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Absorption Rates of Carbon Dioxide in Sodium Carbonate Solutions Effect of Impurities and of Surface Tension J. HAROLD SMITH AND ELTON L. QUINN University of Utah, Salt Lake City, Utah
H E absorption of carbon this fact seems to indicate that A study of the absorption of carbon dioxide dioxide from gaseous mixthe surface tension of the abin sodium carbonate solutions was made tures and its subsequent sorbent is more important in conwith various amounts of sodium sulfate trolling the rate of absorption liberation by certain alkaline added as an impurity. The absorption solutions is a process of great than the viscosity. The only coefficients were found to be greatly affected commercial importance. This direct suggestion of a connecis the basic reaction involved tion between absorption rate by this impurity, which is always present in the principal process used and surface tension is that of in commercial carbon dioxide plants using for making commercial liquid Killeffer (6). He made some sodium carbonate absorbing solutions. carbon dioxide and the industrial “sketchy tests”, prompted Absorption coefficients were also deterrefrigerant dry ice. The abdoubtless by the results of mined in solutions of sodium carbonate, sorbents generally employed are Riou and the other suggestions sodium carbonate, potassium mentioned above, which seemed the surface tension of which was changed t o indicate to him a marked Carbonate, and monoethanolwith the addition of various quantities of amine. The most common gas increase in the absorption rates detergents. mixtures treated are flue gas when surface tension depressors were added. from burning coke, natural gas, and in some cases combustion gases from coal. Any inAfter the experimental part of this work was finished, a formation on the mechanism of absorption seems worth while paper by Sherwood and Holloway (11) appeared. They disbecause it may aid in increasing the yield of the commercial cussed the experiments of Rennolds (9),who studied the process. The authors therefore started a study of the effect effect of wetting agents on the rate of desorption of carbon of sodium sulfate, which is always present as an impurity in dioxide from water solutions. The results obtained by Rencommercial sodium carbonate absorbing solutions, on the rate nolds seem to indicate an opposite effect from that obtained of absorption of carbon dioxide from a synthetic flue gas. It in this research. When the surface tension was decreased also seemed desirable to investigate the relation, if any, bewith formaldehyde and certain commercial wetting agents, tween the surface tension of the absorbent and its absorption the absorption coefficient also decreased. We have no exactivity. planation for this; even if the absorption systems were enThe literature concerned with absorption, absorbing agents, tirely different, one would hardly expect such a difference in and rates of absorption is too great to be reviewed here. results. That having to do with the change of rate of absorption by the addition of foreign substances will be considered briefly. It is Apparatus and Procedure recognized at present that absorption and the nature of the surThe apparatus used for measuring absorption rates is shown face layer molecules of the absorbing liquid are closely related. in Figure 1 and was, in general, much like that described by Therefore anything tending to alter the environment of these Hirst and Pinkel (3) : surface molecules would affect the absorption rate in some way, The viscosity of the absorbing solution has been suggested by The absorption tower, A , was constructed from a Pyrex lass several authors (4, 6,8)as a factor of considerable importance tube, 3.38 cm. in diameter and 107 cm. long. I t was packed for a in determining absorption rates. However, the earlier work length of 88 cm. with glass rings, 0.687 cm. 0. d 0.47 cm. i. d., and 0.69 cm. long (average dimensions). The en&e packing was of Riou and his associates (10) failed to show any direct resurrounded with a glass jacket throu h which water was passed lation between viscosity and absorption, under the condit o keep the temperature constant. $he total volume of packed tions they used. Their studies were made on alkaline carminus the volume occupied by the packing was 504ml. bonate solutions with the addition of such substances as he sodium carbonate absorbing solution was stored in the 5-liter bottle, B from which it flowed t o the ap aratus through the phenols, formaldehyde, sucrose, dextrose, etc., and many of constant-Lead device, C . The flow of the sofution was measured them definitely accelerated the rate of absorption of carbon by flowmeter D in which bromobenzene was used as the manomdioxide. Some of these substances which increased the abeter liquid. After passing through trap E, the absorbent was sorption rate are well-known surface tension depressors, and brought to the proper temperature in preheating tube F and con-
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I N D U S T R I A L A N D E N G I N E E R I N G CHEMISTRY
ducted into the top of the absorber. From the base of the tower the solution overflowed through a siphon which was adjusted to maintain a liquid seal. Samples for titration were taken from tube H .
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continued for about 2 hours, during which all adjustments were made to fulfill the standard conditions. Analysis of the solutions showed that this time was sufficient to bring the whole operation to a steady state. A sample of about 200 ml. was taken in a small flask, which was tightly stoppered immediately. At 10-15 minute intervals two other samples were taken, and alternately an 800-ml. sample was collected for the measurement of the surface tension. Exposure of the samples to the atmosphere was avoided as much as possible.
Analytical Methods
OF ABSORPTIONAPPARATUS FIGURE 1. DIAGRAM
The carbon dioxide-air mixture was produced by mixing metered streams of the individual gases; the air was taken from the compressed air system of the laboratory, and the carbon dioxide was drawn from a cylinder of the liquid through a reducing valve. Water-sealed gasometers were inserted in the lines to ensure constant gas pressures. Frequent analyses of this gas mixture showed it to be constant within a few tenths per cent. This gas mixture was passed up through the absorbing tower countercurrent to the carbonate solution. The resistance of the tower packing to gas flow was found to be slight at the gas velocities used. Temperature measurements were made with thermometers placed in the water jacket, at the bottom of the absorber, and at the top of the absorber. The temperature in the apparatus was controlled by manual adjustment of the hot and cold water from the laboratory taps. This system was found satisfactory, since with careful and continued adjustment the maximum variation of the temperature was not more than 0.2" from the required 25" C. During the runs the temperature of the laboratory was held close t o 25" C. While readings were being made, all of the various factors were held as nearly constant as possible, the absorbing solution alone being varied. The standard conditions adopted at the start of the work and maintained closely throughout all of the determinations may be summarized as follows: temperature at 25" C. (* 0.2" C.); rate of liquid flow, 2.86 liters per hour (318.7 cc. per sq. cm. tower area per hour); carbon dioxide in the gas mixture, 13.4 per cent; velocity of gas flow, 4.8 liters per minute. At the start of each determination the absorbing tower was first flooded with the solution to be used for absorption to ensure thorough wetting of the packing. This solution was then drawn down to the operating level, and the flow of the absorbent and gas mixtures carefully adjusted. The run was
CARBONATE-BICARBONATE CONCEKTRBTIOX OF T H E ABSORBING SOLUTION was determined by a modification of the Winkler method which seemed t o be satisfactory for this work ( 7 ) . However, the large amount of barium carbonate masked the end point during the titration so that the following system of analysis was devised and followed: To a 250-ml. volumetric flask containing a measured volume of standard barium hydroxide solution (more than the necessary amount to convert all the bicarbonate to carbonate), together with sufficient barium chloride to ensure complete precipitation of all the carbonate, was added a definite volume of the carbonate-bicarbonate solution. The volume of the solution was made up to the calibration mark with boiled distilled water; then the flask was vigorously shaken and allowed t o settle (usually by standing overnight). The supernatant liquid was siphoned off, and the excess base determined by back-titration of an aliquot portion of the clear solution with a standard acid and phenolphthalein as indicator. To minimize errors due to the volume of the solid precipitate in the solution, the amount of absorbing solution used was varied so as to keep this volume practically constant. Since we were interested in the relative concentrations of bicarbonate ion in our different solutions rather than in the absolute quantities, the actual volume of this precipitate was never determined. GAS ANALYSESwere made with a standard Burrell gas analysis apparatus which gave readings accurate to about 0.2 per cent of carbon dioxide. 25
35
SURFACE 45 55 TENSION65 DyNE?5
I SODIUM
PER
05
CM'
95
I
SSATE CONCENTRAT'PON,
I
I ~ a .PER
/
L.1.5
FIGURE 2. EFFECTOF SULFATE ION COSCENTRATION AND OF SVRFACETENSION ON ABSORPTIONOF CARBON DIOXIDE
SURFACE TENSION MEASUREMENTS were made with a du Koiiy tensiometer which was first carefully calibrated. I n an attempt to obtain reproducible results, the following procedure was followed: The platinum ring and the watch glasses used for the solutions were cleaned in the usual manner. The solution to be measured was shaken for a moment before being placed in the watch glass. The ring was dropped into the liquid, lowered to the bottom, pulled out of the solution, and brought back to just touch the surface for the measurement.
INDUSTRIAL AND ENGINEERING CHEMISTRY
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surface estimate, The value of V was 790 ml., as this was TABLE I. ABSORPTION RATEOF CARBONDIOXIDE IN SOLUTIONS the total space occupied by the tower packing. Calculation OF VAR1O F SODIUM CARBONATE (2.054 THE of the driving forces, P , involves a correction for the equilibous AMOUNTSOF SODIUM SULFATE" rium vapor pressure of carbon dioxide over the carbonateConcn., NalSO4 CO* 3 COXin %, NaHC03 Surface bicarbonate solutions, since the driving force is equal to the Run No. Equivalents/l. Absorbed, (Calcd,) xit Gas inSoln, Exit Koa partial pressure of the carbon dioxide in the gas phase minus the equilibrium pressure of the carbon dioxide over the soh7 0 3.557 tion involved. This equilibrium pressure was neglected in our 8 0 3.557 Av. 0 3.557 12.6 2.90 0.0405 77.8 use of the equation, since the equation of Harte, Baker, and 10 0.20 3.369 Purcell (2)showed it to be wholly negligible a t the low degree 0.20 3.350 11 of conversion obtained in these measurements. Av. 0.20 3.360 12.6 2.74 0.0383 79.0
Dyz7&,
12 13
1.00 1.00 1.00 1.60 1.50 1.50 1.50+ detergentb (1 g./l.)
Av. 14 15
Av. 16
2.614 2.580 2.597 2.076 2.110 2.093 3.049
12.8
2.15
0.0293 . 7 9 . 6
12.9 12.7
1.76 2.50
0.0236 0.0346
Conclusions
80.3 32.6
a Temperature of-runs 2 5 O C.; average barometric pressure, 650 mm.; rate of liquid flow, ~2.75'liters/m~nute/square meter, rate of gas flow, 4 . 8 liters/minut,e; COz in entering gas, 13.4%; enterin; solution 0 . 1 4 7 converted t o bicarbonate: effective tower volume, 790 mi.; estimate$ total surface 6000 sq. om. b Of 'the lauryl sulfate type.
Under the conditions of this research an increase of sulfate concentration in a solution of sodium carbonate slows up the rate of absorption of carbon dioxide. Moreover, in the range investigated with our apparatus, the decrease in this rate is a straight-line function O f the sulfate-ion concentration. In plant practice this mean a loss Of absorbing power due to an ever-increasing concentration of sulfate as well as to the decrease in the carbonate concentration as absorption takes place. This work has also shown that the rate of absorption is directly related to the surface tension of the absorbing solution. When the surface tension of the absorbing solution was decreased from 78 to 33 dynes per cm., the absorption rate was observed to increase nearly 30 per cent. Such a change
(It was found that this procedure, carrie,d out in exactly the same way each time, was necessary to obtain reproducible results. The degree of agitation of the solution just previous to the measurement has a considerable effect on I its surface tension.) Readings were repeated at the rate of approximately one per minute, TABLE 11. ABSORPTION RATEOF CARBON DIOXIDE I N SOLUTIONS O F S O D I o M and usually five to ten successive readings CARBON AT^ (2.054 N ) WITH ADDITIONS OF SURFACE TENSION DEPRESSORS~ were made before a constant value was reached. Bicarbonate Ten to twenty readings were then taken, and % POz in Conversion, % co2 Surface the average was calculated. This procedure Run Substance Added Exit Gas Entering Exit Absorbed, Tension NO. (Approx.) (Calod.) soln. soln. G./Hr. K g a Dynes/C&. was repeated twice more with fresh solutions 7-8 ......* .. 12.6 0.14 2.90 3.657 0 . 0 4 0 5 77.8 and properly cleaned apparatus, and the surface 17 Detergentb (0.006 d l . ) 12.6 0.14 3.03 3.725 0.0424 45.9 18 Detergentb (O.O15g./l.) 12.5 0.14 3.18 3.926 0.0449 40.9 tension reported is the average Of the 19 Detergentb (0.024 g,/l,) 12.4 0.14 3.61 4.480 0.0514 33.2 3.72 4 . 5 9 3 0.0527 of these three series of determinations. 20 Detergentb (0.44 g./U 12.4 0.16 33.0 21 22
Calculation of Abso+ion Coefficients Table I shows the effect of the addition of
Oleic acid (1 drop/l.) Soap0 (satd. s o h )
12.5 12.6
0.16 0.14
3.38 2.96
4.154 3.638
0.0475 0.0414
36.6 49.3
a Temperature, 25' C . ; average pressure 650 mm: rate of liquid flow 52 75 liters/min~ entering gas: 13:4%; effective ute/square meter; rate of gas flow, 4 . 8 literdminute: ' C O in tower volume 790 ml ' estimated total surface, 6000 sq. om. b A synthetic detergknt of the lauryl sulfate type. A popular white floating soap.
sodium sulfate to the absorbing solution, and I Table I1 gives the effect of change in surface tension o n the rate of absorption. Figure 2A indicates that the rate of carbon dioxide absorption decrease is a linear function of the sulfate ion concentration. Figure 2B shows the relation between the absorption rate and the surface tension of our solutions. The absorption coefficient Koa (Tables I and 11) was calculated by the equation of Comstock and Dodge ( 1 ):
of surface tension can be accomplished by the addition of small amounts of ordinary surface tension depressors, and unless some deleterious effect such as excess frothing takes place, it is possible that the operating efficiency of a plant lye solution might be appreciably increased by this procedure.
Literature Cited where K,,
= absorption coefficient = effeotlve interfacial area per unit of tower volume W / 0 = weight of COZabsorbed er unit time pa". = log mean of terminal giving forces in pressure units
a
I n general, K g and a are not separated, and the product "is also called an absorption coefficient. Values of Kua are expressed as grams/(hour) (ml.) (atmosphere). The effective interfacial area, a, was estimated on the basis of the number of rings in the tower; their average dimensions were determined by measuring thirty rings taken at random. The inner surface area of the tower was included in the total
(1) Comstock, C. S.,and Dodge, B. F., IND. ENG.C H ~ M29, . , 520 (1937). (2) Hart%C. R., Baker, E. M., and Purcell, H. H., Ibid., 25, 528-31 (1933). (3) Hirst, L. L.,and Pinkel, I. I., Ibid., 28,1313 (1936). (4) Hitchcock, L.B., Ibid., 29,302-8 (1937). (5) Hitchcock, L.B.,and Cadot, H. M., Ibid., 27,728-32 (1935). (6) Killeffer, D. H., Ibid.,29, 1293 (1937).
iii ~ (1932). ~
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(9) Rennolds, P.J., Mass. Inst. Tech., S. M. Thesis in Chem. Eng., 1939. (10) Riou and co-workers, Compt. rend., 174, 1017, 1463 (1922); 184, 325 (1927); 186,1543, 1727 (1928). (11) Sherwood, T. K., and Holloway, F. A. L., Trans. Am. Inst. Chem. ETZQTS., 36, 21-70 (1940).
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