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The behaviour of sulfur dioxide, oxygen, sulfuric acid and water in an electrolytic cell Assaly, T.C. 1945

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THE BEHAVIOUR OF SULFUR DIOXIDE OXYGEN, SULFURIC AO ID AND WATER IN AN ELECTROLYTIC CELL. by T.C. Assaly, B.A.Sc. A Thesis submitted i n P a r t i a l Fulfillment of the Requirements for the Degree of Master of Applied Science i n the Department of Chemistry. The University of B r i t i s h Columbia A p r i l , 1945 . I wish to acknowledge the helpful suggestions and assistance of Dr. W.3T» Seyer, under whose direction the present investigation was carried out. TABLE OF CONTENTS Page INTRODUCTION . 1 HISTORICAL •»••••••••»•».•.».«.•««»«...*...««.••»».»• 2 THEORETICAL CALCULATIONS..•»<.••«..•««.«<.»>«.«.<<•«•«* 3 Ionic Mechanism of C e l l Reactions. 3 Thermodynamic Calculations 3 Variation of E.M.F. with Concentration.. 4 Variation of E.M.F. with Pressure. 6 DESIGN AND CONSTRUCTION OF- THE CELL.................... 8 EXPERIMENTAL PROCEDURES 10 Variation of E.M.F. with concentration showing the existence of Hydrates 10 Rate of Approach to the Limiting Value of the .JE • J? * « • • • • • • • •X.'y Effect of Rate of Introduction of.Sulphur Dioxide and Oxygen on the C e l l Potential 21 Variation of E.M.F. with Temperature. ... 22 C OJLTCIJTJSION • • • • • • • • » • • • • « • • « • • • • • » # • • « • « • • • • • « • • • • » • • » • REFERENCES . 23 BIBLIOGRAPHY . 2 6 LIST OF ILLUSTRATIONS, Figure, ' -•• Page, 1, Theoretical E.M.F. of C e l l 7 2, Sulfur Dioxide, Oxygen, Sulfuric Acid C e l l . . . . . 9 3 , Measured E.M.F. of C e l l vs.Concentration, 13 4* Freezing Point Chart of Sulphuric Acid 15 Hydrates, 6. Rate of Change of E.M.F. vs. Time,. 20 7 , E.M.F. vs. Temperature 23 INTRODUCTION The u t i l i z a t i o n of the free, energy of sulfur dioxide i n the form of e l e c t r i c a l energy for the production of sulfu r i c acid i s of great i n d u s t r i a l importance i n the metallurgical and chemical industries where large quantit-A ies of waste sulfur dioxide are produced. Hitherto s u l f u r i c acid has been manufactured either by the Chamber process or Contact process. From a consideration of the heat of reaction A H and the free ensrgy change Aft of the reaction; S 0 2 + 3-/202 + H 2 ° ° H 2 S 0 4 ^ H 2 3°C. i = -52,100cal^G 2 5o 0=-49,100cal. the p o s s i b i l i t y of producing s u l f u r i c acid as well as e l e c t r i c a l energy i n an e l e c t r o l y t i c c e l l i s readily visualized.. In the following experiments i t was proposed to determine the actual electromotive force of the c e l l : Pt, S0 2; H 2S0 4; 0 2 , Pt at various concentrations of su l f u r i c acid and at different temperatures. Hydrates of su l f u r i c acid are known to exist. Their formation i s associated with a large entropy change. In an e l e c t r o l y t i c c e l l the formation of a hydrate would consequently be associated with a correspondingly large potential change. Further, the potential should vary with the nature of the hydrate formed. Hence, i t should be possible to detect the presence and nature of the different sulfur i c acid hydrates by means of the aforementioned c e l l . HISTORICAL From the l i t e r a t u r e i t appears that no previous work has been done d i r e c t l y on th i s problem. Experiments have been conducted by Messrs. 2 ' .. M. de Kay Thompson and U.J. Thompson and by Messrs. M. de Kay Thompson and A.P. Sullivan on the depolarization effects of sulfur dioxide i n an e l e c t r o l y t i c c e l l . They showed that i t was possible to decrease the polarization at the anode and thus decrease the equilibrium voltage of the c e l l by the addition of sulfur dioxide. I t appeared from th e i r experiments that the polarization was due to the accumulation of free oxygen at the anode. On addition of sulfur dioxide to the anolyte a reaction occurred between thi s gas and the free oxygen reducing the polarization effects. Since such a reaction i s accompanied by a substantial decrease i n free energy i t was evident that a c e l l of the type Pt, S0 2; H2S04| 0 2, Pt. should have a definite oxidation-reduction potential. - 3 -THEORETICAL CALCULATIONS Probable Ionic Mechanism of C e l l Reactions. Anode Reactions; The sulfur dioxide f i r s t dissolves i n the sulf u r i c acid electrolyte and then adsorbs on the surface of the electrode? S0 2 (gas) = SOg (solution) The sulfur dioxide then combines with water as i n equation (2) S0 2 + H20 = SOj + 2H+ + 2 (*•) (2) Cathode Reactions: Similarly, oxygen i s dissolved i n the electrolyte and then adsorbed on the surface of the electrode. '&2 ( s a s ) = °2 (solution) (3) Oxygen combines with water as i n equation ( 4 ) . 1/2 0 2 + H20 + 2 C W = 2 OH" (4} Overall C e l l Reaction S0 2 + 0 2 + m H20 = H2S04 (N=X) (3) Thermodynamic Calculations; Here the following nomenclature has been used. AG- = free energy change, calories Z^H = heat of reaction, calories /\ P = pressure change, atmospheres. ZvV = volume change, cubic centimeters, n = valence f = Earaday = 9 6 , 3 0 0 T = temperature,- degrees Centigrade, t = time, hours. M - mole frac t i o n N = normality. X = percent acid J = conversion factor: 1 c a l . = 4 ;183 joules 1 joule = 9 .87 cc.atmosphere Variation i n E.M.F. with Concentration of Electrolyte. The number of accurate measurements from which the free energy of d i l u t i o n of su l f u r i c acid may be c a l c u l -ated i s limited. Tables are available on the free energy of d i l u t i o n of the acid only for dilute solutions. Insufficient data.is known from which to calculate accurate l y the a c t i v i t y of the ions and hence the free energy of di l u t i o n of concentrated solutions. Bronsted^ studied the free energy of d i l u t i o n over a rjather wide range of concen-trations of sul f u r i c acid but at temperatures ranging only up to 9°C. Earned and Stu r g i s p and Lewis and Randall also studied the free energy of this acid, but over a small rang of concentrations. A satisfactory table of the free energy of di l u t i o n at 23°C for concentrations up to . 2 mol frac t i o n s u l f u r i c acid has been worked out by Randall and Cushman'. Their results which are the free energies £± G3. of the reaction: 2 H + + S 0 4 " + mHgO = H2S04(M=X) (6) are given i n the t h i r d column of Table I , and the corresponding values of mole fr a c t i o n M and percent acid X i n the f i r s t and second columns respectively. TABLE I M Z^Gl(cal) ^G2(k cal) E(volts) 1.00 100,00 0 -49.1 1 . 0 6 5 . 2 0 3 7 . 7 0 '•• 3 ' 2 7 0 - 4 3 . 8 . 9 5 0 . 1 3 44 . 9 0 2913 - 4 6.2 1 . 0 0 0 .10 3 7 * 7 0 1643 - 4 7 . 3 1 . 0 3 0 .08 3 1 . 6 0 7 0 2 -48.4 1.048 . 0 6 5 2 7 . 3 3 - 8 3 -49.2 1 . 0 6 7 . 0 3 2 2 . 2 8 - 8 6 5 - 3 0 . 0 1.084 . 0 3 14.42 -2048 -31.1 1 . 1 0 8 -,02 10.00 - 2 7 3 3 -31.8 1.122 .01 3.21 - 3 7 0 2 - 3 2 . 8 1.143 .002 1.08 -3613 - 5 4 . 7 1 . 1 8 5 . 0 0 0 9 .487 - 6 3 4 7 - 3 5 . 6 1.204 . 0 0 0 0 9 .049 -9382 - 5 8 , 5 ' 1 . 2 6 8 The c e l l reaction can be divided into two parts as such: s o 2 + 1/2 0 2 + H2G = H 2 S O 4 (7) 2 H + + S 0 4 ~ + mH20 = H 2S04(N=X) (8) s o 2 + 1/2 0 2 + mHpQ-= H 2S04(N=X) ( 9 ) Addition of equations (7) and ( 8 ) gfeves the overall c e l l reaction. Hence , addition of the free energy of reaction of equations (7) and ( 8 ) w i l l give the overall free energy change within the c e l l . The value of A G for equation (7) i s obtained from the difference of the f r e e energy of form-ation of su l f u r i c acid and the sum of the free energies of formation of-sulfur dioxide, water and oxygen. The free o energy of formation of HgSO^ - 1 7 6 , 0 0 0 calories , of SOg = -71^740 c a l o r i e s 9 , of H20 = - 5 6 , 6 9 0 c a l o r i e s 1 ? of 0 2 = / 1 0 calories, giving for "equation (7) /^ G = -49,100 calories. In equation (8) z^ G i s simply the free energy of d i l u t i o n ofi s u l f u r i c acid. The calculated values of is. G £f or the overall c e l l reaction (equation 3) are given i n the fourth column of Table I. By simple substitution i n to the formula: E = ( ^  Q). (J) (10) •(n) (f) the theoretical values of the emf that would be obtained i n a completely reversible c e l l are obtained. The values are given i n column 3 of the abpve table. In Fig. I the values of the emf given i n Table I are plotted as ordinates against concentration of su l f u r i c acid as abscissa. A hypothetical curve was drawn from . 2 M to 1 M sul f u r i c acid, the range of concentration where information on the free energy of d i l u t i o n was unobtainable. Variation of E.M.F. .with Pressure S0 2(g) + 1/2 0 2(g) + H20(.£) = H 2S0 4 (Z) (11) Equation (11) indicates a volume change from 1 1/2 moles of gas to zero moles of gas, - 8 7 = 1 .3 volumes One mole of gas at N.T.P. = 22,400 cc. ' One mole of gas at 20°G and one atm, = 24,050 vv. V = 1 .3 x 24 ,030 = 3 6 , 0 7 5 cc. 1 joule = 9 .87 co. atm. = 9.87 Apply the formula: _ £ E = A 7 1 1 (12) A P n f J the value ^V3 = 3 6 . 0 7 5 = .01892 12)196,500)19 .87) volts/atm. i s obtained, Miere A E i s the increase i n voltage per z^P unit increase i n pressure. DESIGN AND CONSTRUCTION OF THE CELL. A diagrammatic sketch of the c e l l i s given i n F i g . I I . I t consisted of two stoppered glass vessels, anode and cathode compartments, each containing as electro-l y t e , s u l f u r i c acid of the same concentration. Through the four hole's of each stopper was inserted a gas i n l e t , a gas outlet, a lead for the electrode and an arm of the sulf u r i c acid bridge. At the end of the sulfur dioxide and oxygen i n l e t s were di s t r i b u t o r s . This gave the entering gas a larger surface area, enhancing the saturation process. E l e c t r i c a l contact was made by two mercury f i l l e d glass leads. One end of each lead was sealed to Figure 1 H2.SO.4 SULFUR DIOXIDE, OXYGE1", SULFURIC ACID CELL FIGURE 2 / / the stem of a c y l i n d r i c a l platinized platinum gauze elect-rode, 2 cm. diameter and > cm. high. The potential was read from a potentiometer. The s u l f u r i c acid bridge consisted of a side arm of c a p i l l a r y tubing from each half c e l l leading into a t h i r d vessel also containing a solution of s u l f u r i c acid. The.whole c e l l was set i n a constant temperature bath provided with a constant l e v e l side arm to take care of evaporation. The temperature could be controlled to - .02°C. of the desired value. EXPERIMENTAL PROCEDURES. Variations of the E.M.F. with Concentration of Sul-f u r i c Acid Showing the Existence of Hydrates. The set up of the c e l l for each run was as shown i n Fig.I. The platinum electrodes were coated with a layer of platinum black deposited e l e c t r o l y t i c a l l y from a three percent solution of chloroplatinic acid. A current density great enough to produce a moderate evolution of gas was used and the direction of the current reversed every half minute for approximately one hour or u n t i l a fine velvety coating of platinum black appeared on the 12 surface of the electrodes. The sulfur disxide and oxygen were obtained from pressure tanks. The desired concentration of s u l f u r i c acid for each run was obtained by d i l u t i n g C P . 95% acid with d i s t i l l e d water. The exact strength of the acid i n each case was determined by t i t r a t i n g against.a standard NaOH -solution (approx. 6 N). The sodium hydroxide was standardized against CP. oxalic acid. Phenolphthalein was used as indicator. The two half c e l l s were f i l l e d about three quarters f u l l with s u l f u r i c acid, the side arms of which dipped into a single intermediate vessel also f i l l e d with sul f u r i c a d d to the same l e v e l and of the same strength. The electrodes were only partly immersed i n the electrolyte about one-quarter remaining above the surface. A l l experiments were carried on at 2J?°C and atmospheric pressure. After each run the electrodes were removed and degassed by e l e c t r o l i z i n g them at approximately 8 amperes i n a concentrated solution of sodium hydroxide for one hour, the direction of the current being reversed every ten minutes for the f i r s t f i f t y minutes, then every minute for the l a s t ten minutes. They were f i n a l l y washed i n boili n g d i s t i l l e d water for f i f t e e n minutes. The measurements obtained are given i n Table I I . The f i r s t and second columns give the normality N and the corresponding percent X of the su l f u r i c acid and the t h i r d columns the measured ...enf,. ©f the c e l l . Each row represents a new run. The values of the eemf t given i n the Table are \ the values obtained after the sulfur dioxide had bubbled through the c e l l for approximately 24 hours and the potential had become steady. TABLE I I . s. X(fo) E (volts) «93 3*60 «7339 2.70 1 2 . 2 2 . 7 3 0 x 5.11 21 .70 . 7 2 4 0 69 26 26.82 • 6774 8 . 5 0 3 3 . 3 0 . 6 7 3 ! 1 0 . 7 0 40 . 10 . 671Q 12.40 4 4 . 9 7 . 6 6 8 4 14.14 4 9 . 6 5 «666g 16.71 56.04 . 664Q 17 .30 5 7 . 4 4 - . 660Q 1 8 . 0 0 59.13 . 6 3 0 4 18.47 • 6 0 . 1 7 , . 6 0 5 5 2 0 . 7 1 63.75 . 5 9 0 ^ 24 .00 6 9 . 4 6 .5897 2 5 . 3 0 7 4 . 3 5 . 5 7 4 1 2 6 . 5 0 7 6 . 6 8 • 5 7 2 0 2 7 . 5 0 78.49 »572 G 2 8 . 9 0 81.14 . 5 7 4 x 2 9 . 8 9 8 3 . 0 2 . 547 5 3 0 . 6 0 84.40 .5262 3 1 . 7 0 8 6 . 4 5 • 5 2 2 0 3 6 . 8 0 9 7 . 7 0 . 5 2 2 5 -13-Hgure 3 In F i g . I l l the values of the emf given i n Table I I are plotted as ordinate. The abscissa i s the normality and the corresponding percent acid. The exper-imental results f a l l on a stepwise curve. Each of the four steps corresponds on the abscissa to a hydrate of sul f u r i c acid. The f i r s t step at J0 .7 N(84.48fo) corresponds to HgSO^.HgO; the second at 24.6 N (73.13%) to. H2SG4.2H20« the t h i r d at 17.3 N (57.36f.) to H2SO4.4H20j the fourth at 5.04N (21.40°/.) to HgSO^^O HgO. I t may also, be noted that the slope of each plateau or f l a t portion of the. curve from 36.8 N to 0 N acid increases s l i g h t l y . The formation of a hydrate i s associated with a comparatively large free energy change and entropy change. From the formula £>& - -nFE i t necessarily follows that the formation of a hydrate in. an e l e c t r o l y t i c c e l l would be associated with a large change i n potential. Conversely, the presence of the stepwise increases i n potential presumes a large increase i n free energy or decrease i n the irrevers-i b i l i t y (entropy) and hence, substantially indicates the existence of the four hydrates H2S04.H20, H2S04«2B:20, . HgSO^^HgO, H2S04.20 HgO at 25°C and atmospheric pressure. . Other experimenters have shown evidence f o r the existence of three of these hydrates by other means. Pickering 1^ obtained the hydrate H2SQ4.$H20 by fusion methods. 14 Giron found the hydrate H2S04.2H20 by s o l i d i f y i n g mixtures -15 Figure 4 containing from 68 to 70% of the acid. Donk obtained cr y s t a l l i n e s u l f u r i c acid monohydrate H 2 S O 4 H 2 O by freezing 16 methods. Bourgoin indicated the presence of H 2 S G 4 . 2 H 2 O from observations on electrolysing dilute v i t r i o l . The graph given i n Abegg*s Handbuch der Anorganischen Chemie, IV, 472 (1927) and reproduced i n Fig.IY covers a l l of the hydrates found by the freezing point method. B, D and G-are the cryohydrate points and C, E and H the melting points of the hydrates H 2 S O 4 . 4 H 2 O , H 2 S O 4 . 2 H 2 O and H 2 S O 4 H 2 O respectively. From B to C, D to E and G- to H these respective hydrates s o l i d i f y out. The existence of H 2 S Q 4 . 2 O •..'HjjQO i s not indicated i n the graph. Water supposedly s o l i d i f i e d out from A to B. Probably on much closer exam-ination another cryohydrate point and melting point between A and B would have been detected representing H 2 S O 4 - 2 O H 2 O . Since i n such a hydrate i t would be d i f f i c u l t to detect the H 2 S Q 4 i n the combined state it'may have been overlooked. The e l e c t r o l y t i c method of the writer gives evidence that these hydrates are also stable i n solution up to at least 25°G as-well as at the i r freezing points. I t may be of interest to note the large i r r e v e r s i b i l i t y or unavailable work content of this c e l l by comparing the reversible emf of the c e l l as calculated from free energies (Table I, Fig,I) with the actual measured values obtained from t h i s c e l l (Table I I , F i g , I I I ) , In addition, the calculated emf values do not show t h i s step-wise• phenomenon'.• This may be explained by the fact that the calculated values of the free energy are actually overall values i n which the free energy of formation of the hydrate i s only one factor. Rate of Approach to the Limiting Yalue of the E.M.F. The l i m i t i n g values of the emf given i n Table I I are the values obtained after the sulfur dioxide had bubbled through the c e l l f o r several hours and the values had become steady. That i s to say, a steady state was not approached immediately but required 24 to JO hours. Measurements were made of emf with time for several concentrations of su l f u r i c acid. The measurements for three concentrations are given i n Table.III. The f i r s t column gives the time, the second, third and fourth give the emf measured with 6.26 N, 20.7 N and 31.7 N acid respectively and the f i f t h the rate of change of emf dE t at for 6.26 N acid. TABLE £11 t E ( 6 . 2 6 N ) E (20.7N) E(31.7N hours mins. volts volts volts 0 0 0 0 15 .217 .150 .156 25 . 3 6 0 " v .255 ,240 30 34 .440 40 45 . 4 6 5 .352 .315 1 00 1 15 .525 .410 .359 1 30 1 45 .558 , 4 6 5 . 3 8 0 2 15 2 45 . 5 9 6 .489 .424 3 15 3 45 .618 .508 .453 4 45 5 45 . 6 4 4 .547 , 4 8 6 7 45 9 45 .655 .569 .499 13 15 20 45 . 6 7 5 21 45 22 34 46 45 45 45 .6768 . 6 7 7 4 . 5 9 0 5 . 5 9 0 ^ , 5 2 1 6 .522@ at volts/hour. .872 .868 >858 .533 ,191 .104 . 0 6 6 .038 . 0 2 2 .013 . 0 0 3 • 001 • 0009 ,00002 In Figures V and VI the values of the emf and the rate of change of emf respectively are plotted as ordinates against time as abscissa* I t was deduced from a comparison of the results that the rate of change of emf with time varied only s l i g h t l y over large ranges of concentration as long as the rate of introduction of sulfur dioxide into the c e l l was consistent for 'the different concentrations. From a close examination of the curves i t was assumed that the absorbitivity of the electrodes was very high at f i r s t , approaching a straight l i n e relationship but decreased rapidly as the electrodes approached saturation.. Figures V and VI give only a qualitative picture since the rate of introduction of sulfur dioxide varies the olope of the straight l i n e portion of the curves although leaving the general shape the same. Effect of the Rate of Sulfur Dioxide and Oxygen into the C e l l on the Limiting E.M.F. Although the rate at which equilibrium i s attained i s affected somewhat by the rate of sulfur dioxide into the c e l l the absolute value of the emf after a l i m i t i n g condition i s reachedj i s almost independent of the rate. For example, on increasing the rate of sulfur dioxide into the c e l l from 120 bubbles per minute to over 240 bubbles per minute, the l i m i t i n g value of the emf (after several hours) changed from *7240 to .7252.. This value returned to .724^ on decreasing the rate back to 120 bubbles per minute. Decreasing the rate s t i l l further to 20 bubbles per minute, cut the emf to »7220 which returned to «724Q on increasing the rate back to 120 bubbles per minute. Oxygen acted d i f f e r e n t l y . Increasing the rate from 60 bubbles per minute to 4 times that value, caused the l i m i t i n g value of the emf to drop from «724Q to , 7 2 J i while increasing the rate to well over 8 times caused the emf to drop down to , 7 2 0 i . On decreasing the rate back to 60 bubbles per minute, the emf went up to .72?£ while shutting off the 0 2 i n l e t caused no perceptible change. Hence, i t can.be assumed that at low rates the emf i s independent of the rate while at greater rates, oxygen has a polarizing effect. Variation of E.M.F. with Temperature. Measurements were made using j?»27 N acid at temperatures ranging from 15°C. to 45°C. The measurements are given i n Table IV. TABLE IV, T (°C. ) E (volts) 15 imQ 25 , 7 2 3 0 35 , 6 9 5 4 45 . 6 2 9 7 In Fig.VII the values of the emf from Table IV are plotted as ordinates against temperature as abscissa* The results f a l l on a smooth curve. An asymptotic curve approaching Es,755> seems to result on extrapolation,to lower temperatures. - 2 3 -Figure 7 CONCLUSION I t was shown from the experiments conducted throughout this research that the u t i l i z a t i o n of the free energy of sulfur dioxide i n the form of e l e c t r i c a l energy has considerable p o s s i b i l i t i e s . P l a t i n i z e d platinum electrodes alone were used throughout the investigation. Although they gave good results i t should be possible to decrease the i r r e v e r s i b i l i t y and approach the theoretical emf as calculated from free energies more closely with more reversible electrodes 1^. Certain p o s s i b i l i t i e s are the use of platinum black on an asbestos c a r r i e r , iridium black on platinum or lead elect-rodes. A further study into the stepwise phenomenon of F i g . I l l may unveil some conclusive facts regarding the hydrate formation already proposed and lay the basis for a theoretical explanation of this phenomenon as well as a more thorough knowledge of s u l f u r i c acid i t s e l f . A suggestion would be to study each half c e l l separately against a standard c e l l , such as the hydrogen electrode. - 25 REFERENCES. . 1. • E. Wendell Hewson, Atmospheric Pollution by Heavy Industry, J. Ind. Eng. Chem., Vol.36, No.3, 195 2. Chem. & Met., 15, 677 (1916) 3.. Chem. and Met., 18, 1?8 (1918). 4. Z. Physik. Chem., 68, .693 (1910) 5* «T.A. Oh. S., 47, 94-5 (1925) 6. J . A . Ch.S., 36, 804, (1914-).. 7. J . A . Ch.S., 40, 393, (1918). 8 . Lewis and Randall, Thermodynamics, 554- (1923). 9» Perry, Handbook of Chemical Engineering, Ed.IIj 563 (194-1) 10. Perry, Handbook of Chemical Engineering, Ed.II, 553 (1941). 11. Perry, Handbook of Chemical Engineering, Ed.II, 653 (1941). 12. Findlay, P r a c t i c a l Physical Chemistry, 152 (1923). 13. Chem. News-, 60 (68). 14. B u l l . Soc. Chim., 1913, 13, 1049. 15. Chem. Weekblad, 10, 956, Abst. Am. Chem. Soc. ,1914,2,1926. 16. B u l l . Soc. Chim., ( 2 ) , 12, 433. 17. Shiels, J.P.C., 33, H 6 7 . - 26 BIBLIOGRAPHY. 1. Abeggis Handbucli der Anorganisohen Chemie, IV, (1927). 2. Adam, N.K., Physics and Chemistry of Surfaces. 2. Berkman, S., Morrell, J.C., Egloff, G., Catalysis (1940). 4. Boriskov, S.K. and Ruderman, E.E., J. Phys. Chem. (U.8.S.R.), 14, 161-170. 5. Creighton and Koehler, Electrochemistry, Ed.II, Vol.11 (Applications). 6* Getman and Daniels, Outlines of Physical Chemistry,Ed.VII. 7. Glasstone, Physical Chemistry. 8. Gregg, S. J"., Adsorption of Gases by Solids. 9. Hodgeman, CD., Handbook of Chemistry and Physics. 10. Latimer, Oxidation Potentials* 11* Lewis and Randall, Thermodynamics and the Eree Energy of Chemical Substances. 12, McBain, S.W., Sorption of Gases. 13* Mellor, J.W., A Comprehensive Treatise on Inorganic and Theoretical Chemistry, Vols, I and X. 14. Perry, Handbook of Chemical Engineering, Ed.II. 15. V i n a l , Storage Batteries, Ed.III. 16. Wade, Secondary Batteries. 17. WyM, W., Sulfuric Acid and Sulfur Dioxide. 

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