September 1952
INDUSTRIAL AND ENGINEERING CHEMISTRY
2207
KT TABLE IV.
H.T.U. EVALUATION
1 Run
5
Equilibrium Temp., Temp. (" IC.) Enrichment O C. 10' 24 3.37 1.059 22 3.39 1.041 21 3.40 1.116 22 3.39 1.092 20 3.41 1.136 15 3.47 1.077 1.102 28 3.32 30 (1.164)a 3.30 25 3.36 1.153 17 3.45 1.050 17 3.45 1.056 (1.193)' 50 3.10 51 3.09 1.285, 100 2.68 (4.70) 96 2.72 1.815 86 2.79 1.250 86 2.79 1.364 122 2.53 2.575 Calculated results.
a
N
1.0392 1.0398 1.0400 1.0398 1.0402 1.0417 1.0381 1.0376 1.0390 1.0412 1.0412 1.0332 1.0331 1 .0258 1.0264 1.0276 1.0276 1.0236
1.49 1.03 2.79 2.25 3.22 1.82 2.60 (4.11Ia 3.72 1.21 1.35 (5.39)a 7.70, (60.8) 22.9 8.20 11.4 40.5
H.T.U., Ft. 47.0 68.0 25.1 31.1 21.7 38.4 26.9 (17.0)" 18.8 57.8 51.9 (13.O)a 9.08 (1.15)* 3.06 2.68 1.93 0.543
Press., Lb./Sq. In. 93 88 215 207 199 170 148 155 148 155 159 156 153 153 153 159 156 152
T
perature, T total number of transfer units = gas constant, caloriee per mole per ' K. = temperature, C. and
2
= mole fraction of oxygen
N
NHa G./ldO G. 0 0
R
0 0
0 1.09 3.06 4.42 1.21
g CY
0 1.70
0 1.63 4.06 0 2.13 3.97 2.98
= reaction rate at tem-
' K.
18 in t h e liquid mole fraction of oxygen 18 in the vapor = separation factor or relative volatility =:
BIBLIOGRAPHY
W. W., and White, R. R., Chem. Eng. Progress. 44. 553 (1948).
(1) Akers,
R. R., Zbid.,
46, 563
(1950).
About 99% of the upsets were a result of the ammonia. With no ammonia it was not necessary t o watch the operation. For larger units it will be desirable t o devise a more satisfactory ammonia system.
(5)
ber 23-2, Berlin, Verlag Chemie, 1936. Huffman, J. R., and Urey, H. C., IND.ENG. CHEM.,29, 531
(6)
Mills, G. A,, and Urey, H. C., J . Am. Chew. SOC.,62, 1010
(1937).
ACKNOWLED~MENT
This work was sponsored by the American Cancer Society upon recommendation of the Committee on Growth of the National Research Council. NOMENCLATURE
A* H.T.U.
(3) Boyd, W. T.. Ph.D. dissertation in chemical engineering, University of Michigan, Ann Arbor, Mich., 1951. (4) Gmelin, "Handbuch der anorganischen Chemie," System Num-
= activation energy, calories per mole -- height of column equivalent to one transfer unit, feet
(1 940). (7)Nier, A. O., priv&e communication (1949). (8) Urey, H. C., private communication (1947). (9) Urey, H. C., and Grieff, L. J., J. Am. C h m . SOP..57, 321 (1935). (IO) Reid, A., and Urey, H. C., J. Chem. Phys., 11, 403 (1943). (11) Weber, L. A., Wahl, M. H., and Urey, H. C., ZW., 3, 129 (1935).
RECEIVED for review November 30,
1951.
ACCEPTEDApril 3, 1952.
Ozone by Electrolysis of Sulfuric +
Acid
JUNIOR D . SEADER' AND CHARLES W. TOBIAS University of California, Berkeley, Calif.
I
N 1840 Schonbein (14) discovered t h a t with oxygen, ozone
is also liberated when electrolyzing aqueous acid and salt solutions. Investigations by Soret (16), McLeod ( I 1 ), Targetti (16), Grafenberg (6), and Kremann (7), led t o the conclusion that sulfuric acid of 1.075 t o 1.100 specific gravity gives the highest yields (up t o 18% by weight), using platinum anodes and keeping the temperature of the cell close to 0" C. Fischer and co-workers (3, 4)recognized the importance of low anode surface temperature. The platinum anode surface was believed t o catalyze the decomposition of primarily formed ozone. The highest ozone yields were obtained by using a n inside cooled platinum tube anode of triangular cross section, coating i t with Jena glass and grinding away the glass over one of the edges so as t o expose a horizontal platinum strip, 0.1 mm. wide and 11.5 mm. long. This anode gave a very short contact time for gas bubbles and hence was believed to have reduced the decomposition rate of ozone. The best yield (28%) resulted when the anode was cooled inside by a circulating salt solution t o -14' C.; the current density was 87 A. per square cm., and the electrolyte was sulfuric acid, specific gravity, 1.075. Fischer also studied the ozone process using various other electrolytes. None of the salts, acids, and 1
Present address, University of Wiseonsin, Madison,Wis.
bases, other than sulfuric acid, gave high ozone yields, On the basis of t h e available free energy data, Fischer calculated the potential of the OH-/Os couple t o be -1.69 volts. Wartenberg and Archibald (18) investigated the effect of superimposed alternating current on ozone yield. They claimed to have attained very high (up t o 35% by weight) concentrations of ozone in oxygen, but the energy yield in these experiments was exceptionally poor. Their results were not substantiated by other workers (1, l a ) . Briner and coworkers ( I , 2) were the first to investigate the anodic liberation of ozone a t temperatures around -50" C. These low temperatures required the use of eutectic sulfuric acid-water compositions. The highest ozone yields reported by Briner were below 8.5% by weight. However, by using a 4070 solytion of perchloric acid instead of sulfuric acid as electrolyte, yields up t o 16% were observed. Briner cooled the electrolyte instead of the anode. The experimental procedures, electrode potential data, and theoretical treatment extended by this group invite severe criticism and are partly unacceptable. Putnam and coworkers (IS)undertook a systematic study of the effects of temperature, current density, acid concentration, and pressure on the ozone yield, using perchloric acid as electrolyte. The yield waa found to increase with current density and
Vol. 44, No. 9
INDUSTRIAL AND ENGINEERING CHEMISTRY
2208
concentration. Optimal results were obtained by using 40% by weight acid, 20/, magnesium perchlorate, and 0.2 A. per square cm. anodic current density a t - 53 C. and by reducing the pressure t o 0.1 atm. over the electrolyte. Under these conditions the current yield was 25.9y0, accompanied by an energy yield of 14.5 grams of ozone per kw.-hr. Putnam did not use inside cooled anodes; only the electrolyte was cooled. The electrolysis of sulfuric acid solutions was attempted, but the yields obtained
that the lower the acid concentration, the higher the ozone yield at any given temperature. Since high current density favors persulfate formation, whereas a t low current densities the overvoltage of oxygen is not large enough to permit the liberation of ozone, moderately high current densities should be expected tJo result in the highest ozone yields.
TABLE I.
FORM.4TION O F PERSULF.4TE (Temperature, O o C.)
Current Current Density, Efficiency A,/&. Crn. Specific Gravity of OZONE COLLECTION
SYSTEM
HzSOc Solution 1.15 1.20 1.25
0.05
0.5
1.0
...
7.0 20.9
4.4 29.3
43.5
4HODE COOLING SYSTEM
4 i
I RaRYjERbNT STMIAGE TAW
2
CENTRBUML PUMP
3 HEAT EXCHPNWR 4
ELECTROLYTIC CELL
5
ABSORBERS
6 SAFETY BOTTLE
Figure 1. Experimental Apparatus fol' Ozone Production
did not exceed 5.8%. [During the preparation of this manuscript Lash et al. (8) published the result of experiments conducted with inside cooled anodes using perchloric acid as electrolyte. Energy yields up to 24 grams per Iw.-hr. and concentrations up t o 58 weight % ozone were reported.] The status of anodic ozone liberation, a t the time the present work was started, could thus be summarized in the following facts, which were agreed on by the various investigators: 1. Of the many acids, salts, and bases studied only sulfuric acid and perchloric acid seem t o give reasonable ozone yields. 2 . The only anode material found suitable for this process was platinum (or its alloy). 3. The lower the anode temperature, the higher t h e ozone yield.
Findings of various investigators are in disagreement, if not contradiction, about the effects of current density, electrolyte composition, and electrode and cell geometry. To date no theoretical explanation has been offered about the mechanism of ozone formation. Theoretical considerations are: The standard potential of the water-oxygen couple 2H20 = Oz
+ 4H+ + 4e-, E" = -1.229
The net anodic reaction for the oxidation of water t o ozone 3Hz0 = Os f 6H+
+ 6e-, E" = -1.509
In a sulfuric acid solution, in addition t o these anodic- reactions, sulfate ions may be polymerized 2504'
=
&Os"
+ 2e-, E" =
-2.05 ( 9 )
From these values it is apparent t h a t on an electrode surface, where oxygen discharge is accompanied by a large overvoltagee.g. , on smooth platinum-the liberation of ozone becomes thermodynamically possible. It also follows that ozone may be formed preferentially t o persulfate. If however, the liberation of ozone itself requires such a large overvoltage t h a t the potential required for the persulfate formation is exceeded, oxygen, ozone, and persulfate formation could occur simultaneously. In a comprehensive review of per- compounds, Machu ( I O ) presented data showing t h a t persulfate formation becomes appreciable only a t relatively high acid concentrations and high current densities (Table I). It is t o be expected, therefore,
In addition t o simultaneous formation of oxygen, ozone, and persulfate, the catalytic decomposition of ozone on the electrode surface might also contribute t o lowering the ozone yield. The uncatalyzed decomposition of ozone t o oxygen in the pure gas phase was found to be a second-order reaction. According to Wulf and Tolman ( 1 9 ) the second-order rate equation is
_ -'Os dt
30,000 -_ _ =
kz(Os)2; kz = 2.04 X 10%
RT
( a t 1 atm. pressurej
Appreciable decomposition of ozone t o oxygen a t subzero temperatures can therefore only be expected in the presence of suitable catalysts. The catalytic effect of platinum on the decomposition rate of ozone has not been investigated quantitativelr Such an effect would be expected t o increase sharply with temperature. EXPERIMENTAL
The apparatus used for investigating the effects of current density, acid concentration, and temperature is illustrated in Figure 1. The cell was constructed of six blocks of 1-inch thick Lucite (a methylmethacrylate polymer). The inside dimension? of the cell were 8 X 4 X 3 inches. The cell wall!, top and bottom, were provided with six '/*-inch slots. By inserting Lucite and/or diaphragm plates in these slots, various cell lengths, cathode, and diaphragin locations could be effected. The anode was introduced through a circular hole in one of the end blocks. It consisted of a circular platinum plate, 0.010 inch thick and 1.32 cm in diameter, m-elded to the end of a tantalum tube 2'/2 inches long and having a wall thickness of about 0.01 inch. This tube was supplied with a special fitting so that it could be connected into the refrigeration system, which v a s t o cool the anode internally. The cooling fluid was jetted against the inside surface of the platinum plate anode through a copper tube which projected inside the tantalum tube t o within l/g inch of the anode plate. Such a cooling fluid discharge Tc-ould produce relatively high heat transfer coefficients, thus the temperature drop between the anode surface and the cooling fluid would be small. A plate of either silver-free lead or platinum plate ( 2 1 / 2 X 3l/2 inches) served as the cathode. No provisions were made to cool the cathode. A porous ceramic plate (Ceraphragm D-2, General Ceramics and Steatite Corp.) or a rectangular piece of glass cloth (Fiberglas Cloth So. 181, Owens-Corning Fiberglas Corp.) cemented t o a Lucite frame served as the diaphragm. The gas collection system consisted of two gas absorption bottles nnd a safety bottle in series. On the exit side of the safety bottle reduced pressure was maintained t o facilitate the passage of gases through the absorption bottles. The pressure measured above the electrolyte in the cell was only l/z inch of water below atmospheric, The simple refrigeration system is illustrated in Figure 1. An equimolal mixture of chloroform and carbon tetrachloride was circulated by the small centrifugal pump. I n the Dewar bottle solid carbon dioxide was evaporated in a mixture of similar composition to the above. The electrical system consisted of two 6-volt storage batteries in series with a rheostat, an ampere meter, and the electrolytic cell. PROCEDURE
After arranging the diaphragm and cathode spacings the cell was filled to the desired level with sulfuric acid electrolyte. The
September 1952
INDUSTRIAL AND ENGINEERING CHEMISTRY
2209
refrigeration system for the anode cooling was then brought into operation, and the desired current was applied across the cell. Samples were taken only after attainment of steady state, characterized by constant potential and temperature readings (2 t o 3 hours after the start of electrolysis). The cell gases were bubbled through freshly prepared potassium iodide solution buffered with ammonium and aluminum chlorides (17). The
Figure 2.
Thermocouple Locations
iodine liberated during a measured period of time was titrated by standardized sodium thiosulfate solution. The amount of ozone absorbed in 300 t o 400 cc. of buffered 0.1 N potassium iodide solution was between 20 and 50 mg. Control experiments showed t h a t acidification prior t o titration with 0.1 N sodium thiosulfate was not necessary. The difference of yield obtained from acidified and nonacidified samples did not exceed relative 1%. The duration of experimental runs varied from 15 minutes t o 3 hours, as many as four runs being made consecutively under similar operating conditions. RESULTS
Tables I1 t o V and Figure 3 present data showing the effect of current density, electrolyte composition, and refrigerant temperature on the ozone yield. All runs were carried out with a n anolyte volume of 25 cc. and a catholyte volume of 50 cc. The electrolyte level was 2 mm. above the top of the circular anode plate. The distance between electrodes was 20 mm., the diaphragm being located 10 mm. from the anode. Included in the tables are the following quantities:
TI = temperature of the refrigerant at a point inch from inside of anode wall (thermocouple locations are illustrated in Figure 2) TB = electrolyte temperature inch from one side of anode T 8 = temperature in main body of anolyte 2'4 = temperature in cathode compartment iust in front of cathode i = current aensity a t anode Current efficiency = percentage of ozone obtained based on an equivalent weight of 8 Energy yield = grams of ozone per kw.hr.; energy requirement for refrigeration was not included DISCUSSION OF RESULTS
UNSTEADYSTATE OPERATION. The variation of cell potential with time during the approach t o steady state conditions at a constant current density is shown in Figure 4 for a typical run. The attainment of steady state conditions took approximately 3 hours.
I+&+&+--
1.5
C U R R E N T D E N S I T Y , AMP./SQ. CAM.
Figure 3.
Dependence of Current Efficiency on Current Density Sp. Gr. HzSO,
Curve
TI,"C . -63.5 -63.1 -61.7 -61.3
0=1 A=2
0 =3
x =a + =5
-
2.3
EFFECT OF CURRENT DENSITY. As shown by Figure 3, at -63" C. the four different electrolyte concentrations yielded a similar pattern of dependence of current efficiency on current density. The ozone yield rose sharply with increasing current density, went through a maximum, and then gradually decreased as the current density was further increased. This behavior may be the result of three factors: (1)The overvoltage of oxygen increases with current density, therefore a larger fraction of the
A N D CALCULATED DATAFOR ELECTROLYSIS OF SULFURIC TABLE 11. EXPERIMENTAL ACID AT VARIOUSCURRENTDENSITIES
(Sp. gr. of HzS04 1 2 0 0 ) Run
No 22 21 25 26 24 23 74
7. 2_
76 72 75
"C _ _ _ _ ~ _ _ _Tempeiature, ____TI Tz T8 1.0 4.0 1.6
10.0 11.4 14.0 16.0
- 3.6
30.5 -27.7 -25.2 -20.6 -21.0 -21.2
-
--1.8 1.0 -63.6 -64.0 -62.8 -63.7 -63.3
21.5
10.4 11.4 15.0 16.2 22.0 28.0 -20.6 -19.8 -15.0 -16.8 -14.0
TI 13.6 16.2 21.0 22.4 28.4 38.0 -10 4 - 7 8 - 3 8 - 4 6 - 5.6
Amp./ Sq. Cm 0.293 0.585 0.733 0.733 1.465 2.930 0.249 0.497 0.745 0.827 1.240
2,
E. Volts 4.17 4.82 4.67 4.73 5.47 7.20 8.50 8.90 8.90 9.56 9.95
Current Efficiency
%
0.78 1.03 1.16 1.11 1.25 1.27 23.6 30.3 32.4 32.1 27.5
Energy Yield, Grams/ Kw.-Hr 0.56 0.63 0.74 0.70 0.68 0.63 8.27 9.12 10 84 10 03 8.24
TABLE 111. EXPERIMENTAL AND CALCULATED DATAFOR ELECTROLYSIS OF SULFURIC ACID AT VARIOUSCURRENTDENSITIES (Sp. gr. of HzSOa, 1.2.50) Temperature,
Run No. 71 64 70 68 69
O
C.
____________ TI Tz T8 -64.0 -63.0 -63.6 -61.8 -63.2
-28.2 -21.0 -24.2 -11.0 -18.4
-23.8 -14.8 -19.2 4.6 -12.4
-
T4 -12.6 - 6.3 7.8 - 5.2 - 0.5
-
i, Amp./
Sq. Cm.
0.099 0.866 0.579 0.993 1.242
E, Volts 6.70 9.15 8.85 9.10 9.80
Current Efficiency.
70
19.04 26.80 27.30 23.30 22.10
Enerev Yiel;i, Grams/ Kw.-Hr. 8.47 8.73 9.18 7.63 6.72
2210
INDUSTRIAL AND ENGINEERING CHEMISTRY
Vol. 44, No. 9
with the electrolyte of 1.2000 specific gravity, increase in acid concentra,tion brought about a significant decrease (Sp. gr. of H*SOn, 1.285) in ozone yield. Experiments with acid Ener y Current Yieli, concentrations outside the 1.20 t o 1.35 Run Temperature, C. i, Amps./ E, Efficiency, Grams/ range proved t o be of no value becausc Tz T3 T4 Sq. Cm. Volts yo Xw.-Hr. No. Ti 52 -62.0 -19.0 -16.5 -5.4 0.219 9.50 17.67 5.54 the electrolyte partially froze out. [See 51 -62 0 -17.6 -15.0 -4.0 0.366 10.25 20.35 5.92 phase diagram by Gable et aE. ( 5 ) . ] -2.1 0.511 10.40 54 -61.4 -14.6 -12.8 22.20 6.36 55 -61.2 -12.6 -10.4 0 0.658 10.85 19.89 5.41 It is suspected that a t higher acid 58 -61.1 -13.1 -11.1 -2.4 0.731 10.25 18.60 5.41 concentrations persulfate was formed a t the expense of ozone formation. In TABLE lT.EXPERIMENTAL AND CALCULATED DATAFOR ELECTROLYRIS O F SULFURIC the presentM-orlc no attempt was ACID AT VARIOUSCURRENTDENSITIES to justify this assumption by analysis. (Sp. gr. of HzSOi, 1.380) ( I t would be difficult t o establish quanEnergy Current Yield titative proof because the persulfate can i,Amos./ E, E5ciency, Grams/ Run Temperature, C. undergo hydrolysis and is unstable n.ith No. Ti Ti T8 T4 8s. CM. Volts % Kw.-Hr. 20 1.0 10.6 12.0 14.2 0.219 3.90 1.08 0.79 respect to both anodic and cathodic: 19 0.9 12.0 13.5 16.0 0.585 4.85 0.72 0.45 reactions. ) 12 -62.0 -13.8 -10.0 - 2 6 0.146 7.80 10.21 3.90 13 -61.5 -20.2 -15.0 -8.0 0.146 8.20 10.17 3.69 EFFECT OF TEMPERaTURE. It is be14 -61.2 -20.6 -15.0 0.146 8.20 10.18 3.70 -- 88 .. 02 0.146 8.25 8.89 3.21 licvcd that only the temperature of the 15 -61.2 -20.6 -14.7 27 -62.8 - 9 . 4 -11.2 -10.0 0.183 8.75 11.70 3.98 anode surface is of importance. The 11 -61 2 -16 2 - 04 . 0 0.219 8.85 14.23 6.83 - 6.6 1 -61.6 -11.4 0 0 252 8.80 15.75 5.34 heat transfer between the anode and 2 -61.5 -12.0 -6.6 -0,4 0.252 9.10 15.40 5.05 -1.3 16.75 4.84 the electrolyte has a rather complex -11.3 -8.0 0.366 10.30 50 -61.2 48 -61.5 -14.5 -14.4 6.0 0.438 9.40 17.00 5.38 mechanism due t o individual bubble 47 -61.5 -15.3 -13.8 -6.2 0.511 10.00 17.37 5.18 - 7.0 - 1.3 0.658 39 -60.0 -10.4 11.30 16.50 4.34 formation. S o attempt v a s made t o 45 -61.2 - 5.2 - 4 . 6 5.0 0.658 11.00 15.93 4.32 10.74 2.48 measure the surface temperature since it -10.6 - 0.4 1.007 43 -60.6 -10.0 12.90 was believed that the temperature would not be uniform over the plate, but SHOWING EFFECT OF DIAPHRAGM o s CURRENT EFFICIENCY TABLE VI. DATA would be localized a t the points of bubble IN ELECTROLYSIS OF SCLFURIC ACID, SP. GR. 1.285 nucleation. Experiments conducted with a refrigCurrent Energy Yield, Run I Amps/ Efficiency, Grams/ erant temperature of about 0" C. gave ozone: Kw.-Hr. No. Diaphragm TI. C. dq. Cm. E,Volts % yields lower than 2%. This finding coincides 54 Ceramic -61.1 0.511 10.40 22.20 6.36 69 None -60.0 0.511 5.1 0.72 0.42 with those reported in the literature. 60 Glays EFFECT OF DIAPHRAGM. Without a diaphragm 0 511 9.6 24.25 7.53 cloth -61.9 61 None -59.0 0.511 5.06 0.94 0.55 separating the electrodes, the ozone yiel1 became vanishingly small ( < 1%). The role of the diaphragm in this process is probably rastricted to prevent the transportation of hydrogen bubbles by convection t o the anode. Several types of glass clot,hdiaphragms current might produce ozone a t higher current densities. (2) However, as the current density is being increased, the steady were used with equal success, the ozone yields obtained were state surface temperature of the anode increases also, followed identical within fractions of 1% t o yields obtained by using by a decrease in ozone yield. (3) At higher current densities ceramic diaphragms. If persulfate had been produced in signifia larger fraction of the current produces persulfate ions. cant quantities at the anode, the glass cloth diaphragm woultl EFFECTOF SULFURIC.4CID CONCENTRATION. At -63" C. have been quite inadequate t o prevent its cathodic reduction. refrigerant temperature the highest ozone yields were obtained ENERGY YIELD. No attempts were made t o improve the energy yield in any of the experimental runs. It is believed that by shortening the distance betw-een the electrodes and maintaining the temperature of the electrolyte around room terrperature, the energy yield could be improved by 30 t o 40%.
TABLE 11'. EXPERIXENTAL AND CALCULATED DATAFOR ELECTROLYSIS O F SULFURIC ACID AT VARIOUSCURRENTDENSITIES
O
PLAN FOR FUTURE WORK
At present work is under progress exploring thc temperature region between 0" and 60" C. I t is anticipated that acid concentrations between 1.05 and 1.15 specific gravity q-ould permit higher ozone yields than those reported. Determination of the relative amounts of ozone and oxygen liberated and the persulfate formed during the electrolysis will be important steps toward the clarification of the mechanism of anodic ozone formation in sulfuric acid electrolyte. I 50
I
I00
TIME, MINUTES
Figure 4.
LITERATURE CITED
150
Cell Potential us. Time during Approach to Steady State Conditions at Constant Current Density
(1) Briner, E., Haefeli, R., and Paillard, H., Helu. Chim. Acta, 20, 1510-23 (1937). (2) Briner, E., and Yalda, A., Ibid., 25, 1188-1202 (1942).
September 1952
INDUSTRIAL AND ENGINEERING CHEMISTRY
(3)Fischer, F., and Bendixsohn, K., 2. anorg. Chem., 61, 13, 153 (1909). (4) Fischer, F.,and Massenez, C., Zbid., 52,202,229 (1907). Beta, H. F., and Maron, S. H., J. Am. Chem. SOC., (5) Gable, C. M., 72,1445 (1950). (6)Grafenberg, 2.anorg. Chem., 36,355 (1903). (7) Kremann, Ibid., 36,403(1903). (8) Lash, E.I., Hornbeck, R.D.,Putnum, G. L., and Boelter, E. D.. J . Electrochem. Soc., 98,134 (1951). (9) Latimer, W. M., “Oxidation Potentials,” New York, Prentioe Hall, 1938. (10)Maohu, W., “Das Wasserstoffperoxyd und die Perverbindungen,” Berlin, Springer Verl.. 1937.
2211
(11) McLeod, Chem. SOC.J . , 49, 591, (1886). (12) Malquori, G.,Attiaccad. naz. Lincei, 33,ii, 112-16 (1924). (13) Putnam, G., Moulton, R., Fillmore, W., and Clark, L., TTm8. Electrochem. SOC.,93,211-21 (1948). (14) Schonbein, Pogg. Ann., 50, 616 (1840). (15)Soret, Compt. rend., 56,390 (1863). (16) Targetti, Nuovo cimento, 10,360 (1899). (17) Thorp, C.E.,IND. ENG.CHEM.,ANAL.ED.,12,209 (1940). (18) Wartenberg, H., and Arohibald, E. H., 2. Electrochem., 17, 812 (1911). (19) Wulf, 0. R.,and Tolman, R. C., J. Am. Chem. Soo., 49, 1660. (1927). RECEIVED for review July 13. 1951.
ACCEPTEDFebruary 4. 1H52
Acid Composition of Gum Spirits of Turpentine and Low Wines E. D. PARKER AND L. A. GOLDBLATT Naval Stores Research Division, Naval Stores Station, Olustee, Fla, OMMERCIAL gum spirits of turpentine generally contain a certain amount of acidic materials whose presence may contribute t o the deteriorative changes that take place in the quality of turpentine during storage. The gum naval stores industry is interested in reducing the acidity of turpentine, especially since the Commodity Credit Corp. has placed a maximum limit on the acid content of turpentine (an acid number of 0.5) that is acceptable for a government loan. Consequently, some work of this laboratory has been directed toward solving the problem of turpentine acidity. T o aid in this work, through obtaining a better understanding of the causes and the prevention of acidity, the composition of the acidic components of gum spirits of turpentine was investigated. The acidic components both of low wines-the aqueous phase of turpentine distillation-and of turpentine tailings were studied. To obtain typical samples, low wines and turpentine from commercial distillations were used. I n preliminary work several methods for the analysis of organic acids were tested, including Duclaux numbers (%), analytical steam distillation by the method of Olmstead, Whitaker, and Duden ( 4 ) , the fractional extraction procedure of Bush and Densen ( I ) , and the chromatographic procedure of Ramsey and Patterson (5-7), and of Marvel and Rands ( 3 ) . The most useful were found t o be the procedure of Marvel and Rands and that of Ramsey and Patterson, the latter somewhat modified. These two, with the procedure of Bush and Denson ( I ) , were applied in the work reported. LOW WINES
Low wines of 0.018 N acid PROCEDURE FOR FRACTIONATION. concentration from a commercial still were made alkaline with an excess of sodium hydroxide and were concentrated by distillation to about 3% of the original volume. Five fractions of the sodium salts of this material were obtained as follows: A 10-liter sample was concentrated and saturated with carbon dioxide, and the precipitate was extracted with alcohol t o yield fraction 1. The filtrate was further concentrated and diluted with alcohol and acetone in a stepwise procedure, and the supernatant solution was finally evaporated t o dryness t o obtain fractions 2 through 5 . Each salt fraction wae acidified with sulfuric acid and the acids were separated by distillation in vacuo, The five distillates, or acid fractions (Table I), were obtained as concentrated aqueous
solutions, the per cent acid of the fractions ranging from 37 t o 43.6% (calculated as acetic acid). An over-all yield of 72% of the acids present in the original low wines was obtained. ACID CONCENTRATIONS. Averages of t h e concentrations of acetic and butyric acids, in the five fractions, roughy calculated on the basis of the multiple fractional extraction procedure of Bush and Densen (I), as a preliminary approach t o the problem, showed about 90% acetic and 7% butyric acids. A disadvantage with this method is that the values for acetic acid include any formic acid present. Fraction 5 was estimated t o contttin about 70% acetic acid and 30% butyric acid, disregarding t h r other minor constituents. T o determine more accurately the composition, a coniposite of the five acid fractions was prepared by mixing together a volume of each fraction proportional t o the total volume of that fraction, resulting in a mixture with the same composition as the total distillate. A part of this composite solution, 0.1 ml., was absorbed on a column of 20 grams of silica gel containing 12 ml. of water, and was developed with chloroform and butanol by the method of Marvel and Rands ( 3 ) . Titration of the effluent indicated 5% butyric and higher acids (peak effluent volume or P E V of 50 ml.); 0.2% propionic acid (PEV, 130 ml.); 88% acetic acid (PEV, 160 ml.); and 770 formic acid (PEV, 270 ml.). Fraction 5 , 0.1 ml., was chromatographed in the same way and was found to contain 29% butyric and higher acids (PEV, 40 ml.); 0.9% propionic (PEV, 110 ml.); 40% acetic (PEV, 160 ml.); and 30% formic acid (PEV, 260 ml.). The acids from formic t o butyric were studied by the proce\lure of Marvel and Rands (3). Butyric and the higher molecular weight acids were separated from fraction 5 by a modification of the Ramsey and Patterson
TABLEI. CONCENTRATED AQUEOUSACIDS OBTAINED FROM FRACTIONATED SODIUM SALTS ’
Fraction
1 2 3 4 6
Volume. M1. 402 55 72 44 104
Density, di4 1,046 1.054 1.049 1.044 1.037
Refractive Index, ng0 1.3623 1.3603 1.3595 1.3655 1.3627
Total Acid (as Acid, Acetic), EquivaN % lents 7.62 43.6 3.06 7.47 43.0 0.41 7.31. 41.7 0.63 6.36 36.5 0.28 6.41 37.0 0.67
Concn.,
