T H E ULTRAVIOLET ABSORPTION SPECTRA OF AQUEOUS SOLUTIONS OF SULPHUR DIOXIDE AND SOME OF I T S DERIVATIVES BY FREDERICK H. GETMAS
In a paper treating of the relation between the absorption spectra of acids and their salts, Wright1 was the first to direct attention to the fact that aqueous solutions of sulphur dioxide exhibit a well-defined absorption band in the neighborhood of X = 2 76 pp. Notwithstanding the marked absorption shown by solutions of sulphurous acid, the neutral salt, Na2S03,was found to be diactinic. I n commenting on this fact, Wright suggested that, in view of the doubtful constitution of sulphurous acid, further investigation of the behavior of these substances toward ultraviolet radiation would doubtless prove of interest. At the time of publication of Wright’s paper, the subject of the ultraviolet absorption of solutions of sulphur dioxide and various sulphites was being studied by Garrett2 and also by Schaeffer, Niggemann and Koehler3. Garrett measured the extinction coefficients of HzS03, RbHS03, KHS03, NaHS03, (NH,)2S03, K2S205, Xa2S205,CH3.C(0H) (SO&a) .CH3,NaKS03, Na2S03, CzH6. S O z .0 .C2H5, and SO(OC2H5)~, employing a large Hilger quartz spectrograph in conjunction with a rotating sector photometer. The extinction curve for the acid was well-defined and easily reproducible, whereas with the acid sulphites the absorption was found to increase with time, attaining a maximum value in four or five weeks. This behavior was ascribed to the influence of light, since freshly prepared solutions, and solutions which had been kept in the dark, failed to show selective absorption. Metabisulphites were found to behave similarly to the acid sulphites. The normal sulphites of sodium and potassium, as well as the mixed sulphites of these metals, the symmetrical and unsymmetrical diethyl sulphites and acetone sodium hydrogen sulphite failed to exhibit selective absorption, From these facts the author decided that no conclusive evidence as to the constitution of the sulphites can be derived from the study of their absorption spectra. He attributed the absorbing power of sulphurous acid to hydrated sulphur dioxide, SO2 . nH20, where n is indeterminate, Garrett summarized his conclusions by means of the following equations :( I ) Ka2S206 H 2 0 = 2n’aHSO3, (reaction in solution) NaHS03 NaHS03 = Na2S03 H2S03, (light reaction) (2) (3) HzS03 = S 0 2 . n H z 0 ,(reaction in solution) The relative amounts of hydrated sulphur dioxide present in equally concentrated solutions of the various acid sulphites were calculated and these quantities were shown to increase with increasing strength of the base, or, in
+ +
+
Wright: J. Chem. SOC.,105, 669 (1914). *Garrett: J. Chem. Soc., 107, 1324 (1915). a Schaeffer, Niggemann and Koehler: Z. Elektrochem., 21, 181 (1915).
ABSORPTIOK SPECTRA OF SULPHUR DIOXIDE SOLUTIONS
267
other words, the equilibrium represented by equation ( 2 ) shifts toward the right as the strength of the base increases. Schaeffer and his associates studied the ultraviolet absorption of sulphur dioxide, sulphurous acid and several salts and esters of the latter by means of the familiar Hartley-Baly method, A comparison of the spectrum of a solution of sulphur dioxide in water with that of sodium sulphite indicated that the equilibrium in the reaction, S0z HzO HzS03, lies toward the left side of the equation. On neutralizing a solution of sulphur dioxide with sodium hydroxide, the character of the spectrum was found to change from selective to total absorption in the ultraviolet. An excess of sodium hydroxide was found to exert little or no effect on the degree of absorption, which fact led the authors to conclude that sodium sulphite and its ions possess the same absorptive properties, Sodium bisulphite, prepared by adding the theoretical amount of sodium hydroxide to sulphurous acid, was found to exhibit a weak absorption in the ultraviolet resembling that of sulphurous acid. On the other hand, when sodium bisulphite was prepared by adding the correct quantity of sulphuric acid to the normal sulphite, an entirely different absorption spectrum was obtained which, after the lapse of twentyfour hours, became identical with the spectrum observed with bisulphite prepared by the former method. Solutions of potassium metabisulphite, when freshly prepared, were found to give stronger absorption than sulphurous acid, but after forty-eight hours, the two spectra became identical. In a later communication, Schaeffer and Koehlerl, continuing their studies on the ultraviolet absorption of aqueous solutions of sulphur dioxide, arrived at the conclusion that the true absorbent in solutions of sulphurous acid is a hydrate of sulphur dioxide which appears to be more active than sulphur dioxide alone, Since the normal sulphites and the esters, C2H5. S02.0.CzHC and SO(OC2Hj)z, are both diactinic, it is inferred that sulphurous acid per se is non-absorbing both in the ionized and unionized condition. ' The diminution in the absorbing power of solutions of sulphur dioxide on dilution is traced to a shift from right to left in the equilibrium represented by the equation, SO2 .H20 HzS03. Normal sulphites were found to show only end absorption in the extreme ultraviolet and when a solution of a normal sulphite 'was half-neutralized with sulphuric acid, the resulting acid sulphite at first showed only end absorption and then, after a few days, particularly under the influence of light, the band characteristic of sulphur dioxide was observed to develop. On the other hand, when the acid sulphite was formed by mixing solutions of sulphur dioxide and the normal sulphites, the band manifested itself immediately. A solution of potassium metabisulphite, when freshly prepared, was found to exhibit end absorption only, but on standing an absorption band gradually developed. In like manner, sodium methyl sulphite, which in alcoholic solu-
+
a
Schaeffer and Koehler: 2. anorg. Chem., 104,
212
(1918).
268
FREDERICK H. GETMAN
tion showed only end absorption, in aqueous solution gradually developed the sulphur dioxide band due to hydrolysis into sodium bisulphite and methyl alcohol. Schaeffer and Koehler summarize their views as to the equilibrium conditions in aqueous solutions of sulphur dioxide by the following equations :( I ) SOz.'SOz, H z O ~ ( O z S . O H ) H ~ ( O z S ~ O H )H- + (OzS.OH)(OS ' 0 'OH)- 3 (OS.Oz)-HJ(OS.Oz) = H' (2) They regard the ion, HSO,, as unstable and tending to undergo transformation into H2S03 and SO,=, as represented by the equation, 2H SO3- = SO3= Hzs03. The resulting sulphurous acid is then assumed to come into equilibrium with hydrated sulphur dioxide, as indicated by the equation, HzS03 SOz, HzO. I n this manner, the authors account for the gradual development of the absorption band in solutions of sulphurous acid. The view is also expressed that the ion, SO: has a symmetrical structure and that this structure also obtains in the normal sulphites. The equilibria of solutions of the alkali metal bisulphites was also made the subject of special investigation by Baly and Bailey' in connection with a study of the additive compounds formed by sodium and potassium bisulphites bith aldehydes. These authors concluded that in aqueous solutions of bisulphites there is present in the equilibrium mixture a small proportion of metabisulphite which absorbs light of wave-length X = 257pp. In the absence of oxygen, the solutions were found to be stable in light, but if oxygen or air is present, photo-oxidation of the HS03 ion occurs, followed by an ionic rearrangement in which normal sulphate, sulphurous acid and hydrated sulphur The solution was then found to show the band at dioxide are formed. X = 276pp1 characteristic of hydrated sulphur dioxide. The absorption bands of sulphurous acid, hisulphite and sulphite molecules were shown to lie in the extreme ultraviolet, It was furthermore pointed out, that the characteristic ultraviolet frequencies of sulphur dioxide, hydrated sulphur dioxide and potassium metabisulphite are integral multiples of the fundamental rholecular frequency of sulphur dioxide in the infrared, No evidence was found for the existence of isomerism amongst the sulphite molecules. While all the foregoing investigators are agreed that solutions of sulphur dioxide show a characteristic absorption band at X = 276pp and that solutions of bisulphites, on standing in the presence of oxygen, gradually develop a band at X = 276,144 their statements relative to solutions of metabisulphites are conflicting. Thus, Garrett, referring to the gradual development of a band at X = 2 7 6 ~ ~ in solutions of bisulphites, states that this might be attributed to the formation of metabisulphite according to the equation, NaHS03 NaHS03 = NazSz05 HzO. If this be correct, then, quoting from Garrett's paper, "freshly dissolved metabisulphite ought to exhibit strong selective absorption from the moment
+
e
+
a
+
1
Baly and Bailey: J. Chem. SOC., 121, 1813 (1922).
+
+
ABSORPTION SPECTRA OF SULPHUR DIOXIDE SOLUTIONS
269
of solution. It was found, however, not to be the case. Both sodium and potassium metabisulphite solutions behaved exactly similarly to the hydrogen sulphite solutions, and, in fact, reached the same constant maxima which those solutions had given.” Schaeffer and Koehler, on the other hand, found that freshly prepared solutions of potassium metabisulphite exhibited end absorption only, but that on standing a band gradually developed. Baly and Bailey state, that freshly prepared solutions of potassium metabisulphite exhibit a characteristic absorption band a t X = z57,up and that a similar band was observed in freshly prepared solutions of the alkali metal bisulphites. On exposure to light the band a t X = 257,up disappeared and in its place a new band a t X = 276,up gradually developed with marked increase in the absorptive power of the solution. Again, it is by no means clear whether the progressive change undergone by solutions of bisulphites on standing is due to the action of light in the presence of oxygen, or to the action of light alone, With a view to clearing up these uncertainties the following investigation was undertaken.
Preparation of Materials. In preparing the solutions used in these experiments, due regard was had for the importance of purity of both solutes and solvents in spectrographic work. The sulphur dioxide solutions were prepared with gas obtained by the action of sulphuric acid on chemically pure, crystalline sodium sulphite. The evolved gas mas dissolved in distilled water which had just previously been boiled to exclude air and carbon dioxide, precautions being taken to protect the solutions from oxidation after preparation. The solutions of sodium sulphite and bisulphite were prepared in a similar manner from chemically pure salts. Garrett states that he experienced difficulty in obtaining pure samples of sodium and potassium bisulphites and ultimately was forced to prepare the salts by passing sulphur dioxide into dilute solutions of the carbonates. It should be stated that the samples of sodium bisulphite used in this series of experiments were obtained from a well-known chemical manufacturer and upon analysis were found to correspond accurately with the formula of the salt. Considerable difficulty was experienced, however, in obtaining a satisfactory preparation of sodium metabisulphite. After examining several samples of this salt,obtained from different sources,and finding them all more or less impure, the attempt was made to prepare the salt by passing pure, dry sulphur dioxide into a solution of sodium ethylate in absolute ethyl alcohol. TThile a beautiful crystalline product was obtained, it mas found upon analysis to be no purer than the commercial preparations previously examined. Ultimately, an excellent preparation of the potassium salt was obtained from a reliable chemical firm which, upon analysis, was found to conform very closely to the formula, K2S205. The concentration of the solutions was determined in each case by means of the familiar iodine-thiosulphate method.
Apparatus. The extinction coefficients of the various solutions studied were determined by means of a large size Hilger quartz spectrograph in con-
FREDERICK H. GETMAN
270
junction with a Judd Lewis ultraviolet sector-photometer, as described by the author in a previous paper1. Experimental Data. The experimental results are expressed graphically by means of curves in which logarithms of molecular extinction coefficients are plotted against corresponding wave-lengths expressed in millimicrons, pp, log p The molecular extinction coefficient is defined by t'he expression ml
Y
G FIG. I
A 0.05 m Sodium Sulphite B. 0.05 m Sodium Bisulphite C
m Potassium Metabisulphite D 0.005 m Sulphur Dioxide
_ _
h
0.02
FIG.2
where I, and I are the intensities of the incident and transmitted radiation respectively, and where rn denotes the concentration of the solution in grammolecules per liter and I is the length of the column of absorbing medium in ems. The extinction curve of o . o j m NazSOais designated by the letter A in Fig. I : The position of the absorption band is evidently in the extreme ultraviolet, as was pointed out by Baly and Bailey. Curve B is the extinction curve of freshly prepared 0.05 m KaHSO, and, as with the normal salt, the absorption band is found to lie in the extreme ultraviolet. The extinction curve of freshly prepared 0.02 m KzSzO5 is represented by C. Special preGetman: J. Phys. Chem., 29, 853 (1925).
ABSORPTION SPECTRA O F SULPHUR DIOXIDE SOLUTIONS
271
cautions were taken in the preparation of the latter solution to eliminate both air and carbon dioxide. To this end the apparatus shown in Fig. 2 was employed. A dropping funnel, A, of approximately 70 cc. capacity and fitted with a three-way stopcock, B, was connected through C with a flask, D, containing a large volume of freshly boiled, distilled water. A large clear crystal of potassium metabisulphite, E, was introduced into the funnel, the neck of which was then closed by means of a rubber stopper fitted with a right-angled tube F. The latter was connected with a water pump and B was turned so as to establish communication with the solvent in D. When sufficient water had been aspirated through the funnel to dissolve the surface
FIG.3 A 0.025 m Potassium Metabisulphite B 0.020 m Potassium bletabisulphite C 0.005 m Sulphur Dioxide
layer of E and to insure the removal of the resulting solution, the stop-cock was closed and the crystal permitted to dissolve, When the bulk of the crystal had diminished to about that required for a solution of suitable concentra tion for use in the spectrophotometer, the supernatant solution in A was removed by further aspiration and the crystalline residue allowed to dissolve. connection with the pump was then broken and, after allowing enough of the solution to flow out of the funnel to displace the air in G, the absorption cell was filled and the absorption spectrum immediately photographed. Another portion of the solution was analyzed simultaneously in order to determine its concentration. It will be observed that the solution gives a welldefined absorption band a t X = 263pp. The position of the head of this band was checked by repeated experiments. Curve D, Fig. I , is the extinction curve of freshly prepared 0.005 m SO%. The position of the head of the absorption band at X = 276pp is in agreement with the results of Baly, Garrett
FREDERICK H. GETMAN
272
and others. If due precautions were taken to exclude air and carbon dioxide, the curve for dissolved SO2 could be satisfactorily reproduced. On repeating the experiment with solutions of different concentrations, it was found that Beer’s law does not apply t o aqueous solutions of sulphur dioxide. This is in agreement with the observations of Garrett. Further experiments with potassium metabisulphite are summarized in Fig. 3. Curve A represents the results of two series of measurements of the extinction coefficients of 0.025 m KzSz06. Curve B, which is a reproduction
0.d
0.7
1.0
1.6
1.9 FIG.4 A 0.05 m Sodkm BisulyFite R 0.05 m C 0.025 m ” ” D 0.025 m ” ” E 0.025 m ” ” F 0.025 m ” ” G 0.oOj m Sulphur Dioxide
(freshly prepared) (after 18 hours) (after 5 days) (after IO days) (after 20 days) 0.02 m Sucrose (after 30 days)
+
of curve C of Fig. I , serves to illustrate the inapplicability of Beer’s law to aqueous solutions of metabisulphites. Although the foregoing experiments confirm the statement of Baly and Bailey, that solutions of metabisulphites give a characteristic absorption band, we have been unable to check their result as to the location of the band. They state that the wave-length corresponding to the head of the band is X = a 5 7 p p , whereas we find X = 2 6 3 ~ . Since potassium metabisulphite rapidly decomposes in solution with the formation of sulphur dioxide, it follows that its solutions on standing will develop the band characteristic of dissolved sulphur dioxide at X = 2 7 6 ~ . The results of the measurements of the extinction coefficients of solutions of sodium bisulphite are brought together in Fig. 4. Curve A was obtained
ABSORPTION SPECTRA OF SULPHUR DIOXIDE SOLUTIONS
2 73
with freshly prepared 0.05 m NaHS03. After 18 hours the extinction curve of this same solution was found to have changed to that represented by curve €3. As will be seen, the character of the absorption has changed from general or end absorption to selective absorption. In order to follow more closely the progressive change in the composition of solutions of sodium bisulphite with time, a large volume of 0.025 m NaHS03 was prepared, taking the usual precautions to exclude air and carbon dioxide. With this mother solution, five series of experiments were carried out. In the first series, approximately roo cc. portions of the mother solution were introduced into flasks of 12; cc. capacity and, after stoppering, the flasks were exposed to light for varying periods of time, In the second series, a number of 5 0 cc. bottles fitted with ground stoppers were completely filled with the mother solution. After inserting the stoppers, the necks of the bottles were sealed with paraffine and, as in the case of the preceding series, exposed to light for varying periods of time. The third series resembled the first series in every particular except that to the original solution enough cane sugar was added to make the solution 0.02 molar with respect to that solute. The fourth series was similar ot the second, except that the composition of solution was 0.025 m NaHS03 0 . 0 2 m C12H22011.The last series consisted of partially filled flasks of 0.025 m IVaHSO3 which were protected from the action of light. The progressive change in the composition of the solutions was determined both spectroscopically and analytically by examination of the contents of the different flasks and bottles a t intervals of from five to ten days, The analytical procedure consisted in titrating the solutions against I O cc. portions of X/IO iodine to which a few drops of hydrochloric acid had been added to insure complete liberation of combined sulphur dioxide. The results of the analytical examination are given in Table I.
+
Initial titer, Interval (Ik4
IO
TABLE I cc. N/IO I = 9.6 cc. NaHS03 Titer
I
5
.*..
10
12.9
I1
111
IV
v
9.6 9.5
.... .... ....
9.6 9.5
.... ....
...
9.9
24.5
.... 10.2
IO.7
15
....
...
20
27.0
30 50
62.7
9.6 9.7
31.8
9.8 9.7
....
...
....
...
As will be seen from columns I1 and IV of the table, the titer of the solutions from which air had been excluded remains constant, whereas the figures in columns I and I11 increase, indicating gradual oxidation of the solution. That light plays an important part in the process of oxidation is indicated by the figures of the fifth column. It will be remembered that although this solution was in contact with air it was placed in the dark immediately after preparation. After a period of two months the solution
274
FREDERICK H. GETMAN
had not oxidized as much as the same solution had oxidized in ten days under the influence of light. It has been shown by Mathews and Weeks' that cane sugar and a number of other organic compounds act as negative catalysts in the oxidation of solutions of sodium sulphite. That a similar action is exerted by cane sugar when dissolved in a solution of sodium bisulphite is revealed by a comparison of the data of I and I11 columns of the table. Cane sugar was selected from the list of catalysts enumerated by Mathews and Weeks because it was believed that its presence in a solution of sodium bisulphite would produce little or no alteration in the characteristic absorption spectrum of the salt. The progressive change in the absorption of 0 . 0 2 5 m NaHS03 is depicted by the curves of Fig. 4. It is apparent that after standing for ten days the head of the absorption band becomes constant at X = 2 7 6 ~ the ~ ) wave-length characteristic of dissolved sulphur dioxide. The retarding action of cane sugar on the process of oxidation is clearly shown by the fact that 0 . 0 2 5 m NaHS03 twenty days after preparation has greater absorbing power than the same solution containing 0 . 0 2 gram-molecule of cane sugar thirty days after preparation (curves E and F). Discussion of Results. In reviewing the foregoing experimental data it is apparent that aqueous solutions of sodium bisulphite are stable in light provided oxygen is excluded, but if oxygen or air is present, photo-oxidation ensues and, after five days, the solution shows the characteristic absorption band of dissolved sulphur dioxide a t X = 276pp. Since the absorption curve of 0 . 0 2 5 m NaHS03, five days after preparation, corresponds very closely with that of a mixture of equal volumes of 0.05 m SO2 and 0.0 j m KazS04,it may be fairly assumed that the progressive change in the solution under the influence of light consists in the photo-oxidation of the anion of sulphurous acid, HSO;. The successive steps in the process may be represented by the equations :KaHS03 Na' HSO,, (I)
+
+
+
+
N a C HSO, + N a f H+ S O : , (31 The resulting ions then react with the ions of ( I ) in the following manner :Na' H+ S O , = + N a + HSO,--+aNa++ SO,' Hf+ HSO,, (4) H' HSO; H2S03I502,HzO. (5) Freshly prepared solutions of potassium metabisulphite have been found to give a well-defined absorption band a t X = 263pp, provided precautions are taken to exclude air and carbon dioxide. This fact proves that, initially, solutions of metabisulphites and bisulphites are not identical, as has been claimed by several investigators. Solutions of metabisulphites rapidly hydrolyze as indicated by the equilibrium equation, KzSzOs HzO 2KHS08, (6)
+
+
+
+
+
1
+
e
Mathews and Weeks: J. Am. Chem. SOC., 39, 635 (1917).
ABSORPTION SPECTRA OF SULPHUR DIOXIDE SOLUTIONS
275
and the resulting bisulphite then undergoes photo-oxidation, as indicated in the preceding paragraphs. No indication of the existence of a band at X = 257pp, in freshly prepared solutions of sodium bisulphite, as claimed by Baly and Bailey, was obtained, but the fact that they prepared their solutions of sodium bisulphite by halfneutralization of sulphurous acid with sodium hydroxide, whereas a pure dry salt was used in the present investigation may easily account for this discrepancy, In this connection it may be mentioned that although Garrett took extreme care in the preparation of his solutions he failed to record the existence of any band a t X = 257yy in solutions of the bisulphites. Baly and Bailey estimate that in 0.025 m XaHSO3there is probably less than three percent metabisulphite, indicating that the equilibrium in reaction (6) lies toward the right side of the equation. Baly and Bailey apply to the two absorption bands observed by them at X = 2 j 7 p p and X = 2 7 6 pp, a relationship previously established by the senior author1. By means of a mathematical deduction from the quantum theory, Baly showed that if v x is the wave-number of the characteristic infra-red vibration of a compound, then radiant energy can only be reabsorbed at the wave-numbers VI, v,, v3, etc., where v1, Y,, v3, are successive integral multiples of v,. In the present example, SO2 is the absorbing molecule and this has been found to exhibit a characteristic absorption band corresponding to each of the three following states:-(a) gaseous SO2, (b) SOz.HzO, and (c) KzS03 SOZ. ASBaly and Bailey have shown, the fundamental wave-number of gaseous SO2 is 135.1 and therefore, according to the above generalization, the wave-numbers of the three bands should be integral multiples of this fundamental quantity. The following table is taken from their paper.
TABLE I1 Multiple
25 X 135.1 2 7 X 135.1 29 X 135.1
1/X
3377.5 3647.7 3917.9
X (calc)
X (obs.)
Origin of band
296. I P P 274, I 255.2
296. I P P 276.0 257.0
SO, H2O
so2 *
KzSO3. SO2
If the fundamental wave-number of SOz, 135.1, be multiplied by 28 instead of 29, as in the table, the calculated value of the wave-length of potassium metabisulphite becomes X= 264.3pp, which is in quite as good agreement with our observed value, X = 263.0pp, as is Baly and Bailey’s calculated with their observed value. Summary of Results. The results of the foregoing investigation may be briefly summarized as follows :( I ) Aqueous solutions of sulphur dioxide show a characteristic absorption band at X = 276pp. The absorption is attributed to hydrated sulphur dioxide, SO%.HzO, formed from the anion of sulphurous acid. The extinction curve for sulphur dioxide solutions of the same concentration is easily reBaly: Phil. Mag. (6) 27, 632 (1914); 29,
223
(1914); 30, 510 (1915).
2
76
FREDERICK H. GETMAN
producible, provided oxygen and carbon dioxide are excluded, Beer’s law does not apply to aqueous solutions of sulphur dioxide nor to any of its derivatives. (2) Freshly prepared solutions of potassium metabisulphite give a characteristic absorption band a t X = 263pp. On standing, this band disappears and in its place the sulphur dioxide band a t X = 276pp1 gradually develops. The fact that metabisulphite solutions give a well-defined band at X = 2 6 3 k p , is taken as evidence of the existence of the undecomposed salt in solution. (3) Solutions of sodium sulphite do not show selective absorption. (4) Freshly prepared solutions of sodium bisulphite show end absorption only, but on standing in the light, if oxygen is present, photo-oxidation of the ion, HS08-,occurswith the simultaneous development of the band at X = z76pp characteristic of dissolved sulphur dioxide. The magnitude of the absorption coefficient continues to increase for about thirty days when a constant maximum is attained. ( 5 ) The presence of cane sugar in solutions of sodium bisulphite has been shown to retard the process of oxidation. ( 6 ) Light has been found to be an essential factor in the oxidation of bisulphite solutions. (7) Baly’s relationship connecting ultraviolet frequencies with a fundamental frequency in the infra-red has been found to be applicable t o the bands at X = 263pp and X = 2 7 6 1 ~characteristic ~ of K2S2Oj and S02.HzO respectively. Hzllside Laboratory, Stamford, Coian.
