Rate of carbon Dioxide Absorption
by Carbonate Solutions in a Packed Tower
CHARLES S. COMSTOCK' AND BARNETT P. DODGE Yale University, New Haven, Conn.
Results from the study of various factors affecting the rate of absqrption of carbon dioxide in solutions of sodium and potassium carbonates are presented. The apparatus used was a 3-inch tower, packed with small glass rings and accompanied by the usual auxiliary equipment for measurement and control of the important variables. The effects of the following variables on the absorption rate coefficients were studied : concentration of carbonate solution, temperature, rate of liquor flow, rate of gas flow, per cent of carbon dioxide in the gas, degree of conversion to bicarbonate, pH of the solution. A few examples are given to show how the resulting data may be of use in the design of commercial scrubbing towers.
water. (The difference between molality and moles per liter a t the highest concentration employed is only about 4 per cent.) These variables are not all independent. For example, the pH of the solution depends on variables 1, 2, and 6. Furthermore, variables 5, 6, and 7 change throughout the tower since we are dealing with an integral effect rather than the differential effect a t one level in the tower. The values for these variables will be arithmetic averages of the terminal conditions. With so many variables it would have been almost impossible to cover the field completely, and i t was necessary to limit the number of constant conditions more or less arbitrarily. For example, the effect of concentration of the carbonate solution was investigated a t only a few combinations of the other variables. Even with this considerable limitation of the field, approximately 150 runs were made and data for 111 of them are included in this paper.
Review of Previous Work The present investigation was, in a sense, a continuation of the work of Payne and Dodge (18). These investigators worked only with sodium carbonate solutions; but, since they gave a rather complete review of the literature on carbon dioxide absorption, it will be necessary to mention only those papers which have appeared recently. Harte and Baker (1) working with carbonate solutions from 1 to 2 N in sodium, over a temperature range of 25" to 63" C. in a small wettedwall tower, found that the absorption coefficient, &, was independent of sodium concentration, decreased considerably as the bicarbonate normality increased from 0.2 to 0.6, and doubled for a temperature rise of approximately 22" C . Hitchcock (8, 9) and Hitchcock and Cadot (10) reported results on the effect of alkali concentration on the steadystate initial rate of absorption of carbon dioxide in solutions of sodium hydroxide, potassium hydroxide, sodium carbonate, and potassium carbonate, using a small batch absorption apparatus with stirred liquid phase. Although these data have no direct bearing on the present investigation, they represent some fundamental facts which may prove useful, a t least qualitatively, in interpreting results obtained in a more complex system such as has been studied here. Hatta (6) studied absorption of carbon dioxide from air-carbon dioxide mixtures by potassium carbonate solutions in a small batch absorber similar in principle to that used by Hitchcock. The variables investigated were temperature, composition of the solution, carbon dioxide partial pressure, concentration of the solution, and velocity of liquid stirrer and of gas-phase stirrer. Hatta and Baba (6)investigated the rate of carbon dioxide absorption from bubbles rising through potassium carbonate solutions, varying the carbon dioxide pressure, alkali concentration, and temperature. Hatta and Katori (7)absorbed carbon dioxide by water flowing in a thin layer.
T
HE absorption of carbon dioxide from gaseous mixtures by solutions of sodium or potassium carbonate is an operation of considerable industrial importance, and yet the fundamental factors governing it. have not been well understood, a t least in their quantitative aspect. The objective of the research which this paper summarizes was the elucidation of some of these factors by small-scale experimentation on apparatus which is geometrically similar to that used industrially and then the interpretation of their meaning in terms of a unit of commercial size. A later paper will present similar results for absorption of carbon dioxide by solutions of the ethanolamines. Most of the work was done with potassium carbonate solutions though a few experiments were made with sodium carbonate. The data on rate of absorption in potassium carbonate solutions were the more meager, and furthermore the potassium compound is probably more used industrially at present. Investigation has been made of the effect of the following variables on the rate: 1. Concentration of the carbonatesolution (0.05 to 1.60 molal) 2. Temperature (15' t o 75" C.) 3. Rate of liquor flow (30 to 800 llters per minute per square meter or 0.74 to 19.7 gallons per minute per square foot) 4. Velocity of gas flow (3.6 to 20.3 cm. per second or 0.12 to 0.67 foot per second in an open tower) 5. Percenta e of carbon dioxide in the gas (4.6 to 100 per cent) 6. Degree of conversion t o bicarbonate 7. pH of the solution The figures in parentheses show the maximum range over which each variable was studied. AI1 sohtion concentrations are reported as molalities, or gram-moles per 1000 grams of 1
Present address, Merrimac Chemical Company, Boston, Mass.
520
MAY, 1937
INDUSTRIAL AND ENGINEERING CHEMISTRY
Experimental Method Figure 1 is a diagram of the arrangement of the apparatus.
521
solutions of approximately the same molality. Column 2 gives the total molality on the basis that all the potassium is in the form of potassium carbonate. The temperature in column 3 is the arithmetic mean of the two measured liquor temperatures. Liquor rate and gas velocity are based on the cross-sectional area of the unpacked tower. The carbon dioxide balance (column 10) is the amount of carbon dioxide gained by the liquor, as determined by liquor analysis, divided by the amount lost by the gas, as determined by the gas analysis, times 100. The liquor analyses are considered to be somewhat more accurate than the gas analyses and so the amount absorbed and the coefficient are based on the liquor analyses except in the very few cases where one of the liquor compositions was not determined. Column 12 gives values of an over-all absorption coefficient defined by the equation:
It is similar to that used by Payne and Dodge (18) but improved in various respects : The absorption tower, L, was constructed from 3-inch (7.62cm.) standard pipe and flanged fittings and was lagged to prevent heat loss. It was packed with glass rings 10 mm. 0. d. X 8 mm. i. d. x 10 mm. long, for a length of 10.67feet (352.5cm.) the total volume of packed s ace being 0.548 cubic foot (15.5 liters). A is a 50- allon (189-litery drum serving as a feed tank. It was e uipped wit! a portable motor-driven stirrer and a pipe coil througx which either steam or cold water could be circulated for control of the tem erature. From the feed tank the liquor flowed to constantlever vessel I where the liquid level was held constant by a float control. Different constant-liquor rates were obtained by changing the orifice inserted between flanges at J. The enlarged tube, K , allowed all entrained air bubbles to separate and escape from a vent before the liquor was delivered t o the tower through the gooseneck seal. From the base of the tower the liquor overflowed through a siphon t o tank 0 from which it could be returned to the feed tank by centrifugal pump P . The caibon dioxide-air mixture was produced by mixing metered streams of the indiwhere Ka = absorption coefficient vidual gases, the air being su plied by a small positive-pressure a = effective interfacial area per unit of tower volume blower and the carbon dioxig being drawn from a cylinder of W/' = weight of COI absorbed er unit of time the liquid through a standard reducing valve. The maintenance = active volume of tower 8 5 . 5 liters) of a constant rate of flow of the carbon dioxide was aided by placAps,. = log mean of terminal driving forces in pressure units ing a small electric heater so that it radiated heat to the regulating valve. The air, after measurement by a calibrated orifice meter, Generally K o and a are not separated, and the combination is went to the saturator, 2, where it passed countercurrent to a also called an absorption coefficient. Values of Koa are in stream of water flowing over glass packing. For ruas at elevated grams/(hour) (cc.) (atmosphere). I n most cases the arithtemperatures the water for the saturator was preheated by steam metic mean AP was used, since it differed very little from the in tank A A . The carbon dioxide was added to the air after the saturator, the mixture assloe mean. Calculation of ing through packed cEamthe driving forces requires ber V for thorough mixing. a knowledge of the equiThermometer wells w e r e librium vapor pressure of provided to take the temperature of liquid as it encarbon dioxide over cartered and left the tower and bonate-bicarbonate soluof the entering gas. Gas tions. In the case of samples were taken at four sodium solutions the equaintermediate points in t h e tower as well as at the two tion of Harte, Baker, and t e r m i n a l points. The Purcell (6)was used, and bottles EE, in which the gas for potassium solutions samples were collected, were initially filled with water, the data of Sieverts and and the sample was allowed F r i t z s c h e (60). Both to displace the water comcover a sufficient range so pletely and then to flow for that the d r i v i n g forces some time longer to flush out the first portion of gas could be definitely estabwhich may have had its lished for all of the expericomposition c h ~tn g e d by ments. partial s o l u t i o n in the The correlation of the water. The gas s a m p l e s were present data by this means a n a l y z e d by means of a is admittedly purely emrtable Orsat apparatus. pirical and not related to &e analysis of the gas enthe m e c h a n i s m 'of the tering the tower furnished a check on the flowmeter process. No one has yet readings since the gas comdeveloped a completely position could obviously be s a t i s f a c t o r y theory to caJculated from them; in treat the case where physin e a r 1 all cases the two metho& agreed to 0.2 per cal absorption is accomc e n t c a r b o n dioxide or panied by a r e v e r s i b l e better. Liquor samples chemical reaction, and the were analyzed for carbonate w r i t e r s b e l i e v e that, and bicarbonate as soon as possible after completion of 1a c k i n g such a theory, a run by the usual Winkler the data should be prebarium carbonate method. sented in a s s i m p l e a FIGURE 1. DIAGRAM OF ABSORPTION APPARATUS The method was checked manner as p o s s i b l e to on known solutions made up from pure salts, with make them useful for apvery satisfactory results. dication. Since, as will be shown later, the absorption of carbon Experimental Data and Calculations dioxide by carbonate solutions is controlled by conditions within the liquid phase, it might appear more logical to emTable :I gives the experimental data for potassium carbonate ploy an over-all coefficient based on liquid-phase driving solutions and Table I1 for sodium carbonate. Table I is forces instead of the one based on partial pressures. The subdivided into several groups, each group containing data on
-
522
INDUSTRIAL AND ENGINEERING CHEMISTRY
VOL. 29, NO. 5
TABLEI. DATAON RATEOF ABSORPTION IN POTASSIUM CARBONATE SOLUTIONS~ Run
No.
Molality
Temp.
c.
135 134 128 130 133 129 131 174 173 172 220 219 218 136
0.091 0.091 0,078
182 181 180 183 179 178 177 170 184 185 186 187 188
0.118 0.118 0.118 0.112 0.134 0.134 0.134 0.114 0.112 0.112 0,119 0,121 0.121
20.0 20.0 20.0 20.0 19.5 19.5 19.5 16.5 26.1 33.5 40.0 48.5 57,5
169 168 167 166 138b 139b 142 145 146 147 143 144 148 140b
0.151 0.151 0 151 0.151 0,162 0.162 0.166 0.160 0.160 0.160 0.166 0.166 0.160 0.166
16.0 15.0 15.5 16.0 19.. 0 18.5 20.5 14.0 16.5 15.0 19.8 20.1 15.8 19.5
0,086
0,086 0 078 0.086
0.074 0.074 0,074 0,039 0.052 0 075 0,091
Liquid Rate L./min./sq. m.
15.0 15.0 16.5 18.0 18.0 16.5 18.5 19.0 17.5 17.0
Grts
Velocity Cm./sec.
%
88.0
98.4 101.0 98.7 104.0 96.1 102.1 100.0 99.5 102.3 100.0 101.6 97.9 100.0 97.3
0.0441 0.0560 0.0272 0.0564 0,0564 0.0423 0.0722 0.0739 0.0767 0.0780 0.0119 0,0220 0.0429 0,0967
0.043 0.067 0.071 0.084 0.083 0.078 0.076 0.067
0.0236 0.0471
0.120 0.103 0.097 0.089 0,087 0,081 0.081 0 . 167 0.094 0 . 101 0.103 0.137 0.173
0.0
2.7
10.7 11.4 11.9 11.1 13.1 14.6
6.6
6.4 13.9 19.1 24.0 38.9 49.0 48.2 24.2 24.7 24.5 27.3 25.1 25.5
1.9 1.9 1.9 2.0 2.5 2.5 2.5 2.8 2.0 2.0 3.2 3.3 3.3
17.4 33.0 41.9 56.6 70.2 83.2 83.2 49.0 60.7 65.0 66.1 72.1 80.0
92.4 96.2 101.0 98.7 91.5 99.3 99.6 93.0 95.6 95.4 90.0 98.1 91.7
10.4 10.6 10.5 10.7 10.8 10.7 10.8 10.7 11.3 5.3 12.1 10.9 13.6 10.6
23.4 24.8 24.0 24.6 25.0 25.4 25.6 10.7 14.3 21.0 18.7 25.8 25.6 25.8
22.0 23.2 22,o 21.6 19.6 20.0 19.8 7.0
1.4 1.4 1.4 1.4
101.2 101.0 100.0
0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0
100.0 84.1 71.0 61.1 52.2 54.0 50.2 19,8 26.0 31.0 31.4 42.0 42.2 32.8
25.2 25.2 20.5 17.7 35.1 10.5 18.5 33.4 10.2 33.5 10.9 25.0 26.0
21.5 19.9 16.5 10.4 18.0 3.9 13.1 18.0 3.1 13.0 2.4 9.6 8.9 24.6 23.2 22.0 21.9 20.5 19.6 10.4 7.3 10.0 13.4 20.6 18.8 38.0 100 100 12.6 17.0 10.9 4.9 8.9
0.2 0.2 1.1 1.1 1.1 0.2 1.1 1.8 1.8 13.6 1.2
6.7 11.2 6.7 6.7
159 158 211 214 243 215 212 245 217 244 216 189 190
0.257 0.257 0.340 0.340 0.342 0.340 0.340 0.343 0.340 0.342 0.340 0.280 0.280
21.0 21.9 41.0 40.0 41.9 39.5 56.0 53.7 56.2 54.2 56.0 64.0 75.2
123 187 63 274 322 449 63 191 222 318 449 363 363
10.9 11.0 12.7 12.3
191h 192 193 194 156 155 153 151b 152 149 154 150 196 195 199 21 157 22 197 198
0,530 0.530 0.530 0.530 0.458 0.450 0,450 0.404 0.404 0.404 0.450 0,404 0.532 0.530 0.487 0.515 0.458 0.454 0.532 0.531
19.5 18.0 20.0 20.0 22.0 17.0 16.9 22.0 22.0 22.2 16.8 22.2 22.0 22.0 20.5 21.0 19.6 18.8 60.5 65.7
10.4 50 92 124 152 228 324 323 324 324 324 324 341 343 343 372 540 794 338 330
11.6 11.6 11.6 9.1 10.9 10.7 11.0 11.7 12.2 13.7 10.9 4.1 3.6 3.6 13.1 10.6 11.2 11.1 15.8
25.5 25.5 25.3 26.3 25.2 25.5 21.3 11.2 14.7 18.6 25.8 25.8 73.0 100 100 18.8 25.2 23.1 25.6 27.9
161 160
0.641 0.641
14.5 16.5
148 324
10.7 10.9
25.0 24.9
162 163 235 236 237 238
0.857 0.857 0.814 0.808 0.814 0.818
20.0 19.0 22.0 23.0 23.0 24.5
146 338 322 322 322 318
10.9 10.8
239 241 242
0.817 0.848 0.830
23.8 25.0 25.0
318 318 318
231 232 234
1.051 0.984 0.946
24.0 24.0 23.0
328 322 322
8.6
11.2 13.8 9.9 13.1 9.9 12.9 16.2 20.3
6.6
9.6 18.1 23.9 33.1 32.5 15.0 18.0 17.5 19.0 16.3 15.7
8.8
9.2 13.4 18.1 20.0 15.3
KGa
% 73.5 36.4 57.8 57.9 53.1 74.0 78.4 81.5 82.5 18.8 23.0 31.6 55.9
6.8 6.8
GO2
hbsorption Rate U./eec.
% 0.0 0.0
8.6
30 50 85 120 192 189 236 347 347 347 347 347 347 560
%
c0 2
Balance
21.3 20.1 7.2 13.1 10.2 10.7 18.7 19.0 21.7 23.9 1.5 2.6 6.8 16.6
11.0 9.2 9.5 9.1 10.5
18.0 17.0 14.0
%
Conversion to -Bicarbonate--Entering Exit liquor liquor
25.0 25.0 10.1 18.3 21.1 14.6 25.0 29.9 29.8 30.0 3.0 5.6 12.7 25.0
10.5 10.5 10.8 12.1 5.4 11.3 10.7 6.4
18.0
--
Goa---Enterinn Exit gas gas
5.6
0.0 0.0
5.6 0.0 3.1 3.1 3.1 5.0 4.7 4.0
0.0600
0.0810 0.1180 0.1398 0,1406 0.1438 0.0878 0.0927 0.0964 0,1082 0,1205
0.068
0,067 0.123 0.124 0,113 0.106
112.0 107.8 100.7 96.7 97.0 100.5 100.7 99.5 100.1 105.0
0.0159 0.0195 0.0315 0.0386 0.0642 0.0630 0.0691 0,0396 0.0519 0.0621 0.0638 0.0850 0.0845 0.1076
0.016 0.019 0.032 0.038 0,067 0.064 0.070 0.103 0.103 0.091 0.092 0.089
2.4 2.4 2.9 2.6 2.2 2.9 2.9 2.5 2.9 2.2 2.9 3.0 3.0
49.6 39.4 51.2 26.1 38.1 14.1 60.1 59.9 25.8 45.3 18.5 51.5 56.0
87.8 99.2 83.5 98.7 100.0 99.0 99.9 103.0 97.9 102.0 100.5 95.8 93.8
0.0520 0.0619 0.0358 0.0755 0.1368 0.0589 0.0425 0.1290 0.0640 0.1606 0.0819 0.1775 0.1940
0,052 0.064 0.050 0.131 0.128 0.204 0.075 0.142 0.264 0.188 0.338 0.322 0,450
1.4 1.4 1.7 1.7
38.6 33.0 28.4 23.3 24.0 20.2 11.3
52.5 95.8 106.0 96.9 98.9 97.8 102.5
l2:4 15.2 16.9 19.4 31.9 48.1
103:5 100.1 95.7 99.1 98.5
94.0
0,086
0.119
1.0 2.3 2.4
12.1 10.8 10.5 28.6 33.8
86.9 102.6 87.5 105.9 101.2
0.0128 0,0288 0,0443 0,0487 0.0565 0.0700 0.0549 0.0386 0.0511 0.0633 0 0825 0,0820 0.1869 0.2520 0.3010 0,0706 0,0905 0,1170 0.1613 0,1901
20.1 19.4
1.0 1.0
17.7 11.7
101.7 105,6
0.0541 0,0759
0.059 0.055
25.3 25.2 58.0 46.6 43.3 38.9
20.3 18.4 49.5 37.6 33.5 30.7
1.1 1.1 30.0 36.8 41.0 53.0
16.5 9.7 43.5 47.2 49.9 60.3
90.1 95.3 99.6 99.5 101.0 97.9
0.0660 0 0850 0.1214 0 0940 0 0808 0 0635
0.090 0.090 0.054 0.053 0.051 0.044
6.4 ..
..
42.6 54.0 54.0
31.4 48.2 50.4
54.0 74.9 86.7
60.4 82.0 90.0
102.3 100.1 99.3
0,0577 0.0700 0.0294
0.039 0.036 0.018
10.9 8.6 7.3
24.7 32.7 40.6
18.3 23.3 30.1
3.6 5.6 16.3
9.9 14.2 25.7
100.0 103 0 104.1
0 0745 0 0924 0 0968
0.083 0.078 0.065
5.6
, .
6:4 7.7
0.0
0.0
66.0
... ...
0 0621 98.9 8.1 0.4 25.5 20.4 151 10.8 19.5 1.598 165 0.0826 101.6 18.0 0.4 5,O 19.5 338 10.8 25.3 -. . ~ . . 164 1.598 a Conversion factors: 1 liter/min./sq. meter = 0.0246 gal./min./sq. ft.; 1 em./sec. = 0.0328 ft./seo.: Kua in (grams) (hr.)-'(cc.) -l(atm.) K G in~ (lb.)(hr.)-l(cu. ft:) -l(atm.) -1; KGa In (grams) (hr.) -1(co.) -'(atm.) -1 X 1.42 = KGa in (1b.-mole COz) (hr.) -l(Cu. ft.) -'(atm.) -I. 6 Kaa and the absorption rate are based on the gas analyses. The poor balance for run 191 is beheved t o be due to lack of a good liquid sample of the low liquor rate.
0,063 0.090 X 62.5
-
a8 a result
MAY, 1937
INDUSTRIAL AND ENGINEERING CHEMISTRY
523
TABLE 11. DATAON RATEOF ABSORPTION IN SODIUM CARBONATE SOLUTIONS Run No.
Molality
Temp. a
C.
24.5 0.579 21.5 0.548 25.8 0.749 25.5 0.636 25.0 0.441 20.6 0.208 25.7 0.684 25.7 0.749 24.5 0.460 22.6 0.549 20.0 0.548 22.2 0.216 19.5 0.487 19.5 0.487 24.5 0.669 22.5 0.549 24.7 0.669 23.9 0.493 25.5 0.596 24.5 0.399 25.6 0.499 Based on gas analysis,
56 24 38 44 35 14 40 39 36 27 25 12 16 18 46 26 47 31 53 37 30 6
Liquor Os! Rate Velocity L./min./sq. m. Cm./sec.
Conversion to -BicarbonateEntering Exit liquor liquor
-COS--? Entering
Exit
gas
gas
%
%
%
%
24.2 20.3 26.0 9.8 23.8 22.4 25.1 24.1 23.2 25.4 20.3 19.1 19.6 19.8 24.4 13.7 30.0 21.2 40.9 22.5 21.2
20.0 17.7 22.0 7.9 20.1 18.7 21.4 19.5 18.0 19.7 15.9 15.3 15.8 16.0 20.4 9.9 24.4 17.0 32.4 18.0 15.4
1.3 1.5 1.2 14.2 1.7 3.7 9.4 1.2 1.3 1.7 1.5 4.4 3.5 3.6 21.9 1.7 21.9 5.1" 34.2 1.0 2.3a
22.9 15.5 15.4 20.3 22.6 38.5 21.4 12.2 18.1 13.5 11.3 28.0 13.9 13.0 30.2 8.8 32.9 15.1 51.3 15.3 9.8
11.1 13.1 13.7 11.1 15.0 14.7 13.6 13.6 14.4 10.6 13.1 13.8 13.2 13.2 14.7 12.3 11.3 14.0 9.4 14.5 14.1
difficulty is that all the complication occurs in the liquid phase, and it is practically necessary to assume some kind of a mechanism before one can decide what to take as the driving force. The gas-phase conditions are much simpler and one can easily visualize a driving force without recourse to the multiple films necessary to treat the liquid phase. This is the justification for use of a g a s - f i l m coefficient to correlate the data. The theory of Hatta (4, which is based on the idea that there is a "chemical resistance" in series with t h e d i f f u s i o n a l resistance, leads to an eauation which re. . TDmperoivre 19%. d;ces to the ordinary absorption equation PERCENT CONVERSION based on pure diffuTO BICARBONATE sion, if the rate of FIGURIC 2. VARIATION O F KcU THROUGHOUT T H E TOWERI N ABthe reaction SORPTION OF CARBON DIOXIDEBY is slow. T h e r e is POTASSIUM CARBONATE SOLUTION reason to believe that the chemical reaction rates in the systems here considered are very slow, and hence the controlling factor may be the physical diffusion of carbon dioxide across a diffusion layer. Since Henry's law is closely followed by solutions of carbon dioxide in water or in neutral salt solutions, it seems reasonable to assume it for the free or nnreacted carbon dioxide in the case of carbonate solutions. In cases of purely physical absorption where Henry's law does apply, we know that K L = HKa, where H is the Henry law constant. I n such a case i t does not matter which over-all coefficient is used, since they are proportional. This may furnish additional justification for the use of the gas-film coefficient in the present instance. '
Effect of Conversion to Bicarbonate The absorption coefficient varies not only from one run to another, owing to premeditated changes in certain variables, but also changes markedly throughout the tower in a given run. This effect, which is due to the changing ratio of bicarbonate to carbonate, is illustrated in Figure 2 . It is based on analyses of gas samples taken a t four intermediate points in the tower in addition to the data for the two terminal points.
COS
Balance
% 97.8 99.6 95.9 102.7 104.1 90.0 97.8 97.8 93.0 105.4 103.2 107.0 100.0 104.5 92.3 101.5 103.0
...
96.9 101.0
...
c02
Absorption Rate G./sec.
K aa (U.)(h?.) -1 (cc.) - ' ( a h . )
0.0502 0.0358 0.0564 0.0188 0.0609 0.0512 0.0535 0.0630 0.0708 0.0674 0.0615 0.0604 0.0515 0,0536 0,0640" 0.0450 0.0737 0.05980 0.0954 0.0953 0.0813"
-1
0.052 0.043 0.054 0.048 0.065 0.060 0.054 0.068 0.081 0.069 0.081 0.083 0.060 0.078 0.066" 0.088 0.063 0.074' 0.061 0.108 0.105"
The short horizontal lines show the terminal values of per cent conversion for each section. Such intermediate point data were taken in practically all of the runs, but because of space limitations they are not included in the tabulation of the data. The data presented in Figure 2 are typical of a great many runs and show clearly the considerable effect of bicarbonate in decreasing the coefficient. This was first clearly shown for sodium carbonate solutions by Payne and Dodge (18) and was subsequently confirmed and presented in greater detail by Harte and Baker (1). Because of this important variable it is necessary, in comparing the results of different runs, to have the average percentage conversion about the same in the runs under comparison. Even this involves an approximation when the limits of conversion are wide. Thus an average of 50 per cent conversion might mean terminal conversions of 0 and 100 per cent, respectively, or of 49 and 51 per cent. It would have been desirable from this standpoint to choose conditions which would have given a fairly narrow range of conversion in a given experiment, but other factors governed this choice a t the time the majority of the runs were made. 0.14
A L L P O I N T S CORRECTED , TO 2 5 ' C . A N D
.12
ARITHMETIC
AVERAGE
PERCENT
LIQUOR
CONVERSION
FIGURE3. EFFECT OF AVERAGE PER CENTCONVERSION TO BICARBONATE O N Kca FOR ABSORPTION OF C A R B O N DIOXIDE BY POTASSIUM CARBONATE SOLUTIONS OF DIFFERENT MOLALITY
Figure 3 shows the effect of arithmetic average per cent conversion for the whole tower on the coefficient a t different molalities. The points spatter somewhat, which is not surprising when we consider that each point represents an average for the whole tower; nevertheless, the general effect is clearly shown. Harte and Baker ( I ) have made the most extensive
INDUSTRIAL AND ENGINEERING CHEMISTRY
524
previous study of this factor for sodium carbonate solutions, and the present results are quite similar to theirs. They used a wetted-wall tower so that the factor a could be evaluated, and expressed K Q as a function of bicarbonate normality; on that basis K G was substantially independent of total sodium normality. The present authors recorrelated their results on the basis of per cent conversion rather thannormality; as shown in Figure 4,a separation occurs into fairly well defined curves for each normality.
4 II
0.08,
I
I
I
I
I
I
I
I
some data on the pH of sodium carbonate-bicarbonate solutions a t 30" C. After making a slight correction for the M e r ence in temperatures, the present measured values are in good agreement with those of Kiehl and Loueks. Figure 8 shows the excellent correlation obtained when the coefficients of Harte and Baker (1) are plotted against the pH of the solutions. A similar correlation for some of the writers' own results on potassium carbonate solutions is shown in Figure 9. In this case pH is that of a solution whose composition is the arithmetic mean of the two terminal compositions. The convergence of the extrapolated lines to a common point probably has no theoretical significance, but it is used later in the development of an empirical equation.
Effect of Gas Velocity
u
% PERCENT CONVERSION TO NaHC03
FIGURE 4. RECORRELATION OF HARTE AND BAKER DATA ON ABSORPTION COEFFICIENT FOR CARBON DIOXIDEIN SODIUM CARBONATE SOLUTIONS IN WETTED-WALL TOWER ~
VOL. 29, NO. 5
~
~~
Effect of pH
The negligible effect of gas velocity on the coefficient is shown in Figure 10. This merely confirms the fact, already noted in previous investigations, that the absorption of carbon dioxide by carbonate solutions is entirely controlled by the conditions in the liquid phase. (Even for the absorption of carbon dioxide in solutions of the ethanolamines where the coefficient is much greater and therefore resistance in the liquid much less, gas velocity has only a slight effect.) One may safely conclude that the liquid phase is controlling for all conditions likely to be encountered in the absorption of carbon dioxide by carbonate solutions.
Effect of Carbonate Concentration Payne and Dodge (18) reported that K,u was independent of sodium carbonate concentration over a wide range, and Harte and Baker (1) confinned this observation over a somewhat more limited range. Present results on potassium car-
&a might also be related to the hydroxyl-ion concentration or to the pH of the solutions. Actual pH measurements on such solutions are few but for a limited set of conditions the pH of carbonate and of carbonate-bicarbonate solutions can be calculated from the equations given by Menael ( 1 7 ) , using the activity coefficients of Walker, Bray, and Johnston (22), the dissociation constants of McInnes and Belcher (16) for carbonic I acid, and the principle of ionic strength as enunn ciated by Lewis and Randall (16). No pretense is made that these calculations can yield anything more than approximate values, but they are probably sufficiently near the truth for the present pur105 ,MOLALITY pose which is to show that a good correlation exists between the pH and the absorption coefficient. The results of the calculations are presented in Figures 5 , 6, and 7 . Calculations could not be made for solutions stronger than 0.8 molal because the measurements of the activity coefficient were not carried beyond an ionic strength of 2.5 in molal units, which corresponds to about 0.8 molal concentration of either sodium or potassium carbonate. Although in pure carbonate solutions the pH, increases slightly with increase in concentration, the reverse is true in carbonate-bicarbonate solutions a t any given fraction converted to bicarbonate. Menael (17) made a few pH measurements by both colorimetric and electrometric means, and although the quantitative agreement with the writers' calculated values is only fair, there is good agreement on the general trends. A few pH measurements on potassium carbonate-bicarbonate solutions were made by the authors a t 25" C. by means PERCENT NalCOs CONVERTED of the glass electrode, using the purest salts obtainable. The general form of the curves of Figure 6 FIGURE6 . CALCULATED PH OF PURECARBONATE SOLUTIONS AT 25" C. was substantiated although the observed values of FIGURE6. CALCULATED CHANGE OF PH OF POTASSIUM CARBONATE SOLUpH averaged about 0.25 pH unit higher than the TIONS AT 25" C. WITH CONVERSION TO BICARBONATE BY THE REACTION: K?;COB COz HzO = 2KHCOs calculated values. After the above values of pH were obtained, a FIGURE7. CALCULATED CHANGE OF PH OF SODIUM CARBONATE SOLUTIONS AT 25" C. WITH CONVERBION TO BICARBONATE paper by Kiehl and Loucks (13) appeared with
+
+
MAY, 1937
INDUSTRIAL AND ENGINEERING CHEMISTRY
525 I
I
MOLAL11Y
y
90
s
10.0
ias
Pff FIGURE8. RECORRELATION OF HARTEAND BAKERDATAON BASIS OF PH OF SOLUTIONS
I1 0
MOLAL
0.074 MOLAL;
o'xl
7, C O N V E R T E D . TEMP IS'C, L I Q U O P RATE 325 8.5 % CONL'ERTEO TEMP 17'L L I Q U O R RATE 324. 4 2 . 8 % CONVERTED, TEMP 18.;; L I Q U O R RATTE 349
I
I
I
KtCOs
TION COEFFICIENT
0.086 M O L A L , 2 8 . 8
0 0.450
A
1'
FIGURE11. EFFECTOF MOLALITY OF POTASSIUM CARBONATE SOLUTIONAT C O N S T A N TAVERAGEPER CENT CONVERSION ABSORP-
11.5
110
I
ALL
0.4
MOLAL
c m w c x o TO K J O ~ AND 8.3
1P
I
OI,
J01 d, do
j.5
,io
i.6
i
PH
x"
GAS VELOCITY
FIGURE 9. KQUvs. PH FOR POTASSIUM CARBONATE-BICARBONATE SO L U T I o N s AT 25' C .
FIGURE 10. EFFECT OF GAS VELOCITY ON ABSORPTION COEFFICIENT
bonate solutions at first seemed to confirm this, but it was later found that, when the coefficients were compared a t the same average per cent conversion to bicarbonate, there was an effect of total potassium concentration. This effect has already been illustrated for sodium carbonate solutions by the recorrelation of Harte and Baker's data as presented in Figure 4. Figure 11 shows the effect of concentration of potassium sohtions for three different arithmetic mean per cent conversions; all values were corrected to a constant temperature of 25" C. and a constant liquor rate of 325 liters per minute per square meter. The decrease of the coefficientwith increasing concentration seems surprising a t first sight, but it is entirely consistent with the changes in pH and viscosity of the solutions. Although the pH of a pure carbonate solution increases with increase in concentration, the situation is reversed for solutions containing bicarbonate when the comparison is made a t constant per cent conversion as shown in Figure 6. The viscosity naturally increases with the concentration, and so both factors tend to cause the coefficient to decrease as the concentration increases. This effect is also consistent with the results of Payne and Dodge (18) who found the coefficient to be greater for pure water than for sodium carbonate solutions under comparable conditions. Hitchcock and Cadot (IO) found that the initial absorption rate in pure potassium carbonate solutions increases with concentration a t first and then passes through a maximum. The decrease was accounted for by viscosity, and it is possible that the changes in pH may account, for the increase a t low concentrations. Another factor which might play a part is the decrease in physical solubility of carbon dioxide in salt solutions as the salt concentration is increased.
Effect of Rate of Liquor Flow Liquor rate has a considerable effect on the coefficient as is illustrated in Figure 12 which shows Koa vs. liquor rate a t three different temperatures. An adequate theory of the effect of liquor rate in a packed tower has never been offered, as far as 6he writers are aware, although Hatta (3) recently
'
developed a theory for the case of abI 0,r 5m boo sorption by liquid LIQUOR RATE 23-( flowing in a t h i n layer when no FIGURE 12. EFFECT OF LIQUOR RATE chemical reaction is ON ABSORPTION COEFFICIENT i n v o 1v e d , which might be applicable. Data t o test it are still very meager. It is commonly assumed that increased flow results in greater turbulence and therefore in a thinner film and hence in an increased coefficient. This kind of reasoning is based on an analogy with heat transfer to fluids in forced convection or on studies of absorption in stirred batches of liquids. The conditions, however, in a packed tower are wholly different and the simple picture of a thin film in which all the concentration gradient is concentrated and yet which contains a negligible amount of the total material, with a thoroughly mixed main body of liquid, can scarcely be representative of conditions in a packed tower. All of the liquid in a packed tower is disposed in the form of a relatively thin a m . Using the figures on holdup of liquid published by Payne and Dodge (18) and assuming equal distribution over the surface of the packing, the thickness of the layer of liquid is of the order of 0.3 mm. a t about the mean liquor rate used in these investigations. It is difficult to imagine much convection or turbulence in so thin a layer, and the usual assumption that the concentration gradient is confined to a small fraction of this thickness seems untenable. A more reasonable explanation might be based on the following considerations: An increase in rate of liquor flow increases the thickness of the liquid layer on the packing, and in any given section of the tower more liquid will be available to take up a given quantity of carbon dioxide. Hence the concentration gradient of free carbon dioxide in the liquid will be greater and this will result in a greater absorption rate. Koa has been defined by the equation,
L zb. L i
w/O
Kaa(Po
- PL)V
526
INDUSTRIAL AND ENGINEERING CHEMISTRY
and since PL is calculated on the assumption that chemical equilibrium is established in the solution, (Po - PL)would in most cases be unaltered by the change in liquor rate and its resultant change in layer thickness. As a result, the calculated &a will be greater for the increased rate. Figure 13 may help to make these points clearer. L1 and Lz represent two different thicknesses of liquid layer. We will assume that the rate of reaction is relatively slow so that free carbon dioxide exists throughout the layer. Pa is the carbon dioxide partial pressure in the gas and also the pressure from the liquid a t the interface. (Po - PI,,) represents the carbon dioxide conc e n t r a t io n difference in the layer L1 in terms of gas partial pressures, and (Pa- PLJ that in the s e c o n d layer. The nominal driving force is (Po PL) and it would be substantially the % s a m e i n the two cases. Since W/B is -4 -7 greater in the case Lz THICKNESS OF LIQUID LAYER of Lz,i t i s c l e a r FIGURE 13. EFFECT OF LIQUOR RATE that KGu will also ON COEFFICIENT be greater for this case. If this picture of the effect of liquor rate is correct, one might expect that the coefficients would be the same for equal thickness of liquid layers (other things being equal) even though the actual total liquor rates are different. This g i v y a rough basis for the comparison of different packings: Thus, if packing A has half the surface area of packing B, the two will give the same coefficient per unit of area when the liquor rate over B is twice that 0ver.A. The proposed explanation of the effect of liquor rate might also lead to the conclusion that an increase in viscosity, since it also inci-eases the thickness of the liqdd layer on the packing, would increase the coefficient. This may possibly be true, and the effect previously ascribed to viscosity may be due to some other factor. Riou and Cartier (19)found that the addition of sucrose, dextrose, and glycerol to sodium carbonate solutions increases the rate of absorption of carbon dioxide; and sinoe their absorbent was exposed to the gas in the form of a thin film of liquid, it seems possible that the increase in thickness of the liquid layer due to increase in viscosity may have been the cause.
Effect of Temperature The work of previous investigators showed that temperature had a considerable accelerating effect on the absorption rate. The present investigations confirm this and offer quantitative data over a greater range of conditions. Figure 14 shows the effect of temperature a t various rates of liquor flow for a constant potassium carbonate molality of 0.40 and a constant arithmetic mean per cent conversion of 8.3. The effect of temperature is in general accord with the results of Williamson and Mathews (24), Whitman and Davis ( 2 ~ 3and ) ~ Hatta ( 5 ) . Sieverts and Fritzsche (WI), on the other hand, found little difference in the rate of absorption at 30°, 40°, and 50' C., with the lowest temperature giving highest rate. As will be shown later, the rate of absorption, as distinct from the coefficient, passes through a maximum and the position of the maximum will vary with conditions. This might account for the opposite effects of temperature found by different investigators, though it seems surprising that the maximum temperature should be as low as is indicated by the results of Sieverts and Fritzsche.
VOL. 29, NO. 5
Discussion of Results A word of caution should be given here in regard to the generalization of results on carbon dioxide absorption. Recent unpublished results of H. F. Johnstone, University of Illinois, indicate that carbon dioxide behaves in quite an anomalous manner as far as absorption rate is concerned, and conclusions drawn from the study of its absorption either by alkaline solutions or by water may not be applicable to other systems. The effect of all the important variables a t 25" C. can be compactly represented by the following purely empirical equation ; Koa = (c - b loglo x ) (pH - 8.0) where c, b = functions of liquor rate, s, in liters/min./sq. meter, a8
follows:
c = 0.0127
b
+ 0.000090 s
0.090001237 ~ 1 . 0
I
z = viscosity of solution, centipoises
The pH is taken from Figures 6 and 7 and refers to solutions whose composition is the mean of the two terminal values. The few data on the viscosity of potassium carbonate and of bicarbonate solutions given in the International Critical ALL DATA CORRECTED
TO
0 4 MOLAL
044- K2COJ A N D 0 3 PERCENT AVEPAGL CONVERSION TO KHCO,.
OM-
FIGURES ON
TO LIQUOR
CURVCS REFER RATE (,.+i)&%m,z
036 -
TEMPERATURE
"C.
FIGURE 14. EFFECTOF TEMPERATURE ON ABSORPTION COEFFICIENT
Tables indicate that the conversion of a potassium carbonate solution to a potassium bicarbonate solution is accompanied by practically no change in viscosity. Consequently we may use the viscosities of pure potassium carbonate solutions, as given by Hitchcock and McIlhenny (II), in the equation. The effect of molality on the absorption coefficient is taken care of by viscosity and pH. . This equation was compared with the observed data for eight different runs in which KGUvaried from 0.023 to 0.192 and the average per cent deviation and maximum per cent deviation of calculated data from observed values were 7 and 22 per cent, respectively. Comparison of the results for solutions of sodium carbonate and of potassium carbonate a t the same normality and under otherwise nearly identical conditions shows that the coefficient is slightly greater for potassium carbonate, 10 per cent a t the most. The difference is entirely accounted for by the
,
MAY, 1937
INDUSTRIAL AND ENGINEERING CHEMISTRY
effect of pH and viscosity (as shown by the fact that values of coefficient Kaa, calculated from the equation based on the results with potassium carbonate solutions, agree well with observed values for sodium carbonate solutions a t the same set of conditions). The pH of a sodium carbonate solution is a little less and the viscosity a little greater than that of a potassium carbonate solution of the same molality. The choice between the two absorbents would therefore be determine,d mainly by other factors. The chief factor, other than cost, is the solubility of carbon dioxide in such solutions. A 2.5 N sodium solution a t 30" C. will absorb a maximum of 2.85 cubic feet (at standard conditions) per gallon from a flue gas containing 17 per cent carbon dioxide, whereas a 2.5 N potassium solution will absorb 3.27 cubic feet. Furthermore, a considerably stronger potassium solution can be used because of the greater solubility of the potassium salts, and this results in a still greater difference in capacity for carbon di' oxide absorption. Thus a t 25" C. a sodium carbonatebicarbonate solution of 1.65 sodium normality and with 60 per cent of the sodium converted to bicarbonate is saturated with respect to sodium bicarbonate. The corresponding potassium normality is 4.1 (solubility data from International Critical Tables and from the original literature referred to in those tables). The carbon dioxide absorption in these two cases is 1.48 cubic feet per gallon for the sodium solution and 3.68 for the potassium solution. In other words, if precipitation of a solid phase is the limiting factor, the potassium solution will absorb approximately 2.5 times as much carbon dioxide as the sodium solution. There are no results of other investigators with which direct comparison can be made. Some discussion of the results of Harte and Baker ( 1 ) has already been presented, but numerical values of coeficients cannot be compared since ' their experimwts were made in a wetted-wall tower. Williamson and Mathews (24) investigated the rate of absorption of carbon dioxide by potassium carbonate solutions in a pebble-packed tower. It is difficult to make any quantitative comparison with their results because they reported rates of absorption rather than absorption coefficients, and the data are not given which would permit the conversion to be made. Furthermore, some of their results are given merely as statements without supporting data. They found that tho rate increased markedly with increase in liquor flow, decreased as the conversion to bicarbonate increased, and increased with temperature, passing through a maximum a t about 75" C. These facts are in qualitative agreement with present results. Hatta's experiments were made with stirred batches of liquid and hence cannot be quantitatively compared with those of the writers, but both sets of data are in agreement on the qualitative effects of some of the important variables, particularly those of temperature and degree of conversion to bicarbonate. On the other hand, Hatta found the rate of absorption to be independent of potassium normality over the range from 0.5 to 2.0 N , whereas the writers find that it decreases somewhat, other things being equal. Hitchcock and Cadot (10) found that the rate of absorption by sodium carbonate and potassium carbonate containing no bicarbonate increased a t first with the concentration of the solution, passed through a maximum, and then decreased. These results are consistent with the present ones, although the experimental conditions in the two cases are very different. The increase in rate a t low concentrations may be explained by the fact that pH increases with concentration in pure carbonate solutions whereas it decreases in solutions containing bicarbonate. These investigators also find that potassium carbonate gives a somewhat greater rate of absorption than sodium carbonate a t the same normality, in agreement with present results.
-
527
Application of Data to Plant-Scale Absorption Towers The most interesting application of the data and the objective for which this work was done is the prediction of plantscale tower sizes and operating conditions. Howe (12) gave a description with some operating data of the absorption system used a t a DryIce plant, and certain additional data were secured through a private communication (14). These data may be summarized as follows: Plant capacity, tons (metric tons) Coz/24 hr. 10 (9.06) No. of absorption towers in series 2 Dimensions of packed space in each tower: Diameter, f t . (meters) 10 (3.05) Height, ft. (meters) 90 (27.4) Operating temp., F. ( O C.) 120 (48.9) COZin gas entering towers, yo 19.0 COZ gas leaving towers, % 8.5 Tower packing Coke Av. liquor rate, gal. (liters) /min. 134 (506) Composition of solution, lb./cu. ft. (grams/liter) : Entering absorbers: NazCOa 5.72 (91.6) NaHCOa 2.65 (42.5) Leaving absorbers : NaSCOa 3.32 (53.2) NaHC03 5.81 (93.1)
From these data the following quantities may be calculated: Liquid rate, gal./min./sq. f t . (liters/min./sq. meter) 1.71 (70) Gas rate in open tower (based on dry inlet gas a t standard conditions), ft. (om.)/sec. 0.26 (7.92) Rate of COz absorption, 1b.-moles (gram-moles) /min. : From production figure, assuming 95'% recovery 0.332 (150.7) ' From data on liquid 0.337 (152.6) Conversion of Na t o bicarbonate, 'j&: At bottom of tower 52.5 At top 22.5 Arithmetic mean conversion 37.5 Na normality 2.20 Driving force in tower (difference between partial pressure of COZ in-gas and equilibrium pressure from solution), atm. : Bottom of tower , o . 110 Top of tower 0.0684 Mean A p 0.0892 Packed vol., cu. f t . (cu. meters) 14,120 (400)
Koa : Lb.-mole/hr./cu. ft./atm. = Gram /hr./cc./atm.
0'334 6o = 0.0159 14,120 X 0.0892 0.0112
The agreement between the absorption rate calculated from the production figure and that based on liquor rate and compositions indicates that the data are a t least consistent. In the absence of data on the coke packing i t will be assumed that the superficial area is 20 square feet per cubic foot (65.6 square meters per cubic meter).2 The glass-ring packing used in the present experiments was stated to have an area of 160 square feet per cubic foot (525 square meters per cubic meter). Making the rough assumption that the liquid rates are comparable when referred to the same packing area (based on the picture,.previously presented, of the effect of liquor rate of the coefficient), the rate for the experimental tower to compare with the rate of 70 liters per minute per square meter in the plant tower would be 560. Some rather wide gaps exist in the experimental data, and we must make some assumptions to arrive a t a predicted value of the coefficient. A graph of Kaa vs. liquor rate for sodium carbonate a t 25' C., 0.5 molal, and 10 per cent average
* The larger
of two figures in Perry's "Chemical Engineers' Handbook"
(New York, McGraw-Hill Book Co.. 1934) for 3-inch (7.62-cm.) coke, and presumably a reasonable figure for a coke of assorted sizes varying from 1.6 t o 3 inches (3.81 t o 7.62 om.).
.
INDUSTRIAL AND ENGINEERING CHEMISTRY
528
conversion shows that KQU= 0.10 gram per hour per cc. per atmosphere a t 400 liters per minute per square meter. Extrapolation to a liquor rate of 560 gives KQU= 0.12. A change from 0.5 to 1.1 molal potassium carbonate reduces the coefficient in the ratio of about 0.80, and the same ratio will be assumed for sodium carbonate. An increase in conversion from 10 to 38 per cent reduces the coefficient in the ratio of about 0.65. An increase in temperature from 25" to 49" C. increases the coefficient in the ratio of 2.1. Therefore the predicted value of KQUa t 49" C., 560 liquor rate, 1.1 molal sodium carbonate, and 38 per cent average conversion is 0.12 X 0.80 X 0.65 X 2.1 = 0.131. (The same figure is obtained by using the empirical equation given for potassium carbonate solutions, substituting the proper pH and viscosity for the sodium carbonate solution. This gives a value for 25" C. which must then be corrected to 49" C. as shown.) For the coke packing, assuming that the ratio of coefficients is that of the superficial area of the respective packings, this value becomes 0.131/8 = 0.0164. The agreement between this value and that previously calculated from actual plant data (0.011) is fairly good, considering all the assumptions involved, and lends confidence to the belief that conditions for largescale towers can be approximately predicted from these laboratory results. With the aid of these data and the assumption in regard 'to liquor rate, the effect of several important factors on the design and the performance of carbon dioxide absorption towers may be predicted. To illustrate, a few of the many possibilities are selected. For the calculations to be presented, it is assumed that the tower is packed with 1.5-inch (3.81-cm.) carbon Raschig rings (National Carbon Company No. 5) with a superficial area of 45 square feet per cubio foot (147.5 square meters per cubic meter). Liquid flow rates in the assumed commercial tower and the laboratory tower are compared on the basis of the assumption previously made-namely, that equal k w s over the same packing area give the same coefficients per unit of area, Other conditions assumed for the calculations may be summarized as follows: A t entrance to tower, 30 per cent At exit from tower, 70 per cent 3. Gas enters the tower at the average temperature of the tower, saturated with water vapor and containing 18 per cent carbon dioxide by Orsat analysis 4. Tower is substantially isothermal 5. Total pressure is that of the standard barometer 6. Gas flow rate is 0.5 foot (15.2 cm.) per second in the open tower referred to dry entrance gas at standard conditions 7. Equilibrium data of Sieverts and Fritzsche were used, assuming the constant to be independent of potassium normality 8. Liquor rate is 2 gallons per minute per square foot (82 liters per minute per square meter) 9. Active tower volume calculated frcun the equation
using values of &a estimated from the laboratory data, and log mean A p 10. Estimated values of Koa are as follows:
c.
25
40
50
60
0,0160 0.0113
0.0232 0,0165
0.0334 0,0236
0.0462 0.0326
K aa: Lb.-mole/hr./ou. ft./atm. Gram/hr./cc./atm.
on the dilution of the gases with water vapor, we would expect that there would be a minimum in the curve of height us. temperature. The calculated results are as follows: to
c.
Tower height: Feet Meters
25
40
50
60
393 120
336 102.6
316 96.5
580 177
Therefore, the most desirable temperature lies between 50 " and 60" C. This temperature will vary with the other conditions chosen to be constant, but not greatly. Commercial towers apparently operate a t about this temperature, although whether this is wholly fortuitous or the result of trial and error is not known. The calculated tower heights just given are excessive even a t the optimum temperature, and it is of interest to discover how much the height could be reduced by a reduction in the gas velocity. For this calculation, assume a gas velocity onehalf of the former value, or 0.25 foot (7.62 cm.) per second, a temperature of 50" C., and a carbon dioxide recovery of 66 per cent with other conditions the same as before, except the degree of conversion of the carbonated liquor which is no longer an independent variable and is calculated to be 50 per cent. The calculated tower height for these conditions is 95 feet (29 meters). It is concluded that a reduction in gas velocity will reduce the tower height more than in proportion, other things being equal as far as possible. It aIso shows why the gas velocity used in carbon dioxide absorption towers is relatively low. This velocity is apparently determined, not by pressure drop or entrainment, but by tower height. Suppose that it were desired to remove the carbon dioxide from 18 per cent down to a relatively low figure. Let us examine the possibility of doing this by absorption in potassium carbonate solutions. Equilibrium conditions set a lower limit to the percentage of carbon dioxide in the exit gases and the following tabulation shows the minimum possible amount of carbon dioxide in the gas for a 3 N potassium carbonate solution a t several temperatures and a t two different degrees of carbonation of the liquor entering the tower: to
c.
CO1 in dry exit gas, by vol: When liquor is 3 0 7 converted Whenliquoris 10% converted
1. Absorbing liquor is 3 N otassium carbonate 2. Carbonate converted to {icarbonate:
to
VOL. 29, NO. 5
The assumed conditions call for 6.90 per cent carbon dioxide in the gases leaving the tower or a carbon dioxide recovery of 66.3 per cent. With these conditions constant, the effect of different operating temperatures on the tower height may be calculated. Since the coefficient increases with temperature but the driving force decreases because of the effect both on the equilibrium pressure of carbon dioxide and
25 0.39 0.034
40
50
60
0.61 0.053
0.88 0.076
1.33 0.116
By stripping the liquor to a low bicarbonate content, it should be possible to remove the carbon dioxide to less than 1 per cent, but an excessive tower height would be necessary. For example, assume the carbon,dioxide content were to be reduced from 18 to 1 per cent (by Orsat analysis) by absorption a t 25" C. in 3 N potassium carbonate liquor flowing a t the rate of 2 gallons per minute per square foot (82 liters per minute per square meter) and varying in composition from 30 to 50 per cent conversion. The gas velocity to conform to these conditions would be 0.174 foot (5.3 cm.) per second in the open tower. The calculated height of tower, packed as before with 1.5-inch carbon rings is 364 feet (111 meters). Even if the entering liquor were stripped to 10 per cent conversion and allowed to rise to only 30 per cent converted, with other conditions the same as in the previous example, the required height is estimated to be 237 feet (72.3 meters). Furthermore, in both these cases the gas velocity is fixed a t 0.17 foot (5.19 cm.) per second which means a large-diameter tower for a reasonable throughput. Many other interesting calculations can be made with the aid of the data presented here, but these few illustrations will suffice to show the possibilities. Attention should be called to the great need for data on different types and sizes of packings. Without such data it is difficult to translate laboratory data on coefficients, secured in very small columns, into figures applicable to large towers and useful to the designer of such equipment.
MAY, 1937
INDUSTRIAL AND ENGINEERING CHEMISTRY
Literature Cited Harte, C. R., Jr., and Baker, E. M., IND.ENG. CHEM.,25, 1128-32 (1933).
Harte, C. R.. Jr.. Baker, E. M., and Purcell, H. H., Ibid., 25,
529
Killeffer, D. H., private communication. Lewis, G. N., and Randall, M., “Thermodynamics and the Free Energy of Chemical Substances,” New York, McGraw-Hill Book Go., 1923. McInnes, D. A,, and Belcher, D., J . Am. Chem. SOC., 55, 263046 (1933).
528-31 (1933).
Hatta, S., J. SOC.Chem. I n d . J a p a n , 37, Suppl. Binding, 275-7 (1934).
Hatta, S., Tech. Repts. Tohoku I m p . Univ., 10,613-29 (1932). Ibid., 10, 630-62 (1932). Hatta, S., and Baba, A., Ibid., 11, 365-82 (1934). Hatta, S., and Katori, M., J. SOC.Chem. I n d . , J a p a n , 37, Suppl. Binding, 280-2 (1934). Hitchcock, L. B., IND.ENG.CHEM.,26, 1158-67 (1934). Hitchcock, L. B., Trans. Am. Inst. Chem. Eagrs., 31, 347-64 (1935).
Hitchcock, L. B., and Cadot, H. M., IND.ENG. CHEM.,27, 728-32 (1935).
Hitchcock, L. B., and Mcllhenny, J. S., Ibid., 27, 461-6 (1935). Howe, H. E., Ibid., 20, 1091-4 (1928). Kiehl, S. J., and Loucks, R. D., Trans. Electrochem. SOC.,67 (preprint) (1935).
Menzel, H., 2. physik. Chem., 100, 276-315 (1922). Payne, J. W.. and Dodge, B. F., IND. ENG. C K ~ M 24. . , 630-7 (1932).
Riou, P., and Cartier, P., Compt. rend., 184, 325-6 (1927). Sieverts, A,, and Fritzsche, A., 2. anorg. allgem. Chem., 133, 1-16 (1924).
Ibid., 133, 17-25 (1924). Walker, A. C., Bray, U. B., and Johnston, J., J. Am. C h m . Soc., 49, 1235-56 (1927). Whitman, W. G., and Davis, G. H. B., IND.ENG. CHEM.,16,
264-6 (1924). (24) Williamson, R. V., and Mathews, J. H., Ibid., 16, 1157-61 (1924). RECEIVED January 18, 1937. This paper is based on a dissertation presented by Charles S. Comstock in June, 1935, to the faculty of the Graduate School of Yale University, in candidacy for the degree of doctor of philosophy
Bubble-Cap Column as a LiquidLiquid Contact Apparatus The operation of a small three-plate bubble column was studied in which the
M. C. ROGERS A N D E. W. THIELE Standard Oil Company (Indiana), Whiting, Ind.
interior could be observed when used to extract a heavy motor oil with dichloroethyl ether. The observed plate efficiency was quite low, not over 33 per cent at most. This is attributed to lack of agitation in this type of column.
continuous thread. No droplets were formed, and it was quite certain that the mixing with the solvent was poor. In order to improve the cap, a little “gable,” 0.25 inch wide, was soldered over each notch so that the liquid would be discharged away frpm the cap. This worked quite satisfactorily, and droplets were easily formed.
T
HE study of liquid-liquid extraction processes which has been made in recent years has brought into prominence the analogy between these processes and distillation. At present the apparatus used for the countercurrent extraction of liqujds with liquids usually consists either of a packed column or of a series of mixers and settlers; each combination of a mixer and a settler is the analog of a single bubble tray in distillation. Since the bubble tray is standard and efficient equipment in distillation, it is natural to think of adopting this type of apparatus to the extraction process. The work described here was undertaken with this idea in mind.
)PLATE GL
,TOP OF C A P PERFORATED
Apparatus The apparatus used is illustrated in Figure 1: It is essentially a small, rectangular, three-plate bubble column made of wood, with one rectangular cap per plate; the column is cut in two lengthwise, and the long side of the column is made of glass. By means of steel clamps, the plate-glass side is ’ held tight against rubber gaskets coated with shellac. The half-caps are made of galvanized sheet iron and are fastened t o the back wall. Three types are shown in Figure 2; in any one experiment all three caps were of the same design. The three types studied were V-notch, punched, and slotted. The V-notch type (Figure 2 A ) was first used with simple Vnotches cut into the metal. In operation, however, it was found that the liquid issuing from the top of the V-slots of this cap simply flowed up the metal of the cap and left at the top as a
EXTRACT O U T
u II
FIQWRE 1. LABORATORY BUBBLE-CAP EXTRACTION TOWR
