1220
INDUSTRIAL A N D ENGINEERING CHEMISTRY
be given in a later paper, it is shown that the predicted absorption rates check the experimental rates within 15 per cent for all three gases. It is of interest to note that the ratio of liquid to gas film coefficients
2 is roughly about 15-30 to 1 in the case of gas
bubbles (estimated from Perman’s data) ; whereas the ratio
is only about 1to 1 for gas passed over a liquid surface (using Davis’ data). This difference can probably be attributed more to the exceedingly thin liquid films obtained around gas bubbles than to any marked change in gas film thickness.
CONCLUSIONS Rate of absorption is controlled by rate of diffusion of solute through the surfade films of gas and liquid a t the gasliquid boundary. The relative importance of the two films is determined primarily by the gas solubility and to a lesser degree by the conditions of operation. Gas bubbles give very thin liquid films and are therefore particularly adapted to the absorption of the less soluble ga$es. On the other hand, the rates for the more soluble gases are limited by slow diffusion through the gas, and this effect becomes increasingly evident with gases above sulfur dioxide in solubility. Absorption through free liquid surfaces is enormously increased by stirring the liquid. The same general considerations hold as for gas bubbles, but the liquid film is thicker relative to the gas flm. For this reason the effect of the gas film is less noticeable with sulfur dioxide, although with more soluble gases it comes into prominence.
BIBLIOGRAPHY ‘1-Whitman, Chem. Met. Eng., 29, 146 (1923). %Bohr, A n n . Phys. Chem., 68, 500 (1899). 3--lewis, J . Ind. Eng. Chem.. 8, 825 (1916). 4-Donnan and Masson, J SOC.Chem. I n d . , 89,236 (1920). 5--Haslam, Hershey, and Kean, I n d . Eng. Chem., 16, 1224 (1924). 6--lewis, Much. Eng.,44,445 (1922). 7-Adeneyand Becker, Phil. Mag.,88,317 (1919); S9,385 (1920). 8-Ledigand Weaver, J . A m . Chem. SOC.,46, 650 (1924). g-Perman, J . Chem. SOC.( L o n d o n ) , 67,868 (1895). 10-Adeneyand Becker, Phil. Mag., 42,87 (1921). 11-Becker, I b i d . , 46, 581 (1923). 12-Whitman and Davis, I n d . Eng. Chem., 16, 1233 (1924).
.......... DISCUSSION
In the discussion following this paper, R. E. Wilson suggested that the general absorption equation could be expressed using partial pressures multiplied by the solubility coefficient in place of liquid concentrations for diffusion through the liquid film. Dr. Lewis pointed out that, although this was quite permissible in those cases where the solubility followed Henry’s law, it did not really simplify the problem and could not be used when the solubility relationship was not linear. The chairman noted that this method of presentation waq employed in a later paper (by Professor Haslam) and that the choice between the two systems of expressing liquid concentration would be primarily determined by convenience in using the equations. Dr. Becker communicated $he following discussion: It should be pointed out that the film theory can apply only when the liquid is mixed and in fact that it will have its most precise application when the conditions on both sides of the interface are distinctly turbulent. When the gas and the liquid are not kept in motion other factors appear to become of greater importance and the values of the absorption rates are not then always proportional to the solubility of the gas. The conclusion that for the absorption of less soluble gases a bubble rising through the liquid is the most efficient device is t)orne out by my experiments on stirring. When the liquid was
Vol. 16, No. 12
stirred as rapidly as possible without causing splashing, the value of the rate of solution was only about half the value obtained when the absorption took place from a bubble rising through the liquid. A further important point is the form of the bubble, as this determines the area i t exposes for a given volume, and also the rapidity with which the film surroundingit is renewed as it rises in the liquid. From previous experiments i t has been found that the most effective form is that of a cylindrical bubble ascending a fairly narrow tube, as this gives the maximum of mixing of the liquid for a given speed of rise. The remarkable effect of very slow stirring of the liquid in reducing the effective thickness of the film is of great importance when circumstances are such that a relatively insoluble gas has to be absorbed without bubbling the gas through the liquid, as by this means it is possible to obtain a very large increase in the rate of solution with the expenditure of very little power. Thus a stirrer running a t only 60 r. p. m. increases the value of k~ for oxygen about twenty-five times, as compared with stationary water.
Mechanism of Absorption of Moderately Soluble Gases in Water . By H.G . Becker STANDARD OIL
COMPANY O F NEW JERSEY,
ELIZABETH,
16.J.
WHENTHE LIQUIDIs MIXED
vv
HEN a liquid and a moderately soluble gas are allowed to come in contact and the liquid is kept thoroughly mixed so that it may be regarded as of
uniform composition a t all points, the rate of solution of the gas varies directly as the degree of unsaturation of the liquid. This has been shown experimentally by following the course of the absorption of oxygen and nitrogen in air-free water by a method which allowed the liquid surface to remain unbroken although the liquid was vigorously mixed.’ The method used consisted in allowing a large, cylindrical bubble of the gas to pass up through a tube full of water which was initially airfree, and measuring the reduction in pressure which resulted in the bubble, due to the absorption of a portion of the gas. After each traverse of the bubble, the pressure could be restored to its original value by connecting the gas space with an air reservoir by means of a special tap. The process of absorption could thus be followed step by step from zero concentration of dissolved gas to saturation. The curves shown in Fig. 1 which are for oxygen, show the regular way in which the percentage saturation of the liquid with the gas increases with time, and it is found that when the rate of solution a t various points on the curve is plotted against the corresponding gas content of the liquid a straight line results. The rate of solution can therefore be expressed in the form
-- a dw dt
-
Integrating, w
- bw
- a / b = ce-bt
For the experiments with air-free water, Whent = 0 , w = 0 , c = - a / b Therefore w = a / b (1 - e-b8) and when t = infinity, a / b = w = w8,the saturation value Hence, w = w, (1 e-bt) where w = weight of gas per cubic centimeter of liquid a t time, t w. = weight of gas per cubic centimeter of liquid a t equilibrium
-
When the initial gas content wt is not zero, this expression becomes 1
Adeney and Becker, Phil. Mag., 88, 317 (191s); 89, 385 (1920).
IND UXTRIAL AND ENGINEERING CHEMISTRY
December, 1924 w = w,
-
(w,
- w1)e-b'
The experimental constant a represents the rate a t which the molecules of gas are entering the liquid continuously. Since this will be proportional to the partial pressure of the gas p (provided the latter obeys Henry's law, which is approximately the case for gases of moderate solubility) and will be dependent on the ratio of area A to volume V , we may 100
80
60
40 20 0
0
4
2
6 FIG.1
write a = SPA, where S may be called the "entrance coefficient." Similarly the constant b will be of the form b = where $ is the exit coefficient, since it represents the rate a t which the molecules of gas are leaving the liquid continuously. Experiment shows that for the usual range of atmospheric temperatures the value of S is little affected, as might be expected, since it must be primarily dependent on molecular velocity. On the other hand, the value of f varies very greatly with temperature, increasing considerably for a moderate rise in temperature. This disparity in the effect of temperature on the rates of entrance and exit of the gas is reflected in the diminishing solubility of most gases in water with increase in temperature. The values of these constants f'or oxygen and nitrogen and air are given in Table I for a range of temperature from 2.5"to 35" C. yj
TABLB I
Temperature O
c.
Value o f f
Value of s
s
=fs
Oxygen
2.5 8.8 15.5 20.2 25.2 30.3 35.1
0.373 0.434 0.499 0.545 0.591 0.641 0.687
3.5 11.2 15.0 20.2 24.2 30.4 35.1
0.372 0.448 0.490 0.543 0.593 0.647 0.696
3.6 11.4 15.0 20.2 25.0 29.6 34.3
0.352 0.441 0.476 0.525 0.574 0.623 0.672
0.04390 0.03719 0.03206 0.02955 0.02732 0.02465 0.02270
0.0164 0.0161 0.0160 0.0161 0.0161 0.0158 0.0156
0.02203 0.01820 0.01701 0.01549 0.01456 0.01322 0.01220
0.0082 0.0081 0.0083 0.0084 0.0086 0.0085 0.0085
0.02700 0.02260 0.02120 0.01930 0.01780 0.01660 0.01560
0.0095 0.0099 0.0100 0.0101 0.0102 0.0103 0.0104
1221
librium in a gas-liquid system is very greatly accelerated, not only because the solubility is lowered, but also because the rate of exchange is greatly increased. It is important to note that the values off are practically identical for the three gases a t any definite temperature. If the liquid is not kept mixed but is allowed to come in contact with the gas in a stationary condition, the simplicity of the foregoing relation may be disturbed, since the liquid may no longer be considered of uniform composition throughout. However, between the extremes of stationary and perfectly mixed liquid there may be an indefinite number of stages and with a view to studying the relation between the rate of solution and the mixing of the liquid further experiments were carried out.2 These were performed in an apparatus in which the water cduld be stirred a t different rates by a stirrer introduced through the bottom of the vessel, while the absorption of gas could be measured by the constant-pressure manometer connected to the upper end of the vessel. Observations were made on the absorption of oxygen from air by de-oxygenating the water by means of a little ferrous hydroxide precipitated in the liquid. As long as any ferrous hydroxide remained the liquid could be regarded as oxygenfree, but as soon as it was all oxidized the liquid would start to re-oxygenate by the usual process of absorption. The method thus allowed a study of the rate of solution of oxygen into water of zero oxygen content, and the results show that under such conditions the absorption of oxygen is proportional to the time-i. e., the rate of solution is constant. When the oxygen starts to accumulate in the water, however, the rate of solution again folldws the logarithmic curveshowingthatthe rate is proportional to the degree of unsaturation. Table I1 gives the values of the absorption rates obtained with different speeds of stirring. Speed of stirrer R. p, m.
1000 590 340 140 80 60 0
Estimated area of surface Sq.cm.
TABLE I1
19.0 11.1 8.8 8.1 8.1 8.1 8.1
Rate of solution Cc./min.
Rate per unit area Cc./min.
0.030 0.014 0,008 0.004 0.003 0.002 0.0001
0.0016 0.0013 0.0009 0.0005 0.0004 0.00025 0.00001
When the rates of solution are plotted against the rate of stirring a straight line (11) in Fig. 2 is obtained, but when the S C ~ L - B /a0
200
so0
400
5ou
600
Nilronen
A.ir .
When the values o f f are plotted against temperature, a straight, line is obtained for each gas, and the equationsof these lines are f = 0.0096 ( T
f f
= 0.0103 = 0.0099
- 237) for oxygen
( T - 240) for nitrogen ( T - 239) for air
1
where T represents the absolute temperature in each case
It will be observed that the value off nearly doubles for a rise of temperature of about 35" C., and this emphasizes the fact that a t the higher temperatures the approach to equi-
Fro. 2
effect of the variation in area produced by stirring is taken into account and the rates per unit area are plotted, a curve (I in figure) results; which shows that with increased stirring the rate does not increase indefinitely but tends to reach a maximum. The rate previously obtained for oxygen was 0.0161 cc. per minute per square centimeter; therefore, for 'the oxygen in the atmosphere the rate should be 0.0161 X a Becker, Phil. M G K . , 46, 581 (1923).
IND L'STRIAL AND ENGINEERING CHEMISTRY
1222
21/100 = 0.0034 cc. per minute, since the rate of solution is proportional to the partial pressure. This value is more than double that obtained with the highest rate of stirring used in these experiments, thus showing that to secure a maximum rate of solution extremely efficient mixing is essential. TABLE 111
Mean value of S Value o f f Gramdsa. cm./hour 13.8 X 1 0 - 4 Oxygen 0.575 (T-236.5) Nitrogen 0.618 (2'-240.0) 6 . 3 X 10-4 Air 0.594 (T-239.3) 7.7 x 10-4 Value of S for oxygen for different rates of stirving Rate of stirring, r. p . m. 0 60 80 140 340 590 1000 S , grams/sq. cm./ hour ( X 10-4) 0 , 0 4 1 1.02 1.63 2 . 2 0 3.67 5.30 6.52
On the other hand, the unusual efficiency of very moderate stirring in hastening the rate of solution in otherwise stationary water is shown by the large increase in rate produced by revolving the stirrer as slowly as 60 r. p. m. This produces an inMJonmw nmnnrw crease of the rate to c,
value when the water is still. The results of these e x D e r i m e n t s are expressed using units of cubic centimeters and minutes, but it is sometimes more convenient to express them in terms of grams of gas and hours. Table 111 shows the constants in these units. R'HEN
THE LIQUID I S
NOTMIXED If an attempt made to study the rate of solution of air, for instance, by exposing vessels full of air-free water to the atmosphere for different times, and analyzFIG 3 ing them for gas content, it is found that the values obtained are liable to be very irregular in character. When the surface of the liquid is exposed to the atmosphere these discrepancies may be attributed to such agencies as drafts of air disturbing the liquid or cooling the surface by evaporation; but in one experiment made with stationary water in the apparatus described above it was noticed that the rate of solution altered abruptly in the course of the observations for no apparent reason. A series of experiments was therefore made3 using a modified form of the earlier apparatus, in which a column of air-free water was exposed to oxygen in such a way that the disturbing effects of temperature variations or evaporation from the surface of the liquid were eliminated. The whole apparatus was kept at a constant temperature by means of the water jacket through which water from a thermostat was circulated, and the gas was given the correct vapor pressure before being introduced so that no evaporation from the liquid could take place. The results showed that while the gas content of the water was low, the rate of solution followed a fairly regular course agreeing approximately with a logarithmic curve, but that when the content of the water reached 60 to 70 per cent of saturation the rate tended to become quite 8
Becker and Pearson, Proc R o y . SOC.(Dublin), 17, 197 (1923).
Vol. 16, No. 12
irregular, and beyond this point consistent observations could not be obtained. The values off obtained in this way lay between 0.386 and 0.600, and these values compare fairly well with some approximate values previously obtained from some results of experiments with exposed tubes of water. The values in the latter case were f = 0.388 for fresh water and f = 0.509 for sea water, and it will be seen that the values are of the same order of magnitude in both cases. It may therefore be concluded that during the earlier states of reaeration of stationary water the process of absorption proceeds as if the water were very slowly mixed, and the rate of solution varies accordingly. Under these conditions the value of f may vary somewhat owing to accidental disturbances. With the object of determining to what extent this irregular absorption generally occurred in the absorption of a gas by a liquid, a series of experiments was begun in which observations were made on the behavior of gases of widely different solubilities when exposed to air-free water. The apparatus used is shown in Fig. 3 and is a modification of somewhat similar apparatus used by Wroblewski.* It consists of a glass tube, A , of about 25-mm. bore, fitted at the lower end with a rubber stopper and a wide bore stopcock, B. The upper end of the tube is ground off flat and cemented into a hole drilled in a piece of plate glass, CI, which is ground t o fit a second similar sheet of glass, Cz, so that the two slide over each other air-tight when slightly greased. The upper plate, CZ,carries a short length of the same size tubing which serves as a gas space; this is fitted with a rubber stopper carrying an inlet tube, D, and an outlet tube, E, as well as a manometer, M , with which tomeasure the decrease in pressure. The whole apparatus was immersed in a thermostat during the course of an experiment. The tube A and tap B were filled with a clean mercury, and the glass plate CZslid on in the position shown in the figure. The vessel containing the air-free water (previously prepared) was then connected to B so as to leave the system full of water and mercury only. On opening the tap, B, the mercury slowly flowed into the lower vessel, and the air-free water replaced it in A without being exposed to the air. The apparatus was then transferred to the thermostat and the upper chamber filled with the gas to be used, by passing a stream of gas through until a test showed that all the air had been driven out. The taps D and E were then closed, F opened, and the top plate slid over until the two tubes were in line, the time being noted with a stopwatch. The gas was in this way suddenly exposed to air-free water, and the course of the absorption could be observed on the manometer from the very start. The results obtained with a few gases are shown in Fig. 4. It will be seen that there is a wide variation in the behavior of different gases when being absorbed by water. Under certain conditions some gases are absorbed in such a way as to form a saturated layer a t the surface of the liquid, which causes the rateof solutionto falloff very rapidly, but thislayeris frequently very unstable and any slight disturbance causes it to sink in the liquid, thus causing a corresponding increase in therate of solution. Thus in the case of carbon dioxide the initial rate of solution is high, but this rapidly falls off until it becomes nearly zero. On disturbing the apparatus slightly, however, this rate immediately regains its high value, which again falls off; and it is interesting to note that this falling off in rate does not occur a t the same degree of saturation in all cases and in some cases does not occur a t all (especially if the apparatus is so supported as to be subject to slight vibration). Thus in Experiments I and I I a with carbon dioxide the absorption continued a t a rapid rate far above the concentration which in other experiments caused a flattening in the curve. 4
A n n . P h y s i k , 4, 268 (1878).
December, 1924
INDUSTRIAL A N D ENGINEERING CHEMISTRY
I n the case of Experiment I with carbon dioxide it will be seen that when about 0.37 cc. was absorbed the rate tended to diminish, but a slight shake restored it to nearly its previous value and it finally fell off to 0.63 cc. The apparatus was again disturbed, and the gas started to enter the liquid at a rapid rate and continued for some time with but slight falling off. Again with Experiment I1 the rate showed a tendency to fall to zero a t a somewhat earlier stage, while the disturbance had but slight effect and the rate fell off again rapidly. Asimilar effect is noticeable with hydrogen, when the rate fell tozeroat about 0.5 cc. absorption, but was again accelerated by a fairly vigorous shaking. The effect was very marked with hydrogen sulfide, which started a t a very rapid rate as might be expected from its solubility, but this rate fell off and became zero a t about 1.4 cc. absorption. In this case, although vigorous shaking caused a slight increase in the rate, still no appreciable amount of gas beyond this point could be introduced into the liquid. The surface layer in this case seems to be particularly stable and resistant. In the case of nitrous oxide three experiments showed practically the same rate of solution up to 0.5 cc. absorption. Beyond this there are slight deviations but the rate is fairly well maintained throughout. In the last experiment there is evidence of the formation of a saturated layer in the rather sharp falling off shown a t T , which was stopped by a slight shake a t this point. The rate then resumed its former value. With nitric oxide and water the rate is high and is continued throughout in Experiment I, but in Experiment I1 the rate
1223
When chlorine is exposed to water it is very rapidly absorbed, and the rate of solution maintains its high value throughout the period of these experiments without any evidence of the formation of a saturated layer. The amounts absorbed in this case are so large that the scale in the figure is ten times its indicated value for the chlorine graphs. The initial rates of solution of still more soluble gases were observed with this apparatus, but as the time interval over which the absorption could be measured was very small, the figures obtained can only be regarded as approximate. Table I V gives the values of the initial rate of absorption for the different gases examined (at 22’ C.)-that is, the rate a t which they enter when the water is free from gas of any kind and before any saturated layer has formed. (By “saturated layer” is meant a film of liquid of measurable gas content which should not be confused with the saturated surface film of molecular thickness which is probably always present.) TABLE IV GAS
Hydrogen Nitrous oxide Carbon dioxide Nitric oxide in water Nitric oxide in ferrous sulfate Chlorine Hydrogen sulfide Carbon dioxide in caustic potash Hydrogen sulfide in caustic potash Sulfur dioxide in water Hydrogen chloride in water
Initial rate of solution Cc./min./sq. cm. 0.0043 0.0120 0.0168 0.0375 0 I0250 0.0984 0.0950 2.50 5.56 2.00 54.30
Grams/hour/ sq. cm. 0.000023 0.000141 0.000198 0.000302 0.000201 0,001870 0.000866 0.0295 0.0506 0.0343 5.3
I ’9
80
100
/i
FIG.4
fell off rapidly to a very small value owing apparently to the formation of a very resistant layer. With nitric oxide and a solution of ferrous sulfate the rate is somewhat lower than with water (Experiment 111).
These experiments show that the absorption of a gas by stationary water is a complex phenomenon and may occur in different ways according to the nature of the gas and its effect on the density, surface tension, and viscosity of the
INDUSTRIAL A N D ENGINEERING CHEMISTRY
1224
water. Thus certain gases, such as carbon dioxide, hydrogen, and hydrogen sulfide, tend to saturate the surface layers of the water fairly readily, and the layers so formed are of considerable permanency and not easily dislodged. On the other hand, gases such as nitrous oxide, nitric oxide, and chlorine show no tendency to form such saturated surface layers and are absorbed at a correspondingly higher rate. Since in most cases the effect of dissolved gas on the density of the water is to increase it, we may form a picture of the process by assuming that the first action is to form a saturated layer which is heavier than the rest of the liquid and tends to sink. At the same time the surface tension and the viscosity of the liquid tend to maintain the surface in its original condition. Whether the layer will sink and a t what rate is then determined by the relative magnitude of these opposing agencies, which also determine whether the layer will be stable or unstable-i. e., whether the rate of solution will be greatly affected by slight disturbances of the liquid or not. The effect of dissolved gas in the surface tension and viscosity of the liquid is of importance even when the liquid is used in a scrubbing tower, as it has been pointed out by Hurtersand also by Donnan6 that these properties may affect the effective interface of the gas and liquid owing to their effect on the formation of drops, splashing, and the wetting of the packing material of the tower. The matter is therefore of considerable importance and worthy of further study.
Vol. 16, No. 12
with single bubbles of known area, and that the chief difficulty in estimating absorption rates for this case would come in determining the actual area of the many bubbles formed by such means. Mr. Wilson remarked that experiments on absorption from bubbles by oil indicated that the type of nozzle used to deliver the bubbles had very little effect upon the rate of absorption, since the bubbles aggregated to a definite size a t a point just above the nozzle. I n this connection T. A. Boyd stated that the type of nozzle was very important in experiments where gas was bubbled through water. He had noted a certain amount of aggregation of bubbles through water, but in dilute sodium chloride solution there was no such effect. W. H. McAdams noted the similarity between Dr. Becker’s curve showing the effect of velocity on the rate of absorption coefficient and the curves obtained in studies on the effect of velocity on the coefficient of heat transfer. This point further emphasized the parallelism between the two processes of absorption and heat transfer.
Effect of Gas Velocity and Temperature on Rate of Absorption’ By R. T. Haslam, R. L. Hershey, and R. H. Keen
..........
MASSACHUSETTS INSTITUTE OR TGCHNOLOGY, CAMBRIDGE, MASS.
DISCUSSION
I n discussing the first halfof Dr. Becker’s paper, R. P. Russell pointed out the similarity between Dr. Becker’s results with stirred solutions and the theory presented by Dr. Lewis. Since the atmospheric gases are only slightly soluble, the absorption of these gases is determined by liquid diffusion. Although Dr. Becker had developed his equation with a different nomenclature and had finally expressed it in integrated form, the basic principle was the same in both treatments. In fact, the coefficientf,when multiplied by 60, was the same as Dr. Lewis’ coefficient kL, the factor 60 being necessary to convert from minutes to hours. The results obtained when the liquid was free from agitation could not be interpreted so readily, since it was evident that stratification effects were pronounced. Such results could not be subject to the formula proposed by Dr. Lewis, which assumed uniformity of composition in the main bodies both of the liquid and of the gas. However, it was possible that the initial absorption rates tabulated by Dr. Becker would conform to the theory, provided the influence of stratification was unimportant a t the start. Mr. Russell had therefore calculated absorption coefficients from these initial rates, assuming that the rate was controlled entirely by liquid film diffusion. He believed that it was permissible to neglect any gas film because Dr. Becker worked with pure gases. These calculations gave values of k L varying from 1.1 to 15-4. e., a 1Cfold range for the pure gases in water, whereas the initial rates themselves varied over 200,000 fold on a weight basis. Admitting that the agreement of the coefficients was only qualitative, the results, nevertheless, seemed to check reasonably with his assumptions. Mr. Russell asked if Dr. Becker believed that stratification was evident a t the start, which might explain the discrepancies in the coefficients, and Dr. Becker replied that it probably did affect the initial rate. Mr. Melligan asked whether Dr. Becker had any data on absorption rates when air was bubbled up through a porous plate. Dr. Becker replied that his work had been carried out
T
HE research herein reported was conducted in an effort
to discover whether the two-film theory of absorption of a gas in a liquid would quantitatively describe the rate of absorption in a given system under varying conditions. Accordingly, it was proposed to operate a definite systemthat is, a system of a given gas absorbed in a given liquid in a standard apparatus-first, under conditions such that the liquid film offered the major resistance; and second, under conditions such that the gas film offered the major resistance. Furthermore, two such systems, one having very soluble solute gas and the other having a less soluble solute gas, were to be employed. Thus the effect of solute gases of widely different properties was to be studied. The similarity between the rate of transfer of heat by conduction and the rate of transfer of material by diffusion has already been indicated. This investigation has as the second object the further extension of this relationship.
EXPERIMENTAL METHOD APPARATUS-The apparatus used is shown diagrammatically in Fig. 1. The tower was a Pyrex- glass tube 3 inches in diameter and 3 feet long, lagged with magnesia pipe covering. The liquid was introduced into the tower through a distributor a t the top. The tower was of the wetted wall type, without packing. The gas was blown up through the tower. T h e solute gases were drawn from cylinders; the air, which was used as a “carrier” gas, was forced through the tower by means of a pressure blower. Flowmeters and an orifice were used to measure the gas. The liquid, which was water in all cases, was heated where necessary by blowing steam into it. The final adjustment of water temperature was made by leading the water through a coil in a constant temperature water bath. The gas temperature was adjusted by leading the gas through a coil in the same bath. Thermometers were inserted in the water and gas lines both before and after the tower. Presented under the title “A Study of the Mechanism of Absorption.” A. E. Marshall, Consultiag Engineer, Corning Glass Company. 1
6
J. SOC.Chem. I n d . , 4, 641 (1885). Ibid , 39, 236T (1920).
* Obtained through courtesy of
