ordinary temperatures, but rapidly on heating. Iodine vapour on heating passes from a violet colour to a deep indigo blue; this behaviour was investigated by V. Meyer (Ber., 1880, 13, p. 394), who found that the change of colour was accompanied by a change of vapour density. Thus, the density of air being taken as unity, Victor Meyer found the following values for the density of iodine vapour at different temperatures:—

T° C. | 253 | 450 | 506 | 842 | 1027 | 1570  
---|---|---|---|---|---|---  
Density | 8.89 | 8.84 | 8.73 | 6.08 | 5.75 | 5.67  
  
This shows that the iodine molecule becomes less complex in structure at higher temperatures.

Iodine possesses a characteristic penetrating smell, not so pungent, however, as that of chlorine or bromine. It is only very sparingly soluble in water, but dissolves readily in solutions of the alkaline iodides and in alcohol, ether, carbon bisulphide, chloroform, and many liquid hydrocarbons. Its solutions in the alkaline iodides and in alcohol and ether are brown in colour, whilst in chloroform and carbon bisulphide the solution is violet. It appears to combine with the solvent (P. Waentig, Zeit. phys. Chem., 1909, p. 513). Its chemical properties closely resemble those of chlorine and bromine; its affinity for other elements, however, is as a rule less than that of either. It will only combine with hydrogen in the presence of a catalyst, but combines with many other elements directly; for example, phosphorus melts and then inflames, antimony burns in the vapour, and mercury when heated with iodine combines with it rapidly. It is completely oxidized to iodic acid when boiled with fuming nitric acid. It is soluble in a solution of caustic potash, a dilute solution most probably containing the hypoiodite, which, however, changes slowly into iodate, the change taking place rapidly on warming. When alkali is added to aqueous iodine, followed immediately by either soda water or sodium bicarbonate, most of the original iodine is precipitated (R. L. Taylor, Jour. Chem. Soc., 1897, 71, p. 725, and K. J. P. Orton, ibid. p. 830). Iodine can be readily detected by the characteristic blue coloration that it immediately gives with starch paste; the colour is destroyed on heating, but returns on cooling provided the heating has not been too prolonged. Iodine in the presence of water frequently acts as an oxidizing agent; thus arsenious acid and the arsenites, on the addition of iodine solution, are converted into arsenic acid and arsenates. A dilute solution of iodine prevents the decomposition of hydrogen peroxide by colloidal platinum (G. Bredig, Zeit. phys. Chem., 1899, 31, p. 258; 1901, 37, p. 323).

Iodine finds application in organic chemistry, forming addition products with unsaturated compounds, the combination, however, being more slow than in the case of chlorine or bromine. It rarely substitutes directly, because the hydriodic acid produced reverses the reaction; this can be avoided by the presence of precipitated mercuric oxide or iodic acid, which react with the hydriodic acid as fast as it is formed, and consequently remove it from the reacting system. As a rule it is preferable to use iodine in the presence of a carrier, such as amorphous phosphorus or ferrous iodide or to use it with a solvent. It is found that most organic compounds containing the grouping CH3·CO·C— or CH3·CH(OH)·C— in the presence of iodine and alkali give iodoform CHI3.

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).

In medicine iodine is frequently applied externally as a counter-irritant, having powerful antiseptic properties. In the form of certain salts iodine is very widely used, for internal administration in medicine and in the treatment of many conditions usually classed as surgical, such as the bone manifestations of tertiary syphilis. The most commonly used salt is the iodide of potassium; the iodides of sodium and ammonium are almost as frequently employed, and those of calcium and strontium are in occasional use. The usual doses of these salts are from five to thirty grains or more. Their pharmacological action is as obscure as their effects in certain diseased conditions are consistently brilliant and unexampled. Our ignorance of their mode of action is cloaked by the term deobstruent, which implies that they possess the power of driving out impurities from the blood and tissues. Most notably is this the case with the poisonous products of syphilis. In its tertiary stages—and also earlier—this disease yields in the most rapid and unmistakable fashion to iodides; so much so that the administration of these salts is at present the best means of determining whether, for instance, a cranial tumour be syphilitic or not. No surgeon would think of operating on such a case until iodides had been freely administered and, by failing to cure, had proved the disease to be non-syphilitic. Another instance of this deobstruent power—“alterative,” it was formerly termed—is seen in the case of chronic lead poisoning. The essential part of the medicinal treatment of this condition is the administration of iodides, which are able to decompose the insoluble albuminates of lead which have become locked up in the tissues, rapidly causing their degeneration, and to cause the excretion of the poisonous metal by means of the intestine and the kidneys. The following is a list of the principal conditions in which iodides are recognized to be of definite value: metallic poisonings, as by lead and mercury, asthma, aneurism, arteriosclerosis, angina pectoris, gout, goitre, syphilis, haemophilia, Bright’s disease (nephritis) and bronchitis.

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.

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