THE DECOMPOSITION OF HYDROGEN IODIDE' BY H. AUSTIN TAYLOR
The general problem of chemical reaction and reactivity has received considerable attention during the past decade, and numerous theories2 have been advanced which have attempted with varying success, to account for the many facts already observed. The treatment of reactivity from the point of view of chemical kinetics has met with apparent success in numerous cases but the time is rapidly approaching when it will be essential to verify practically all the data accumulated in the light of more recent advances before they can be used to confirm a theory. The writer has recently3 called attention to the necessity of a modification of the earlier ideas of the general stability of molecules in view of the great influence of polar substances on various reactions, and on this account considerable doubt must be held regarding data with respect to supposedly homogeneous reactions. The recent attempts to obtain a solution of the problem of unimolecular reaction velocities remain open owing to the lack of data wherewith to confirm them. All the reactions which were supposed until recently to be unimolecular have now been shown to be a t least bimolecular and except for examples of radioactive change (for an explanation of which the aid of such theories is unnecessary and inapplicable) not one example of unimolecular decomposition can be cited. Resulting from this and from the extreniely general influence of catalysts, recent work by Rideal, Norrish and others4has led to the conclusion that most chemical reactions are catalytic in nature, and that the number of bimolecular reactions which are known to occur in the gaseous state without a catalyst is rapidly decreasing. Attempts a t a theoretical evaluation of bimolecular reaction velocities have made use in many cases6 of data of reactions which were assumed to be homogeneous gas reactions but which are now being shown not so. I n the attempt by Lewis to calculate the velocity constant of the bimolecular decomposition of hydrogen iodide, although, as Dushman has pointed out, the success of the result in no wise confirms the radiation hypothesis, the agreement between calculated and observed velocity constants is taken as satisfactory. An empirical factor for the distance within which two molecules approach during collision is employed in the deduction, a very slight change in which would cause complete non-agreement between the calculated and observed reaction velocities. I n the calculations of Dushman, the heats of actiContribution from the Laboratory of Physical Chemistry. Princeton University. Trauts: 2. anorg. Chem., 106, 81 (1909); Lewis: J. Chem. Soc., 105, 2330 (1914)' Perrin: Ann. Phys. 11, 5 (1919);Dushman: J. Am. Chem. SOC., 43, 397 (1921). a J. Phys. Chem., 28, 510 (1924). J. Chem. Soc., 123, 697, 3006 (1923). 6Lewis: J. Chem. SOC.,113, 471 (1918);Dushman: loc. cit. 1
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vation a t various temperatures are obtained which show no falling trend as would be expected from the definition of heat of activation, the mean value being compared with that obtained from the average temperature coefficient of the observed velocity constants. The data employed in both cases are those due to Bodensteinl, on the assumption that the velocity constants therein obtained pertain only to the purely gaseous decomposition of hydrogen iodide. One of the main results of this investigation is that in glass vessels this decomposition is not a purely gaseous reaction but is influenced effectively by the surface in contact with the gas phase. Consequently it would seem that the data are inadequate for the purpose required and the agreement obtained between observed and calculated results fortuitous. Nor is it possible at the present time to make use of any data whereby such calculations may be unequivocally tested. I n a forthcoming textbook2 Edgar has considered the mass action equilibria in the following gaseous systems, complete references of which are also given: ( I ) H2+C02 +CO+H20 -+- 2N0 (2) N2+Oz (3) CO+I/202 ----f COZ (4) H2+1/202 ---t. HzO (5) Nz04 ---t. 2NOz ( 6 ) NO2 +NO+1/202 (7) 1 2 -+- 2 1 (8) (CH,COOH)2+ 2 CHsCOOH ( 9 ) I / ~ N ~ + J ; €+ I ~ ”3 s02+1/20, -+- SO3 (11) HCI+1/402 -+- 1 / 2 C 1 ~ + 1 / 2 H ~ 0 (IO)
I n the majority of these reactions no data are available of actual reaction rates wherewith a direct test of the mass action expression may be obtained. I n those cases where data are available it has been shownrecently that such are not reliable if interpreted as purely gaseous reactions, for example in the case of the oxidation of hydrogen to steam and of carbon monoxide to carbon dioxide the influence of water vapour has been indicated previously and no reaction velocity measurements are available in absence of water vapour. Similarly in the synthesis of ammonia, all available velocity data have been obtained in presence of various catalysts. With reference to the more recent reactions studied, RideaP has shown that the formation of hydrogen sulphide is influenced by the surface of the sulphur undergoing reaction, whilst Daniels4 points out that the decomposition of nitrogen pentoxide is autocatalysed by nitrogen dioxide. Z. physik Chem., 29,295 (1899). “Treatise on Physical Chemistry.” Vol. I Chapter VIII. (1924). 3 J. Chem. SOC., 123, 704 (1923). 4 J. Am. Chem. SOC., 42, 1131 (1920); 43, 53 (1921); 44, 757 (1922). 1 2
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The conclusion is to be drawn therefore, that for the aforementioned theories of bimolecular gaseous reactions, as also for the law of mass action as applied to such reactions, no evidence is available whereby a direct test of the deductions may be made.
Experimental The general method adopted to study the decomposition of hydrogen iodide was a dynamic one, pure hydrogen iodide flowing a t a constant rate through a tube kept a t constant temperature, the effluent gases passing into an aqueous solution of potassium iodide, the resulting iodine and hydriodic acid being titrated with standard sodium thiosulphate and potassium hydroxide respectively. From these data the total amount of hydrogen iodide passing in a given period of time (twenty minutes was found to be a convenient time and was employed throughout) and the percentage decomposition could be calculated, as also could the reaction velocity knowing the volume of the tube in the furnace. Pure hydrogen iodide was prepared by treating red phosphorus and iodine with water in the usual manner, purified by passing through moist red phosphorous and through a U-tube kept in an ice bath to remove further condensible products, and finally dried over phosphorus pentoxide. The gas was then collect,edin the ho?der by immersing this in a bath of solid carbon dioxide and ether. When sufficient solid hydrogen iodide had been collected the holder was sealed and kept in the carbon dioxide bath. I n making an experiment the holder was allowed to warm up to o°C and maintained thereat by immersing in an ether bath in melting ice. This developed a pressure in the holder amounting to approximately four atmospheres and by allowing the gas to flow slowly through a side tube having a fine capillary which allowed about twenty cubic centimetres to pass per minute, a constant flow of gas could be maintained as long as any liquid hydrogen iodide remained in the holder. This side tube was connected to the reaction tube inlet by a short piece of rubber tubing. The reaction tube was fixed in the centre of an electric furnace which could be maintained a t constant lines are not obtained (Curves I and 11). -,.o As previously assumed, the velocity constants observed are in all cases made up of a gas and a surface reaction. Since. IO004 however, the amount of gas reaction is very approximately the same in the two cases namely in the empty tube and with FIU.I glass powder, the volumes of the reaction Curve 1, velocity of HI decomposition in empty tube. tubes being almost the same, if the velo- Curve 11, in presence of glass powder. .city constants of the empty tube be sub- Curve 111, velocity of surface reaction. tracted from those with the glass powder the resulting values should be independent of gas reaction and the logarithms plotted against the reciprocal of absolute temperature should give a straight line. That such is the case is seen in Curve I11 shown in the figure. The slope of this line gives the heat of activation of the surface reaction and thence
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the temperature coefficient in the range studied. The va1u.e so obtained is 10,700 calories for the heat of activation, yielding a temperature coefficient for IO' in the region of 490'C of I . I O . From the figures in Table IV it will be seen that for the data with glass powder the apparent temperature coefficient for IO' between 500 and 5zo'C is I . 2 2 whilst that between 400 and 460'C is 1.17 and the decrease would seem due to the increase in relative amount of surface reaction at the lower temperature. The value I . I O obtained above would therefore appear to be a reasonable one in the limiting case of pure surface reaction. I n this connection it was of interest to find in an approximate manner what particular constituent in the glass was the effective catalyst. The possibility that during the early passage of hydrogen iodide through the tube, neutralisation might occur of some alkali in the surface layer, the influence of sodium iodide on the percentage decomposition of hydrogen iodide was studied in a manner similar to that used previously. It was found, however, that the resulting decomposition was comparable to that occurring in the empty tube. Using pure silica, however, which was powdered to the same extent as the glass employed above, the percentage decomposition at 480'c was 3.21, whilst with the glass powder a t the same temperature 5.95 percent was decomposed and in the empty tube 1.24 percent. The catalytic effect of silica therefore is manifest. The greater influence of glass powder may be due to the fact that the silica in glass may assume the role of a promoted catalyst and, whilst the soda itself is inert, the possibility exists of its ability to promote the catalytic influence of the silica adjacent to it in the glass surface. On the other hand the effect may be simply due to differing surfaces in the glass and the silica. As regards the purely gaseous reaction nothing can be said a t present owing t o lack of knowledge as to how the complex data nhove are to be divided. The straight line of surface reaction is of no assistance since the percentage of surface reaction at each temperature is not known. Suffice that, since from the curve the amount of gas reaction in the data for the empty tube, is rapidly increasing, the tangent to the curve at 520'C should give a very approximate measure of the temperature coefficient of the gaseous reaction. The value obtained is approximately 2 . The value of the temperature coefficient obtained by Bodenstein, for what must have been a complex reaction comparable to that in the present experiments with the empty tube was 1.53 in the temperature region employed, that calculated from the data in Table I V between 500 and 520'C has exactly the same value. However, even if it were possible so to divide the complex data actually observed into two parts namely a surface reaction and a gaseous reaction and from this division to calculate the absolute velocities of the two reactions, the evidence is not yet conclusive that the reaction velocity so calculated for the gaseous phase may be interpreted as relating solely to the gaseous bimolecular decomposition of hydrogen iodide. Further experiment may show that even this decomposition in the gas phase is a catalysed reaction and until such is
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proved incorrect the data even if available from a mode of treatment such as that above, may not be used for a strictly rigorous test of expressions for bimolecular gas reaction velocity. Of the many so-called homogeneous gas reactions so far studied but one would appear to be truly so. The formation of phosphorus pentachloride as studied by the author recently (Zoc. cit.) is a true gas reaction being unaffected by the walls of the containing vessel as also by water vapour, but in this case, it was shown too that liquid chlorine and phosphorus trichloride react with violence even a t -78"C, which would preclude the evaluation of the rate of combination. With reference to the hydrogen bromide decomposition it would seem certain by analogy with the foregoing results that in this case also the glass surface of the reaction chamber would considerably affect the rate. It remains only for similar work as herein stated for hydrogen iodide to be repeated with hydrogen bromide to prove such a statement. The same should be true also of the synthesis of hydrogen bromide and iodide from their elementary constituents.
Summary (I). The decomposition of hydrogen iodide has been studied by a dynamic method at temperatnres from 400 to 520OC. The effect of a change in the extent of surface maintaining the vol(2). ume of reactant constant, has been fully demonstrated, and the decomposition shown to be composite of a surface reaction and a reaction in the gas phase. (3). The temperature coefficient of the surface reaction is 1.10 for 10' whilst that for the reaction in the gas phase would seem to be approximately 2 in the temperature range employed. (4). The general bearing of such facts on the recent theoretical considerations of gas kinetics is indicated. Princeton, N . J.
