1698
MARTINBLANK
then appears to relax and return to the structure that il; had previously a t this temperature. Upon deuteration the protein seems to be structurally more stable because it shows much smaller changes in adsorptive capacity than did the hydrogen prok i n when the temperature was raised and then
Vol. 65
lowered. Lysozyme also shows that mono site coverage by water molecules as proved by deuterium exchange may be determined by extrapolating straight line portions of the isotherms out to the abscissa for the saturated vapor pressure a t a given temperature.
THE EFFECT OF VAPORS ON MONOLAYER PERMEABILITY TO CARBON DIOXIDE* BY MARTINBLANK Department of Physiology, College of Physicians and Surgeons, Columbia University, New York, N . Y . Received March 7. 1961
Changes in the vapor pressure of water or methanol, which are present in the subphase, greatly affect the permeability of an or:tadecanol monolayer to carbon dioxide. Water vapor decreases the permeability while methanol vapor increases the permeability a t low concentrations and decreases it a t high concentrations. Both vapors are present in the monolayer as additional resistances to the passage of carbon dioxide but it is not known if the vapor is trapped or if the carbon dioxidevapor collisions are more effective in a monolayer. The permeability of octadecanol t o carbon dioxide appears to vary linearly with water vapor pressure so it is possible to extrapolate the values to a zero vapor pressure state. This result is compared with octadecanol permeability to water and the ratio of the permeabilities agrees a p roximately with a Boltzmann expression where the energy is the work that must be done by the gas to pass through t i e monolayer. Since this energy depends on the diameter of the penetrating molecule. the effect of molecular size appears to be an important factor in the energy barrier to penetration.
I. introduction cules and a decrease in the cohesion of the moncRecent studies',? on monolayer permeability to layer, bot,h resulting in a sharp increase in perseveral gases have indicated that water vapor is meability. (This is observed in the case of water operationally present in the monolayer as an addi- vapor permeability.) However: the temperature tional resistance. In brief, the supporting evi- rise also causes a rise in the vapor pressure of the water, increasing the amount of water operationdence is the following. ally present in the monolayer during the penetra(a) The permeabilities of several monolayers to carbon dioxide, oxygen and nitrous oxide are tion of carbon dioxide and nitrous oxide. The approximately two orders of magnitude lower than net result, a smaller change in permeability, has a the previously observed permeabiiity to water magnitude that can be predicted2 semi-quant,iThese two measurements, however, tatively on the basis of the above assumptions. (c) The permeability of monolayers on very condo not relate to monolayers under the same condicentrated solutions such as glycine buffers, hemotions. When one measures the permeability of a monolayer to water, the system includes a mono- globin solutions, etc., are much greater than on the layer on an aqueous subphase. In the experi- more dilute solutions. This may be partly due to ments with ot,her gases, the gases are always addi- other factors but the more concentrated solutions tional components since an aqueous subphase is also have lower vapor pressures which would give always necessary for the existence of a monolayer. less of a resistance due to water vapor and, thereAs a r(:sult, the water that must be in the mono- fore, a greater permeability. (d) Langmuir and Langmuir6 estimated the layer, due to the equilibrium between the aqueous subphase and the gas phase, can act as an extra amount of water vapor present in a monolayer during a steady state evaporation process. Recalresistance t,o the passage of other gases. (b) On raising the temperature, the large in- culation using newer data shows approximately crease in monolayer permeability observed in the the same concentration of water in the monolayer case of water is not observed for the other gases. as in saturated air. This supports the idea that The activation energies for the processes are about the amount' of water operat'ionally present in the 15 kcal./mole for water, 1 kcal./mole for carbon monolayer can be given by the vapor pressure. (e) The presence of water vapor in monolayers dioxide and 2 kcal./mole for nitrous oxide. The explanation offered in line with the above hypothe- has been suggested as the cause of a peculiarity sis is t'hat a rise in temperature has two opposing in another physical property, namely, the limiting effects on the total process. First, it causes an area of molecules. On a surface it, is about 10% increasc in the kinetic energy of the penetrating mole- greater than the cross sectional area of long straight hydrocarbon chain compounds in crystals. I%. K. * This inve~tigationwas supported by a Senior Research Fellow- Adam7 stated that the water in t'he subphase ship (SF-482) from t h e U. 8. Public Health Service. causes a '(. . . disruptive force antagonistic to the (1) M. Blank a n d F.J. W.Roughton, Trans. Faraday Soc.. 66, 1832 lateral adhesions between the long chains." He (1960). (2) M. Blank. A.C.S. Symposium "Transport Procemes Through . . it is probable that a also suggested that ((.
Monolayers," in press. (3) R. J. Archer a n d V. B.La Mer, J . P h y s . Chem.. 69. 200 (1955). (4) H. I,. Rosano and V. K. La Mer, ibid.. 6 0 , 348 (1956). (R) G. T. Barnes and V. k'. La Mer A C S. SyrnposiuiIi a s in ref. 2.
(6) I. Langmuir and D. Langmuir. J . Phys. Chcm., 31, 1719 (1927). (7) N. K. Adam. "The Physics and Chemistry of Surfaces." Oxford, 1941 p. 52.
Oct.. 1961
EFFECT OF VAPORS ON MONOLAYER PERMEABILITY TO CARBON DIOXIDE
certain number of water molecules are entrapped among the film molecules. . Although the expansion of monolayers in a surface has been ascribed generally to the effect of bonding with the subphase, the above possibilities remain. The hypothesis of water in a monolayer affecting permeability raises an important problem since foreign substances in monolayers have been shown to increase rather than decrease the permeability. Archer and La Mer3 found that benzene impurities greatly increased the permeability of a monolayer to water. Rosano and La Mer4 found that mixed monolayers always had lower resistances than the more impermeable monolayer. Extensive work by Robbins and La Me9 showed that occluded solvent molecules disrupted the monolayer and their results conrelated with water evaporation resistance measuirements. Cook and Riess also showed the delet’eriouseffects of several solvents on the low surface pressure region of stearic acid isotherms. These studies demonstrate that foreign substances in a monolayer have a disrupting effect that is accompanied by a greater permeability. However, there is a considerable difference between water vapor and! the hydrocarbon molecules, such as benzene, that were studied and that are known to affect cohesion in monolayers. In this paper the effect of water and methanol vapors on monolayer permeability will be determined. The :results will be discussed in terms of the process of monolayer permeability and, specifically, the effect of molecular size on the energy barrier. 11. Experimental Results
1GOB
. .”
To test the hypo1,hesis outlined in the Introduction, octadecanol was used as the monolayer since it is relatively impermeable and, therefore, has a fairly large effect on the gas transport rate. Many physical properties of octadecanol monolayers have been investigated and the permeabilities to water vapor and several other gases are known. They have a further advantage in that the OH group is relatively insensitive to changes in the subphase, e . g . , octadecanol permeability is unaffected by subphase pH over a very wide range. In the experiments, an excess of octadecanol was deposited from a spreading solution in 40-60” petroleum ether onto the surface of a ca,rbon dioxide absorbing solution. The solvent was pumped away leaving an octadecanol monolayer at its maximum surface pressure and a small solid phase which was shown riot t o affect the gas absorption. This method ensured the deposition of a monolayer at a reproducible surface pressure in a cell which did not permit quantitative deposition. The measurements involved the use of a temperature compensated differential manometer (described in detail in a previous publication’) to detect the absorption of carbon dioxide a t a monolayer free gas-solution interface and a t an interface with an octadecanol monolayer. From the two rates, it was possible t o calculate a permeability for the monolayer, assuming an initial steady state. This permeability is equivalent to that calculated in experiments on wat,er evaporation. Carbon dioxide was chosen as the gas since i t has been studied fairly extensively in this system. The absorption, which is free from convection effects within rather wide limits of the rate, it%easy t o control by varying the pH of the absorbing solution since this does not affect the monolayer. To see if one could vary the amount of water operationally present in the monoJayer by varying the vapor pressure of the subphase, LiC1 ciolutions a t various concentrations were ( 8 ) M . L. Robbins and V. K. La Mer, J . CoZZoid Sci., 16, 123 (1960). ( 9 ) H. D. Cook and II. E. Rice, J . Phys. Chem.. 60, 1533 (1956).
I
I
I
I.0
2.0 TIME (min).
3.0
Fig. 1.-Absorption rate of carbon dioxide on two different subphases with and without octadecanol monolayers (25.0’).
I
...
40-1
\
..
0
\
9
IO
20
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1
I
30
40
50
60
70
ec
1
90
7
100
y . V A P 0 R PRESSURE ( o f LI CI mlutlori),
Fig. 2.-The measured monolayer permeability to carbon dioxide as a function of the subphase water vapor pressure (25.0’). used. LiCl is among the most soluble salts and therefore give8 a wide range of vapor pressure decrease. The water vapor pressure as a function of LiCl concentration is taken from the “Handbook of Chemistry and Physics.” (The data are given for 100” but the fractional decreases in vapor pressure are independent of temperature.) Figure 1 gives the uptake of carbon dioxide as a function of time for two subphases of differing vapor pressure. (Curves b are displaced by one minute from curves a to make the diagram clear.) The curves represent averages of a t least three determinations and the individual points arp good to about =k 2 to 370. In the absence of LiCl, the vapor pressure is about the same as for 0.1 M solutions previously used to absorb the carbon dioxide, and it is possible to observe a large difference between a clean and an octadecanol covered interface. (See curves b for a borate subphase.) On the 8 M LiCl subphase shown in curves a , the data for the octadecanol monolayer and for the clean surface fit approximately on the same curve. This solution has a much lower vapor pressure and the data show a much greater permeability of the monolayer to carbon dioxide. Experiments were done on a number of different subphases and Fig. 2 is a plot of the monolayer permeability (Pr) to carbon dioxide as a function of the yo vapor pressure of the subphase (with pure water taken as equal to 100yo).The graph shows that the permeability increases as the vapor pressure decreases and that the variation is approximately linear. The data have been extrapolated to zero vapor pressure t o give a value of about 50 X 10-8 cm./sec. To examine the possibility that gases other than water vapor might have similar effects on the permeabiiity of a monolayer to carbon dioxide, experiments were done with subphases of water methanol mixtures. The method of varying the concentration of methanol in a monolayer by adding it to the spreading solvent leads to variations in concentration as a result of different degrees of evaporation, solution in subphase, diffusion rate in subphase, etc. With subphases of varying methanol concentrations, it is possible
MARTINBLANK
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Vol, 65
channeled into this “hole,” thereby increaaing the chances for collisions. C o ~ o n ais0 s can be more efFective if the monolayer chains act as intermediaries in collision pn>cessea In other words, in cases where carbon dioxide and vapor molecules just avoid each other in rmczw, in 8 monolayer the carbon dioxide may collide with a CH, p u p which in turn collides with the vapor molecule. Thw elTecta may be considerable despite the small number of vapor molecules simultsrneousy present in the monolayer, since the turnover rate is very
103 P, rs. % METHANOL
high. L
Ib
I
I
20 30 % METHANOL.
I
40
Fig. 3.-The m d monolayer permeability to carbon dionde aa a function of the concentration of methanol in the subphase (25.0°). to know the equilibrium vapor pregsure of the methanol and, therefore, the amount of methanol “in” the monolayer at dl times. Methanol, which is somewhat eimilsr to the hydrocarbons used aa solvents in monolayer studiea, is a
simple molecule that is miscible with water in all proportiom. The solutions exhibit a wide range of v a p r presgurea which are known over the complete range of &urea. The experiments on these s u b p h yielded the same t of raw data as are given in Fig. 1. The resulte for 5% me% an01 are similar to those for 8 M LiCl (curyea a), and the data for 20% methanol are similar to thoee gven for-no LiCl in the subphase (curves b). From experiments wlth subphaaea of varying water m e t e o l ratio, the permeability of octadecanol to carbon dionde waa calculated and these valum are given in Fig. 3 as a function of methanol concentration in the pubphase. The graph ahom that the pres enoe of small amounts of methanol in the monolayer increams the permeability very markedly. Thb parallels the d t a . mentioned eadier, of studies on water permeability where d l amounts of molecules such as benzene greatly increase monolayer permeability. However, when one increases the vapor presaure of methanol, the monolayer bec o m e ~leae permeable to carbon dioxide. Apparently, after one subtracts the initial disrupting deet of methanol on monolayer coheaion, increaaing methanol concentration has the same effect as increasing water vapor concentration. When the vapor presaure of methanol is approximately equal to that of the water vapor, the monolayer permeabiity drops below its initial value, demonstrating that the blocking effect is not j u t a return to the initial value after the monolayer disruption but, rather, a definite deerease in permeability.
III. Discussion A. The Effect of Vapors-The observed decrease in monolayer permeability to carbon dioxide in the presence of water and methanol vapors is not a matter of simple collisions between the vapor and carbon dioxide because a comparison of the results obtained with octadecanol monolayers and those obtained in the absence of monolayers should keep the number of collisions constant. However, the monolayer may affect the number of collisionsbetween the carbon dioxide and the vapor molecules in a non+pecific manner. It also may trap some of the vapor molecules and, therefore, muse a change in the permeability. Let us consider some changes in the pattern of molecular collisions that may occur as a result of the p m n c e of a monolayer. For example, the point at which a vapor molecule enters the monolayer can become a preferred site for the entrance of another molecule, since the vapor has already done part of the work of pushing the monolayer chaiis apart. A carbon dioxide molecule may be
Let us now consider the possibility that molecules of the vapor are trapped in the monolayer. One can aak why should the monolayer trap the vapor and not the carbon dioxide. If the carbon dioxide were trapped there would be a dif€erence in the Pf depending on the partial pressure and this is not found ovei a tenfold range of pressure. Therefore, one also w d d have to explain specificity if the above explanation is used. Furthermore, a caiculation of the number of water molecules present simultaneously in a monolayer, assuming the same concentration as in saturated air [see I(d) above], gives about one molecule to every lo4 monolayer molecules. The amount is negligible in terms of the ability to comprees the monolayer and also in terms of the ability to act as a physical barrier. However, the idea of occlusion of vapor in the monolayer is supported to some extent by the data of Fig. 3. At low values of methanol vapor pressure the permeability of octadecanol is greatly enhanced, while at higher values the permeability decreases and one observes the blocking effects shown by water vapor. Considerable evidencc favors the idea that an increase in permeability is associated with the occlusion of a foreign molecule in the monolayer. If there is occlusion at low vapor pressures, it seems equally posaible at higher vapor pressures. We, therefore, have two possibilities for explaining the extra resistance due tcr water or methanol vapors which involve either short (collision effects) or long (occlusion effects> lifetimes in the monolayer. In either case, there must be coilisions between the carbon dioxide and the vapor (that is either transient or trapped}, so collision effects is the simpler hypothesis and, therefore, the more desirable one. The variation of Pr with water vapor pressure pig. 2) appears to be linear and the data have been treated as such. The independent variable, the % vapor pressure, can be replaced by the actual vapor pressure or the number of molecules per unit volume in the gas phase to give equivalent graphs. This shows that the value of Pf is proportional to the number of molecules of vapor in a unit volume of tbe gas phase and, presumably, in a unit volume of monolayer. Therefore, a reasonable explanation for the observed linear variation is that the decrease in permeability is due to collisions of carbon dioxide with the vapor and that the number of collisions variea linearly with the concentration of vapor, as in the c ~ s eof binary collisions. [Although this serves as a rationalization for the extrapolation to zero vapor presrmre, the as-
Oct.. 1961
EFFECT OF VAPORS ON MONOLAYER PERMEABILITY TO C.4nuo.u DIOSILK
surned linearity of Pf has not been proven. In fact, the linear behavior has no meaning above lUOoI, vapor pressure since Pf will approach but will not equal zero nor will it become negative as implied by the straight line. However, points above 100% vapor pressure cannot be obtained on water 3ince one must raise the temperature which in and of itself increases Pf. l f one adds a volatile substance such as methanol, the data (Fig. 3) do show a levehng off of Pf when the vapor pressure becomes very high. This may demonstrate a nonlinear mechanism, but it may be relevant only a t higher concentrations of vapor. I n addition, these data are complicated by having both water :tiid methano! vapors as well as a disrupting effect on the monolqper so they may not be relevant at all. However if one considers the possibility that the originally proposed linear variation is not the case, then the extrapolated value may be higher. ] The value of Pf extrapolated to zero vapor pressure represents the l’f that would be obtained using a pure carbon dioxide gas phase. Since the most reliable points of kig. 2 are those of lowest Pi, the extrapolation is made subject to error over and above that due to the scatter of the points and to the fact that bhey extend over only half of the range. Nevertheless, the trend of the points is clear and an error iri the extrapolated value should not, affect the qualitative conclusions to be drawn. B. The Effect of Molecular Size.-The movement of gases through monolayers cannot be described as a cla,ssical diffusion process because the calculated diffusion coefiicients vary with the monolayer thic-kness. Ynwever, the monolayer does impede the movement of the gas, and it is necessary to formulate the effect in some way. Langmuir and co-worker+I9 originally described the effect of a monolayer in terms of an extra energy barrier to the passage of gas. This has proved to be a fruitful approach and there have been many attempts t u efiaracterize the energy barrier by studying the effeccs of surface pressure, the chain length and nature of the polar group of monolayers, as well as the eEect of temperature. Rerently, Blank and La Mer“ attempted to char:wttrize the molecular processes involved during penetration by calculating the energy for the penetration, by water, of a single layer of CH2 groups in fatty alcohol inonolayers as a function of the separation of the CHZ grovps. They found a dependence on separation distance which indicates that the water molecule must compress the CHa groups against the repulsion forces. This is ‘reasonable in close packed monolayers and it is also in line with Langmuir’s original hypothesis of the energy barrier where a gas must do work against the surface pressure (the repdsive forces) of a monolayer. Langmuir’s formulation of the energy barrier was a t near equilibrium conditions, where a Boltzmann factor gives the variation in the fraction (f) of molecules that can penetrate a monolayer in terms of extra energy ( E ) needed by the molecules. E can be evaluated in terms of monolayer proper-
170 1
ties by setting it equal to the work done by the gas (of cross-sectional area ao) against the surface pressure (n) of the monolayer. Therefore f = e - E / k T = G-rcadkT
(1)
where k = Boltzmann constant and T = absolute temperature. The rate of penetrating the monolayer is proportional to f, so equation 1 can be looked upon as (proportional to) a permeability and the inverse as a resistance. Since the early studies were confined almost exclusively to water penetration, a. was constant and in succeeding papers it was neglected. Recent studies with several penetrating gases have made it possible to examine the applicability of equation 1 to the penetration process. The first results’ on carbon dioxide permeability could not be compared directly to the water vapor permeability measurements since they were affected by the water vapor. However, the permeability extrapolated to zero water vapor pressure (50 X cm./sec.) represents the permeability of a “water free” octadecanol monolayer to carbon dioxide and is comparable to the measurements of water permeability. Two values are in the literature for octadecanol permeability (at maximum surface pressure) to water. They are approximate!~‘~ 300 X and 250 X 10-* cm./sec. when converted to the same units. The ratio of Pf’s therefore is about 5-6 to 1, the monolayer being more permeable to water. The values of a0 for water and carbon dioxide are not known accurately since they depend on the techniques used in their evaluation. One can get an approximate idea from data summarized by Moelw.m-Ilughes,12 which include effective radii and incompressible radii based on gas viscosity measurements and so constitute a fair range of values. Substituting values for the cross-sectional area as measured by the square of the radius, (into equation I) one obtains a permeability ratio of about 3-6 ta 1, the monolayer being more permeable to tx ater. This is a reasonable demonstration that the effect, of molecular size can be predicted approximately from the equation originally proposed by Langmuir. The carbon dioxide values are the most reliable at the present time for this kind of test but data for nitrous oxide also indicate reasonable agreement. Nitrous oxide has approuimately the same radius as carbon dioxide and octadecanol has approximately the same permeability to both gases. The available data for oxygen, which has the same permeability as carbon dioxide but a smallcr radius, do not agree with equation 1. However, the statements about nitrous oxide and oxygen assume that the interference due to water vapor has the same effect as in the case of carbon dioxide. This assumption is not valid in the case of oxygen if the interference effects are due to collisions with transient or occluded water vapor. Carbon dioxide and nitrous oxide have the same molecular weight but oxygen is considerably lighter and, therefore, makes many more collisions a t the eame pressure. The greater number of collisions
(10) I. Langmuir imd V. J. Schaefer. J . Franklin Inat.. 136, 119 (1943).
(11) M. Blank m d V. K. La Mer, A.C.S. Sympoiium PI in ref. 2.
(12) E. A . i\loelwyn-Eiighed. “Physical Chemistrl 1957, p. 597.
”
Peroarnon
MARTIN BLANK
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7 2m
0 ‘
__-
5
Ib
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I5 20 25 30 35 SURFACE PRESSURE(dyneslcm)
Fig. 4--The calculated energy for water penetrating a layer of CH2 groups (in fatty alcohol monolayers) as a function of the surface pressure (25.0”).
would c:tuse greater interference, hence a higher permeability on extrapolation to zero pressure and the pos:ibility of agreement with equation 1. Therefore, the generality of equation 1 with regard to the dependence of the permeability on molecular size of penetrating molecule is still to be demonstrated b u t it does seem to apply to the available data. It might be added in passing that dependence of the rate on the size of the penetrating molecule makes the monolayer perniselective (ie., it endows the monolayer with functional “pores”). In addition to predicting an effect due to molecular size, equation 1 predicts that the logarithm of the permeability or the resistance (r) should vary linearly with the surface pressure. This has been observed in some cases. Archer and La Mer3 showed this for the S phase of fatty acids Rosano and La Mer4 and Barnes and La Mer5 have given data for the high pressure phases of fattv alcohols which agree approximately with the prediction Blank and La Mer’: showed that the \ d u e of E for penetrating a CH1 group, E(CHz), ir: the high pressure region of fatty alcohols is approximately h e a r uith P (see Fig. 4). (This ccliidition i.; equivalent to In T being proportional T . ) in these cases of agreement as to functional dependence. the proportionality wnstant leads to v,iriable values of a0 that are about an order of magnitudp too small and that. therefore, indicate a !nvk of quantitative agreement There are many case6 where epen the qualitative dependence of the permeability on t is not as predicted. so that equation 1 carinot be n general statement of the energy barrier. 1:quation 1 appears to be but one of the fnctors in the mergy barrier and one can expect agreemen, only under conditions that will become evident u3on examination of the energy barrier as n whole. C. The Energy Barrier to Monolayer Penetration.--Aside from the molecular size factor given in equation L, Langmuir’O found the need for a monolayer chain length term. Archer and La Mer3 iound that each CH2 group in a11 LC phase fatty acid monolayer contributes about 300 cai./mole to the energy barrier for water penetration. The magnitude of the contribution is P independent, but in tht S phasc- nherc it i. al+c~P Independent thr rnagii~tudr~I \ ittmut 300 mi. niolc Blank”
Vol. 65
found that for the penetration of carbon dioxide, oxygen and nitrous oxide, about 350 cal./mole were added per CH2 group. It was not possible to check the ?r dependence here but Blank and La Mer’’ showed that the contribution per CH2 group can be ?r dependent. Figure 4 shows that E(CH,) varies with ?r in fatty alcohol monolayers and that the variation can explain the three different values quoted as being due to different states of monolayer packing. The ?r dependence of the monolayer chain length factor can account for the inability to determine the correct value of a. from the ln r vs. n data and it also points up the complexity of n as a variable in penetration experiments. There are other factors that enter into the energy barrier since monolayers of the same length (number of CHP groups) and at the same n do not have equal permeabilities to the same penetrating species. Rosano and La filer4 found that incompressible monolayers are far less permeable than compressible ones, but compressibility per se does not determine permeability. Cholesterol monolayers are relatively incompressible and the surface pressure can exceed 40 dynes/cm., yet these films have little or no effect on the passage of gases. Since the saturated straight chain compounds are the least permeable, the important factor that appears to t F: significant is the effect of the structure of the polar and hydrocarbon parts of the molecule on the ability to pack in an interface. (The polar group is also important in causing an orientation of water moleeuies in the liquid, which should affect the permeability of a monolayer as well as its stability.) This factor, which can be called intrinsic to the monolayer type, may also be n dependent, especially if the molecule ran be easily deformed upon compression. It is possibie to summarize our knowledge of the energy barrier by combining the various factors that have been discussed. E
=
“flS( X ’ > 1( X ) ,
aaa]
(2
where K ( n ) is the intrinsic factor, E(*) is the chain length factor and nao has been given in equation 1. By using the same gas so that TU,, is constant aiid by using a homologous series so that K ( n i 1s (>onstant, it has been possible to determiiie l(Pj. I n this paper, K ( a ) and 1(n) were kept Coiistani so that the nu0 dependence could be investigated. I’or these and other reasons E has been broken up into a number of factors even though the surface pressure, monolayer chain length and packing are all inter-related. -1s a matter of fact. ;;TI definitely varies with the monolayer state3 so it probably would be more meaningful to lump K i n ) and Z(S) together as a monolayer factor [that is distinct from T% which depends on the gas that is penetrating]. If E , as in equation 2 or as modified h y combining K ( T ) and Z(a) in a single term. is substituted into equation 1, the effects of the above variables on the penetration rate become apparent and one also sees the framework for considering the temperature dependence. Bside from the interfering effects of mater vapor, there is a further complication in the analysis of the permeability that is due to the possibility of hydrarion of the gas in the aclueous phase. Hawke and
Clc ., 196;
EFFECT O F VAPORS O N hl0;h;OLAYER PERMEABILITY TO CARBOX
-4lexander l 3 determined monolayer permeability to cwbon dioxide and hydrogen sulfide where the gas moved from the aqueous phase to the gas phase In the rases where the systems were the same as tilose dudied by the author’s technique, the results agreed as to order of magnitude but the permeabilities were always lower. Furthermore, the values for hydrogen sulfide were lower still. Carboil dloxtde arid hydrogen sulfide have approximately equal diameters12 so that all the factors in equation 2 are approximately equal. However, if one con-idcrc that there can be some penetration as a hydrated species thus increasing ao, or what seems more ,rkely, that there is an additional restraining 4tep due to the need for dehydration before the gas car) penetrate, then one can accouilt for the diflwrnccs observed. The Bunsen coefficient wuuld t~ ‘i measure of the affinity for water and, Einw the value is about three times higher for hydroeui sulfide (than for carbon dioxide,, i; n o u l j have less of a tendency to dehydrate. l’hereic ]re riydrogeri sulfide would pass through more slov I>* and carbon dioxide mould pass more slon lv frc rn the aqueous phase than from T h c are many data availabie for thc gas phast thc penel railoil of phyma membranes (which arc’ 11 3)
7
1
i l i i ~ i i rani3
1
‘Ilrxmder
1C
C
S>rnrosiun>ae
iii
ref
DIOXIDE
1703
bimolecular layers about 75 A.thick) that can be explained on the basis of a process measured by the partition coefficient. This would support the above explanation and account for the differences observed. A . S. ~IICHAE.I,S (Jlassachusetts Institute of Technology;. -1s there any possibility that convective processes (created by tempmitiire, density, or surface tension gradients) in the liquid filni in the absence of a deposited monolayer arc: eliminated iri the presence of a monolayer, thereby increitsing the persistence t,o mass transfer independent of any specific harrier action of the nionoiayer itself? 31. BLa.uli.-‘l’his question has been considered (ref. 1 ) and it seems unlikely that convective processes play a role i r i thv systems studied. I n brief, the reasons are: 1. Calculatioris of the absorption rate when no monolayer is present (assuming it liquid diffusion limited process) coincide \\.it11 the ohserved rate data. This indicates that there are no convective processes during absorption in the aimnce of monolayers. The reason for this may be that the iricrease in temperature due t o heats of solution and reaction cancals the effect of an increase in density due to mass absorption. The absenct, of convectioIi may also be due to the .short timrs of thc experiments. rtain monolayers offer no measurable diffusion imrier i)ut nevertheless affect the transfer across interfaces iti stirred systems iiy interfering with convctctive processes n m r thr interface. I3ovine serum albumin arid cholesterol :Ire euainples of thest suhstanres and th(by were found to have 110 effect 011 the trmsport of gases,
