December, 1924
INDUSTRIAL A N D ENGINEERING CHEMISTRY
pressure difference in the system caused by the absorption and indicates only the rate a t which this pressure difference causes the liquid to move in the capillary. The true rate of absorption must be considerably greater than the apparent rate shown by these curves, because the low pressures produced as absorption begins must seriously slow down the rate of absorption during the remaining period. However, a study of this series of curves and of the vapor pressure data on ammonia solutions brings out the interesting fact that the slope of the curves is proportional to the total gas pressure (1 atmosphere) minus the vapor pressure of ammonia above the solution.
1233
Comparative Absorption Rates for Various Gases By W. G . Whitman and D. S. Davis MASSACHUSETTS INSTITUTE
OF
TECHNOLOGY, C A M B R I D G E , MASS.
HE work described in this paper represents an experimental study of absorption through a free surface of liquid when the liquid is stirred. To determine the relationships between absorption rates for different gases, four gases of widely varying solubilities were selected for investigation. These gases, oxygen, sulfur dioxide, ammonia, and hydrogen chloride, represent a range of about three hundred thousand fold in solubility under the conditions of the experiments. I n addition to the absorption runs, experiments on the rate of escape of sulfur dioxide from aqueous solutions by air were performed, including runs on the effect of rate of stirring and viscosity of solution.
T
EXPERIMENTAL METHOD
BuabLE L&LUM€ F I G . 8--EFFEC’I
OF C O N C E N T R A T I O N U P O N
(CC)
The apparatus is shown diagrammatically in Fig. 1. The absorption chamber, A , is an 8-liter porcelain jar with side outlets for gas entrance and exit tubes, a and b, for a glass coil for cooling water, c, and for a tube for withdrawing liquor samples, d, and holding the liquor thermometer. A plate glass cover with rubber gasket is clamped over the jar, and the paddle stirrer, e, for gas and liquid passes through a stuffing box set in this plate. This stirrer is rotated a t 60 * 5 r. p. m. in all runs except when otherwise noted. For ammonia, sulfur dioxide, and oxygen the gas was bled from commercial cylinders, whereas the hydrogen chloride was generated by dropping sulfuric acid into concentrated hydrochloric acid solution. This gas stream, after passing through flowmeter B, mixes in chamber C with a stream of air which has passed through a pulsation chamber, D,and a flowmeter, E. The temperature of the mixed gas is taken at a.
RATEO F ABSORPTION OF
AMMONIA
.......... DISCUSSION
Professor McAdams asked whether the higher rates with potassium hydroxide as compared with sodium hydroxide could not be explained by the lower solubility of sodium bicarbonate. Mr. Ledig stated that this might be the case, particularly as fine crystals were formed in certain of the experiments with sodium hydroxide. Mr. Wilson mentioned that carbon dioxide was very difficult to absorb from dilute gases, as he had observed in studying the absorption of various gases in gas mask research. The chairman asked whether the viscosity of the absorbing solution had been considered; furthermore, had the ammonia data been investigated with reference to the rate varying with the degree of unsaturation? Mr. Ledig replied that the data had been obtained very recently and had not yet been subjected to a thorough analysis. Milwaukee Tests Rubber Paving A strip of rubber paving for test purposes has recently been laid on the approaches to the Sixth Street Viaduct, Milwaukee, Wis. The rubber blocks employed were similar to those laid some time ago on the Northern Avenue Bridge in Boston. Milwaukee is the fourth large city to test the rubber blocks, similar strips having been laid in Boston, St. Louis, and Chicago.
u
u I1I
FIG.1
The procedure in an absorption run was to determine by liquor analyses the rate of absorption in 4 liters of water when a steady stream of gas of constant composition was passed over the liquid. In the cases of hydrogen chloride, ammonia, and sulfur dioxide, a 4-gram liquor sample was withdrawn a t intervals; for oxygen the sample was 250 cc. All samples were stoppered immediately after withdrawal. Hydrochloric acid solutions were analyzed with standard alkali and phenolphthalein, ammonia with standard acid and methyl orange, sulfur dioxide solutions with standard iodine and starch, and oxygen by the standard WinMer method. The procedure in runs to determine rate of escape of sulfur
