The specific gravity of the crystals of iodine is 4·95. It melts at 114° and boils at 184°. Its vapour is formed at a much lower temperature, and is of a violet colour, whence iodine receives its name (ἰοειδης, violet). The smell of iodine recalls the characteristic smell of hypochlorous acid; it has a sharp sour taste. It destroys the skin and organs of the body, and is therefore frequently employed for cauterising and as an irritant for the skin. In small quantities it turns the skin brown, but the coloration disappears after a certain time, partly owing to the volatility of the iodine. Water dissolves only1⁄5000part of iodine. A brown solution is thus obtained, which bleaches, but much more feebly than bromine and chlorine. Water which contains salts, and especially iodides, in solution dissolves iodine in considerable quantities, and the resultant solution is of a dark brown colour. Pure alcohol dissolves a small amount of iodine, and in so doing acquires a brown colour, but the solubility of iodine is considerably increased by the presence of a small quantity of an iodine compound—for instance, ethyl iodide—in the alcohol.[63]Ether dissolves a larger amount of iodine than alcohol; but iodine is particularly soluble in liquid hydrocarbons, in carbon bisulphide, and in chloroform. A small quantity of iodine dissolved in carbon bisulphide tints it rose-colour, but in a somewhat larger amount it gives a violet colour. Chloroform (quite free from alcohol) is also tinted rose colour by a small amount of iodine. This gives an easy means for detecting the presence of free iodine in small quantities. The blue coloration which free iodine gives withstarchmay also, as has already been frequently mentioned (seeChapterIV.), serve for the detection of iodine.
If we compare the four elements, fluorine, chlorine, bromine, and iodine, we see in them an example of analogous substances which arrange themselves by their physical properties in the same order asthey stand in respect to their atomic and molecular weights. If the weight of the molecule be large, the substance has a higher specific gravity, a higher melting and boiling point, and a whole series of properties depending on this difference in its fundamental properties. Chlorine in a free state boils at about -35°, bromine boils at 60°, and iodine only above 180°. According to Avogadro-Gerhardt's law, the vapour densities of these elements in a gaseous state are proportional to their atomic weights, and here, at all events approximately, the densities in a liquid (or solid) state are also almost in the ratio of their atomic weights. Dividing the atomic weight of chlorine (35·5) by its specific gravity in a liquid state (1·3), we obtain a volume = 27, for bromine (80/3·1) 26, and for iodine also (127/4·9) 26.[64]
The metallic bromides and iodides are in the majority of cases, in most respects analogous to the corresponding chlorides,[65]but chlorine displaces the bromine and iodine from them, and bromine liberates iodine from iodides, which is taken advantage of in the preparation of these halogens. However, the researches of Potilitzin showed that areversedisplacement of chlorine by bromine may occur both in solutions and in ignited metallic chlorides in an atmosphere of bromine vapour—that is, a distribution of the metal (according to Berthollet's doctrine) takes place between the halogens, although however the larger portion, still unites with the chlorine, which shows its greater affinity for metals as compared with that of bromine and iodine.[66]The latter, however,sometimes behave with respect to metallic oxides in exactly the same manner as chlorine. Gay-Lussac, by igniting potassium carbonate in iodine vapour, obtained (as with chlorine) an evolution of oxygen and carbonic anhydride, K2CO3+ I2= 2KI + CO2+ O, only the reactions between the halogens and oxygen are more easily reversible with bromine and iodine than with chlorine. Thus, at a red heat oxygen displaces iodine from barium iodide. Aluminium iodide burns in a current of oxygen (Deville and Troost), and a similar, although not so clearly marked, relation exists for aluminium chloride, and shows that the halogens have a distinctly smaller affinity for those metals which only form feeble bases. This is still more the case with the non-metals, which form acids and evolve much more heat with oxygen than with the halogens (Note13). But in all these instances the affinity (and amount of heat evolved) of iodine and bromine is less than that of chlorine, probably because the atomic weights are greater.The smaller store of energy in iodine and bromine is seen still more clearly in the relation of the halogens to hydrogen. In a gaseous state they all enter, with more or less ease, into direct combination with gaseous hydrogen—for example, in the presence of spongy platinum, forming halogen acids, HX—but the latter are far from being equally stable; hydrogen chloride is the most stable, hydrogen iodide the least so, and hydrogen bromide occupies an intermediate position. A very high temperature is required to decompose hydrogen chloride even partially, whilst hydrogen iodide is decomposed by light even at the ordinary temperature and very easily by a red heat. Hence the reaction I2+ H2= HI + HI is very easily reversible, and consequently has a limit, and hydrogen iodide easily dissociates.[67]Judging by the direct measurement of the heat evolved (22,000 heat units) in the formation of HCl, the conversion of 2HCl into H2+ Cl2requires the expenditureof 44,000 heat units. The decomposition of 2HBr into H2+ Br2only requires, if the bromine be obtained in a gaseous state, a consumption of about 24,000 units, whilst in the decomposition of 2HI into H2+ I2as vapour about 3,000 heat units areevolved;[68]these facts, without doubt, stand in causal connection with the great stability of hydrogen chloride, the easy decomposability of hydrogen iodide, and the intermediate properties of hydrogen bromide. From this it would be expected that chlorine is capable of decomposing water with the evolution of oxygen, whilst iodine has not the energy to produce this disengagement,[69]although it is able to liberate the oxygen from the oxides of potassium and sodium, the affinity of these metals for the halogens being very considerable. For this reason oxygen, especially in compounds from which it can be evolved readily (for instance, ClHO, CrO3, &c.), easily decomposes hydrogen iodide. A mixture of hydrogen iodide and oxygen burns in the presence of an ignited substance, forming water and iodine. Drops of nitric acid in an atmosphere of hydrogen iodide cause the disengagement of violet fumes of iodine and brown fumes of nitric peroxide. In the presence of alkalis and an excess of water, however, iodine is able to effect oxidation like chlorine—that is,it decomposes water; the action is here aided by the affinity of hydrogen iodide for the alkali and water, just as sulphuric acid helps zinc to decompose water. But the relative instability of hydriodic acid is best seen in comparing the acids in a gaseous state. If the halogen acids be dissolved in water, they evolve so much heat that they approach much nearer to each other in properties. This is seen from thermochemical data, for in the formation of HX in solution (in a large excess of water) from thegaseouselements there isevolvedfor HCl 39,000, for HBr 32,000, and for HI 18,000 heat units.[70]But it is especially evident from the fact that solutions of hydrogen bromide and iodide in water have many points in common with solutions of hydrogen chloride, both in their capacity to form hydrates and fuming solutions of constant boiling point, and in their capacity to form haloid salts, &c. by reacting on bases.
In consequence of what has been said above, it follows thathydrobromic and hydriodic acids, being substances which are but slightly stable, cannot be evolved in a gaseous state under many of those conditions under which hydrochloric acid is formed. Thus if sulphuric acid in solution acts on sodium iodide, all the same phenomena take place as with sodium chloride (a portion of the sodium iodide gives hydriodic acid, and all remains in solution), but if sodium iodide be mixed with strong sulphuric acid, then the oxygen of the latter decomposes the hydriodic acid set free, with liberation of iodine, H2SO4+ 2HI = 2H2O + SO2+ I2. This reaction takes place in the reverse direction in the presence of alarge quantityof water (2,000 parts of water per 1 part of SO2), in which case not only the affinity of hydriodic acid for water is brought to light but also the action of water in directing chemical reactions in which it participates.[71]Therefore, with a halogen salt, it is easy to obtain gaseous hydrochloric acid by the action of sulphuric acid, but neither hydrobromic nor hydriodic acid can be so obtained in the free state (as gases).[72]Other methods have to be resorted to for their preparation, and recourse must not be had to compounds of oxygen, which are so easily able to destroy these acids. Therefore hydrogen sulphide, phosphorus, &c., which themselves easily take up oxygen, are introduced as means for the conversion of bromine and iodine into hydrobromic and hydriodic acids in the presence of water. For example, in the action of phosphorus the essence of the matter is that the oxygen of the water goesto the phosphorus, and the union of the remaining elements leads to the formation of hydrobromic or hydriodic acid; but the matter is complicated by the reversibility of the reaction, the affinity for water, and other circumstances which are understood by following Berthollet's doctrine. Chlorine (and bromine also) directly decomposes hydrogen sulphide, forming hydrochloric acid and liberating sulphur, both in a gaseous form and in solutions, whilst iodine only decomposes hydrogen sulphide in weak solutions, when its affinity for hydrogen is aided by the affinity of hydrogen iodide for water. In a gaseous state iodine does not act on hydrogen sulphide,[73]whilst sulphur is able to decompose gaseous hydriodic acid, forming hydrogen sulphide and a compound of sulphur and iodine which with water forms hydriodic acid.[74]
If hydrogen sulphide be passed through water containing iodine, the reaction H2S + I2= 2HI + S proceeds so long as the solution is dilute, but when the mass of free HI increases the reaction stops, because the iodine then passes into solution. A solution having a composition approximating to 2HI + 4I2+ 9H2O (according to Bineau) does not react with H2S, notwithstanding the quantity of free iodine. Therefore only weak solutions of hydriodic acid can be obtained by passing hydrogen sulphide into water with iodine.[74 bis]
To obtain[75]gaseous hydrobromic and hydriodic acids it is mostconvenient to take advantage of the reactions between phosphorus, the halogens, and water, the latter being present in small quantity (otherwise the halogen acids formed are dissolved by it); the halogen is gradually added to the phosphorus moistened with water. Thus if red phosphorus be placed in a flask and moistened with water, and bromine be added drop by drop (from a tap funnel), hydrobromic acid is abundantly and uniformly disengaged.[76]Hydrogeniodide is prepared by adding 1 part of common (yellow) dry phosphorus to 10 parts of dry iodine in a glass flask. On shaking the flask, union proceeds quietly between them (light and heat being evolved), and when the mass of iodide of phosphorus which is formed has cooled, water is added drop by drop (from a tap funnel) and hydrogen iodide is evolved directly without the aid of heat. These methods of preparation will be at once understood when it is remembered (p.468) that phosphorus chloride gives hydrogen chloride with water. It is exactly the same here—the oxygen of the water passes over to the phosphorus, and the hydrogen to the iodine, thus, PI3+ 3H2O = PH3O3+ 3HI.[77]
In a gaseous form hydrobromic and hydriodic acids are closely analogous to hydrochloric acid; they are liquefied by pressure and cold, they fume in the air, form solutions and hydrates, of constant boiling point, and react on metals, oxides and salts, &c.[78]Only the relativelyeasy decomposability of hydrobromic acid and especially of hydriodic acid, clearly distinguish these acids from hydrochloric acid. For this reason, hydriodic acid acts in a number of cases as a deoxidiser or reducer, and frequently even serves as a means for the transference of hydrogen. Thus Berthelot, Baeyer, Wreden, and others, by heating unsaturated hydrocarbons in a solution of hydriodic acid, obtained their compounds with hydrogen nearer to the limit CnH2n+2or even the saturated compounds. For example, benzene, C6H6, when heated in a closed tube with a strong solution of hydriodic acid, gives hexylene, C6H12. The easy decomposability of hydriodic acid accounts for the fact that iodine does not act by metalepsis on hydrocarbons, for the hydrogen iodide liberated with the product of metalepsis, RI, formed, gives iodine and the hydrogen compound, RH, back again. And therefore, to obtain the products of iodine substitution, either iodic acid, HIO3(Kekulé), or mercury oxide, HgO (Weselsky), is added, as they immediately react on the hydrogen iodide, thus: HIO3+ 5HI = 3H2O + 3I2, or, HgO + 2HI = HgI2+ H2O. From these considerations it will be readily understood that iodine acts like chlorine (or bromine) on ammonia and sodium hydroxide, for in these cases the hydriodic acid produced forms NH4I and NaI. With tincture of iodine or even the solid element, a solution of ammonia immediately forms a highly-explosive solid black product of metalepsis, NHI2, generally known asiodideofnitrogen, although it still contains hydrogen (this was proved beyond doubt by Szuhay 1893), which may be replaced by silver (with the formation of NAgI2): 3NH3+ 2I2= 2NH4I + NHI2. However, the composition of the last product is variable, and with an excess of water NI3seems to be formed. Iodide of nitrogen is just as explosive as nitrogen chloride.[78 bis]In theaction of iodine on sodium hydroxide no bleaching compound is formed (whilst bromine gives one), but a direct reaction is always accomplished with the formation of an iodate, 6NaHO + 3I2= 5NaI + 3H2O + NaIO3(Gay-Lussac). Solutions of other alkalis, and even a mixture of water and oxide of mercury, act in the same manner.[79]This direct formation ofiodic acid, HIO3= IO2(OH), shows the propensity of iodine to give compounds of the type IX5. Indeed, this capacity of iodine to form compounds of a high type emphasises itself in many ways. But it is most important to turn attention to the fact that iodic acid is easily and directly formed by the action of oxidising substances on iodine. Thus, for instance, strong nitric acid directly converts iodine into iodic acid, whilst it has no oxidising action on chlorine.[79 bis]This shows a greater affinity in iodine for oxygen than in chlorine, and this conclusion is confirmed by the fact that iodine displaces chlorine fromits oxygen acids,[80]and that in the presence of water chlorine oxidises iodine.[81]Even ozone or a silent discharge passed through a mixture of oxygen and iodine vapour is able to directly oxidise iodine[82]into iodic acid. It is disengaged from solutions as a hydrate, HIO3, which loses water at 170°, and gives an anhydride, I2O5. Both these substances are crystalline (sp. gr. I2O55·037, HIO34·869 at 0°), colourless and soluble in water;[83]both decompose at a red heat into iodine and oxygen, are in many cases powerfully oxidising—for instance, they oxidise sulphurous anhydride, hydrogen sulphide, carbonic oxide, &c.—form chloride of iodine and water with hydrochloric acid, and with bases form salts, not only normal MIO3, but also acid; for example, KIO3HIO3, KIO32HIO3.[83 bis]With hydriodic acid iodic acid immediately reacts, disengaging iodine, HIO3+ 5HI = 3H2O + 3I2.
As with chlorine, so with iodine, aperiodic acid, HIO4, is formed. This acid is produced in the form of its salts, by the action of chlorine on alkaline solutions of iodates, and also by the action of iodine on chloric acid.[84]It crystallises from solutions as a hydrate containing 2H2O (corresponding with HClO4,2H2O), but as it forms salts containing up to 5 atoms of metals, this water must be counted as water of constitution. Therefore IO(OH)5= HIO4,2H2O corresponds with the highest form of halogen compounds, IX7.[85]In decomposing (at200°) or acting as an oxidiser, periodic acid first gives iodic acid, but it may also be ultimately decomposed.
Compounds formed between chlorine and iodine must be classed among the most interesting halogen bodies.[86]These elements combine together directly with evolution of heat, and formiodine monochloride, ICl, oriodine trichloride, ICl3.[87]As water reacts on these substances, forming iodic acid and iodine, they have to be prepared from dry iodine and chlorine.[88]Both substances are formed in a number of reactions; for example, by the action of aqua regia on iodine, of chlorine on hydriodic acid, of hydrochloric acid on periodic acid, of iodine on potassium chlorate (with the aid of heat, &c.) Trapp obtained iodine monochloride, in beautiful red crystals, by passing a rapid current of chlorine into molten iodine. The monochloride then distils over and solidifies, melting at 27°. By passing chlorine over thecrystals of the monochloride, it is easy to obtain iodine trichloride in orange crystals, which melt at 34° and volatilise at 47°, but in so doing decompose (into Cl2and ClI). The chemical properties of these chlorides entirely resemble those of chlorine and iodine, as would be expected, because, in this instance, a combination of similar substances has taken place as in the formation of solutions or alloys. Thus, for instance, the unsaturated hydrocarbons (for example, C2H4), which are capable of directly combining with chlorine and iodine, also directly combine with iodine monochloride.