PCLPHUKIC ACID AKD HTDRIODIC ACID* BY F L O R E S C E BUSH
Pome years ago Professor Phelps of Marshall College called the attention of Professor Bancroft to a laboratory experiment which was done regularly by the students at Tufts College in the introductory course. To show a difference in stability between hydrochloric acid and hydriodic acid, potassium chloride crystals were added in one case, and potassium iodide crystals in another case, to concentrated sulphuric acid. K i t h potassium chloride hydrochloric acid is given off, while sulphuric acid oxidizes hydriodic acid to iodine. This was what the experiment mas intended to show, and did show. The interesting thing was that hydrogen sulphide is said to come off first, and sulphur dioxide later. The gas evolved by the reaction was tested immediately with lead acetate paper and a positive test for hydrogen sulphide was obtained. I t seemed desirable to find out why the more strongly reduced compound was apparently produced first. According to lIellorl, Gay-Lussac found that concentrated sulphuric acid and hydriodic acid react according to the equation
HnSO,
+ z H I = SO2 + ~ H 2 0+ I?,
while Soubeiran noted that this reaction reverses when the concentration of the sulphuric acid is sufficiently low. With concentrated sulphuric acid and an excess of hydriodic acid, the reaction product is hydrogen sulphide,
HZSO,
+ 8H1 = HZS + 4Hz + 4Iz
because hydriodic acid reduces sulphur dioxide to hydrogen sulphide. It, should therefore be possible to get either hydrogen sulphide or sulphur dioxide from the reaction between sulphuric acid and hydriodic acid, depending on the relative amounts in the reacting zone. The following experiments were made with the purposes of testing this qualitatively. d small beaker was used, fitted with a cover for stirring and testing gases. Concentrated sulphuric acid was treated with sodium bisulphite until the solution smelled strongly of sulphur dioxide. Potassium iodide crystals were added without stirring and a test for hydrogen sulphide was obtained a t once. I n this case the effective concentration of sulphurous acid was high enough so that it was reduced by the hydriodic acid to some extent. The formation of sulphur in considerable amounts shows that the hydrogen sulphide concentration \vas high enough to cause some of it to react with the excess of sulphurous acid [or sulphur dioxide) according to the equation 802 zH?S = 3s 2H20.
+
+
'This work vas done as part of the senior thesis under Professor Bancroft in the autumn
of 1926. "A Comprehensive Treatise o n Inorganic a n d Theoretical Chemistry," 2 , 204, (1922) .
210
614
FLORESCE BUSH
The experiment was repeated, stirring vigorously the solution which contained an excess of sulphur dioxide. A test for hydrogen sulphide was obtained at once and sulphur was also formed, showing that stirring does not change the reaction qualitatively. Concentrated sulphuric acid was next saturated with an excess of crushed iodine crystals and this mixture was stirred. Potassium iodide crystals were added in small amounts with constant stirring. X test for hydrogen sulphide was obtained immediately. Powdered potassium iodide crystals were added in small amounts to a concentrated sulphuric acid solution which was kept stirred constantly. Sulphur dioxide alone was detected until a very large amount of iodine had formed. Further additions of powdered pot'assium iodide then gave a t,est for hydrogen sulphide. The addition of small crystals and the vigorous stirring kept the hydriodic acid concentration low and prevented the formation of hydrogen sulphide until a sufficiently high concentration of sulphurous acid had been built up. That iodine itself affects the course of the reaction is shown by the following experiment. Concentrated sulphuric acid was saturated with an excess of crushed iodine crystals and was kept stirred vigorously. On addition of potassium iodide crystals in small amounts, a test for hydrogen sulphide was obtained a t once. This may mean that the presence of iodine stabilizes hydriodic acid to some extent, so that the concentration of the latter builds up sufficiently to permit the hydrogen sulphide reaction to take place, or it may mean that iodine catalyzes the reaction between hydriodic acid and sulphurous acid. This result was quite unexpected and calls for further investigation. The fact that hydrogen sulphide is obtained a t once when potassium iodide crystals are added to concentrated sulphuric acid shows that there is something which causes a relatively high concentration of hydriodic acid. Initially we have a high concentration of sulphuric acid and neither sulphur dioxide nor iodine. The difference must occur at the surface of the potassium iodide crystals. With moderately large crystals we shall have a t their surfaces a film of potassium bisulphate, hydriodic acid, sulphuric acid, and potassium iodide, in which the concentration of sulphuric acid may easily not be high enough to oxidize hydriodic acid rapidly, and which may not be high enough to oxidize it a t all. In that case, oxidation will occur when the hydriodic acid in relatively high concentration diffuses out of the surface film.
If that is the case we ought to get no hydrogen sulphide if we stir vigorously or if we make the potassium iodide crystals small enough. I t has already been stated that only sulphurous acid is formed when powdered potassium iodide crystals are added wit,h vigorous stirring. We have also found the same result when we add coarser crystals of potassium iodide with vigorous stirring or very finely powdered crystals without stirring.
SULPHURIC ACID AND HYDRIODIC ACID
61 j
These experiments throw light on some work done by Faraday‘ in 1834 on the electrolysis of sulphuric acid. “On experimenting with sulphuric acid, I foul d no reason to believe that it was by itself a conductor of, or decomposable by, electricity, although I had previously been of that opinion. When very strong it is a much worse conductor than if diluted. If then subjected to the action of a powerful battery, oxygen appears a t the anode, or positive electrde, although much is absorbed, and hydrogen and sulphur appear a t the cathode, or negative electrode. Kow the hydrogen has with me always been pure, not sulphuretted, and has been deficient in proportion to the sulphur present, so that it is evident that when decomposition occurred water must have been decomposed. I endeavoured to make the experiment with anhydrous sulphuric acid; and it appeared to me that, when fused, such acid was not a conductor, nor decomposed; but I had not enough of the dry acid in my possession to allow me to decide the point satisfactorily. M y belief is, that when sulphur appears during the action of the pile on sulphuric acid, it is the result of a secondary action, and that the acid itself is not electrolyzable.” ‘‘Pure liquid sulphurous acid does not conduct nor suffer decomposition by the voltaic current,* but, when dissolved in water, the solution acquires conducting power, and is decomposed, yielding oxygen at the anode, and hydragen and sulphur at the cathode. “A solution containing sulphuric acid in addition to the sulphurous acid, was a better conductor. It gave very little gas a t either electrode: that a t the anode was oxygen, that a t the cathode pure hydrogen. From the cathode also rose a white turbid stream, consisting of diffused sulphur, which soon rendered the whole solution milky. The volumes of gases were in no regular proportions to the quantities evolved from water in the voltameter. I conclude that the sulphurous acid wm not a t all affected by the electric current in any of these cases and that the water present was the only body electrochemically decomposed; that, a t the anode, the oxygen from the water converted the sulphurous acid into sulphuric acid, and a t the cathode, the hydrogen electrically evolved decomposed the sulphurous acid, combining with its oxygen, and setting its sulphur free. I conclude that the sulphur a t the negative electrode was only a secondary result; and, in fact, no part of it was found combined with the small portion of hydrogen which escaped when weak solutions of sulphurous acid were used. “I have already given my reasons for concluding that sulphuric acid is not electrolyzable, ax. not decomposable directly by the electric current, but occasionally suffering by a secondary reaction at the cathode from the hydrogen evolved there. In the year 1800,Davy considered the sulphur from sulphuric acid as the result of the action of the nascent hydrogen.3 In 1804, “Experimental Researches in Electricity,” 1, 201, 223 (1839). See also De la Rive: Bibliotheque Vniverselle. 40, 2 0 5 ; or Quarterly Journal of Science, 27, 407. Sicholson’s Quarterly Journal, 4, 280, 281
616
F L O R E S C E BUSH
Hisinger and Rereelius stated that it was the direct result of the action of the voltaic pile,' an opinion which from that time Da and which has since been commonly received by all opinion requires that I should correct what I have already said of thr tic.composition of sulphuric acid in a former series of these Researches ( j g ? ) : I do not now think that the appearance of the sulphur at the negative ~ l e c trode is an immediatc consequence of electrolytic action." The opinion which Faraday had forrncrly held is to be found on p. 1 6 0 (,1833). "The theory I have advanced accords in a most satisfactory iiianncr with the fact of an element or substance finding its plact, of rcst, or rather of evolution. sometimes at one pole and sonir+nies at the other. Sulphur illustrates this effect very well. When sulphuric acid is dccomposcd by tlic, pile, sulphur is evolved at the negative pole: but when sulphurrt of silver idecomposed in a similar way, then the sulphur a p p w r ~at thc positive polr : and if a hot platina pole be w e d so as to vaporize the sulphur evolved in the latter case, then the relation of that pole to the sulphur R i exactly the saiiic as the relation of the same pole to oxygen upon its iinnicmion in Tvatcr. I n both cases the elenieiit evolved is liberated at the pole, but not retained by i t ; but by virtue of its clastic, unconibinable, and immiscible condition passes away into thp surrounding medium. Thc sulphur is evidently determined in thew opposite directions by its opposite chemical relations t r i oxygen and silver: and it is to such relations generally that 1 h a w referrod all electro-chemical phenomena. Where thry do not csist, no electro-cheniical action can take place. IYhere they are ,strongest, it i; most powrful; whcrc they are reversed. the direction of transfer of the sulxtancr is reversed with them. " K e know now that the shortage of oxygen at the anodr, when sulphuric acid is electrolyzed: is due t o the formation of peranlphuric acid. It ir e to see that sulphuric acid can be reduced a t the cathode to sulphurous acid and to hydrogen sulphide: b u t no direct reduction of sulphur t o hydrogen sulphide is probable. The answer is apparently that the electrolytic reduction is nornial, giving sulphurous acid and hydrogen sulphide, which then react to form wlphur according to the equation, SO2 + 2H2S = 3s
+~€3~0.
This is confirnicti hy the fact that a great deal of sulphur ic fornird, +IS Farnday himself sayc, whm one electrolyzes a sulphurous acid solution to which some sulphuric acid i. added. Il'ith a lead cathodc in *ulphui.ic acid, th(x amount of sulphur fornied in a givrn time decreases niarkedly with rapid I Y XisI *reduced stirring, because then ~ e r ylittle of the requiting ~ L ~ ~ ~ ~ U acid further. It is probable that n.ith a platinuiil cathode and conseqnent1)- :i smaller over-voltagSr, the amount of sulphiu fornied for the same curwnt density could be tlccrcased still more: b u t this cxprrinicnt TYX? not t i k l
S V L P H C R I C A C I D A S D HYDRIODIC ACID
617
Khile concentrated sulphuric acid oxidizes hydriodic acid to iodine, it is vie11 knoivn that in dilute solution' sulphur reduces iodine to hydriodic acid. "Sulphurous acid is oxidised quantitat'ively to sulphuric acid by S 1 0 iodine, without the xparation of sulphur. The intermediate formation of the yellnw coinpound, SO,.HI, which occurs in solutions of moderate concentration, has no influence on the final result. The low results obtained vihen sulphurous acid is exposed to the air during the titration are due entirely to evaporation of sulphur dioxide, the amount of atmospheric oxidation being negligible. "Sodium sulphite solut,ion is inore readily osidised than sulphurous acid; consequentlv atmospheric oxidation is a more disturbing factor when sodiuni sulphite is titrated with iodine. Since the reaction b e t m e n sulphurous acid and iodine is not reversed under the state of dilution obtaining in volumetric analysis, the addition of sodium hydrogen carbonak to neutralise hydriodic acid, as when solutions of arsenious compounds are being titrated, is unnecessary; and Pince a sulphite solution is osidised so quickly, it is not necessary to allow a time intwval for such oxidat,ion to be completed." Since concentrated sulphuric acid reacts with hydriodic acid to form iodine and sulphur dioxide, and since this reaction reverses when the sulphuric acid is more dilute, there must be a concentration of sulphuric acid at which the reaction reverses. As a first approximation fifty percent sulphuric acid is about the limit at which sulphur dioxide will still reduce iodine. I t might be interesting to deterniine this point electroinetrically. From the experiments described in the first part of this paper it is evident that there is a range of concentrations over which hydriodic acid reduces sulphurous acid to hydrogen sulphide, which may then react with exces; sulphur dioxide to give sulphur. That, sulphurous acid or sulphur dioxide .should be a stronger oxidizing agent in a more strongly acid solution is merely another instance of the almost universal fact that acid increases the oxidizing poiver of oxidizing agents and that alkali increases the reducing power of reducing agents, quite regardless whether the oxidizing or reducing agent i-: or i.: not an electrolyte in the ordinary sense of the term. Wardlaw, Carter and Clews? are not satisfied with so simple a statement as the one just given. "The facts that sulphur dioxide reduces most readily in a dilute acid Inetliuni, and that it oxidises most readily in a strong acid ~ n r d i u r n ,may ~ he correlated if oxidation and reduction are explained on an ionic basis, oxidation bring represented by the surrender of positive charge.: and reduction by the transfprence of negative charges. "Sulphur dioxide in aqueous solution is generally regarded as a modcrately weak acid, ionising principally into H', HBO.,',and BO," ions. [It is not JIacaulny: .J. C'hem. Soc.. 121. 5 5 2 i 1922 .J. Chem. S i x . . 117. 124: , 1 9 2 0 , . [ T h e y inem t h a t snlphiir ilii,si(le rrdur,eb or i)sitli.-e. wmcthing el-e-just w h a t t h e y actually *ay.j 8.
ili
t h e oppwite
618
FLORENCE BUSH
clear how this differs in principle from the ionization of sulphuric acid, which is usually classed as a strong acid.] It is in this condition that it reacts as a reducing agent. Thus: 2Fe’ ’ ’
+ SOa + H?O
2Fe’ ’
=
+ SO,” + z H ‘
In strongly acid solution containing a large number of hydrions, the concentration of SOd”ions will be reduced, and, on the above assumption, its power of reducing should be diminished. This is in accordance with the experimental results. ‘‘NOWlet i t be assumed that sulphur dioxide is capable of ionizing to an extremely minute extent as a base, yielding a correspondingly minute amount of sulphur ions. I t has been shown that the sulphoxides, the organic analogues of sulphur dioxide, have basic pr0perties.l This tendency will be all the greater the larger the number of hydrions present in solution. Thus: H20
SO,e OS(OH)2*SO’ H20
SO‘’
H,OS‘
’ ’ ’
’
+
2
(OH)’,
+ z(0H)’.
In view of the large number of hydrions present in the solution, the concentration of hydroxyl ions would be reduced to a very low value and the reaction toward the right favoured. Oxidation is now represented
S’‘ ‘ ’ + 4Fe’ ’ +4 Fe‘ ’ ’ If oxidation takes place due to the ion SO‘ *, zS0’
‘
+ S.
+ 4 Fe’ +4Fe’ . + z S 0 , + H20 +HZS?OS, ‘
zSO
which represents an intermediate stage in the reduction of sulphur dioxide to sulphur. Thiosulphuric acid would break up into sulphur dioxide and sulphur. I t may be observed that the latter h-ypothesis is in many respects a re-statement of the thionyl chloride hypothesis* applied ionically and more generally.” Wardlaw is apparently adopting a view similar to that advanced by Faraday in 1833 and discarded by him in 1834. He ignores the whole question of hydrogen sulphide. The reaction between hydrogen sulphide and sulphur dioxide has been studied recently by matt hew^.^ “To bring about the decomposition of a mixture of sulphur dioxide and hydrogen sulphide in either the gaseous or the liquid state, the addition of a third substance is necessary, and it has been demonstrated that this third substance must be in the liquid phase. There is no rigid relabionship between the values of the dielectric constants of substanccs and their chemical activity as measured by their ability to bring about the interaction of hydrogen sulFromm and Iiaiziss: Ann., 374, 90 (191o); Fromm: 396, 7 j ?Kardlam a n d Clews: J. Chem. Boc., 117, 1093 (1920). j J. Chem. Soc., 1926, 2270.
(1913).
SULPHURIC ACID AND HYDRIODIC ACID
619
phide and sulphur dioxide. The assertion that hydrogen sulphide and sulphur dioxide when in a liquid state react vigorously, even when dry, has been disproved. It is suggested that the activity of a substance in causing decomposition is dependent on the solubility of the two gases in the substance when liquid, or on the solubility of the solid in the liquid mixture of the two gases.” Atmospheric oxygen oxidizes hydriodic acid more rapidly than it does sulphur dioxide. Consequently we may have the oxidation of sulphur dioxide by atmospheric oxygen taking place in two stages:
+ 4HzO + 2H2SOa + 4HI 4HI + = zH20 +
ZSOZ
and
21%=
0 2
212.
This is, therefore, an induced reaction1 with oxygen as the actor, hydriodic acid as the inductor, and sulphur dioxide or sulphurous acid as the acceptor. Since the hydriodic acid is regenerated as long as there is an excess of sulphur dioxide and as long as the sulphuric acid concentration stays below about fifty percent, hydriodic acid catalyzes the reaction between sulphurous acid and oxygen, though not very rapidly. The general results of this paper are: I . Concentrated sulphuric acid oxidizes hydriodic acid to iodine. With a relatively low concentration of hydriodic acid the sulphuric acid is reduced to sulphurous acid; with a relatively high concentration of hydriodic acid the reduction goes to hydrogen sulphide. Sulphur is due to the reaction between hydrogen sulphide and sulphurous acid.
2. When potassium iodide crystals are dropped into concentrated sulphuric acid, the relative amounts of sulphurous acid, hydrogen sulphide : and sulphur can be varied within rather wide limits by changing the sizes of the potassium iodide crystals and the rate of stirring.
3. Iodine crystals seem to catalyze the reduction of sulphurous acid to hydrogen sulphide; but this has not been proved to be the case.
9. The formation of sulphur when concentrated sulphuric acid is electrolyzed is due to the reaction b e h e e n hydrogen sulphide and sulphurous acid. There is no evidence for the existence of a sulphur or a sulphur-oxygen cation. j . Since sulphurous acid reduces iodine to hydriodic acid in weakly, and oxidizes hydriodic acid t o iodine in strongly, acid solutions, there must be a concentration of sulphuric acid above which sulphurous acid does not reduce iodine. This seems to come at about fifty percent sulphuric acid.
6. The change of sulphurous acid from a reducing agent to an oxidizing agent with increasing concentration of sulphuric acid seems to be another Luther and Schiloa: Z. physik. Chem., 46,
j j j
(19033
620
F L O R E S C I : BUSH
instance of the practically general case that a reducing agent is iiiore powcIful in an alkaline solution and that an oxidizing agmt iq more powerful in an acid solution. ;. Since atmospheric oxygen oxidizes hydriodic acid more rapidly than it does sulphurous acid, hydriodic acid acts as a catalyst in the oxidation of sulphurous acid by atmospheric oxygen so long as sulphurous acid is present in excess relatively to the hydriodic acid and so long as the sulphuric acid concentration is well beloiv fifty percent. 8. If we consider the preceding reaction as an induced reaction, oxygen is the actor, hydriodic acid the inductor, and sulphurous acid the acceptor. This is a clean-cut case of catalysis in a homogeneous systeni with the intcrmediate formation of a definite cheriiical subst>ance,iodine. Corriell l7~iiwrsi1~/,
