Charcoal Adsorption Equilibria for Light Hydrocarbons C. J. WALTERS California Research Corp., La Habra, Calif.
DSORPTION on activated charcoal has been used by the petroleum industry for a number of years in field and laboratory tests for heavier hydrocarbon content of natural gases, Also of importance is the separation of light hydrocarbon mixtures on a commercial scale by adsorption on charcoal. The development of these processes depends on equilibrium adsorption data for the mixtures involved. Few such data appear in the literature. The purpose of the present investigation was t o provide additional equilibrium data and to determine whether characteristics of a charcoal could be established which would a t a given temperature and pressure permit prediction of vaporadsorbate equilibria for light hydrocarbon mixtures of varying composition. A similar investigation by Lewis and others (W,3)is reported in the literature. These authors used both thermodynamic and adsorption potential methods in an effort to correlate equilibrium data on pure components with data on mixtures. The adsorption potential method was the more successful but was found generally unsatisfactory for mixtures in which the molecular weights of the constituents differed appreciably. Inadequate data prevented further development of their correlations. A review of the adsorption potential theory (1)indicated that a limited amount of data could be used to make B satisfactory evaluation of the method if a broad enough range of molecular weight and mixture composition was used. Methane, propane,
G
H
Figure 1.
To vacuum pump and McLeod gagc Mercury reservoir Manometer Thermal conductivity cell E. Adsorption bell F. Compensator 0. Gas buret H. Leveling bulbs I. Capillary tubing
EXPERIMENTAL PROCEDURE
Apparatus. The experimental requirements were that the vapor be in equilibrium with the charcoal a t a measured temperature and pressure and that the vapor composition be uniform throughout the system. A schematic diagram of the apparatus used is shown in Figure 1. The adsorption temperature was controlled to within 0.5" F. by a bath surrounding the adsorption tube. Pressure was measured to the nearest 0.5 mm. on a mercury manometer. The uniformity of composition was determined by a thermal conductivity cell mounted in the system. Circulation and mixing of the vapor were achieved by a method described by Lewis and others (2)(Figure 1). Mercury from reservoir B dropping into capillary leg I of the apparatus caused the vapor to circulate continually through the system. Materials. Phillips Petroleum Co. research grade hydrocarbons and a single sample of reactivated Columbia G charcoal were used in all determinations. The reactivation process consisted of treating fresh charcoal for 6 hours with 1000° F. superheated steam. The surface area of the charcoal as determined by nitrogen adsorption was 1368 square meters per gram. Pure Component Measurements. The 60" F. isotherms of methane, propane, and n-butane were determined over the pressure range of 0 to 760 mm. Preparation of the charcoal before each run consisted of evacuation to less than 0.01 mm. at 300" F. for 1to 2 hours. Vapor was admitted to the system from the buret in measured amounts and the pressure read after each addition. Mixture Measurements. Vapor mixtures were blended in external containers and introduced to the system in measured amounts until the system pressure was 760 mm. A flow of mercury was started in the circulating system and the pressure was maintained at 760 mm. by means of a leveling bulb. After 30 to 40 minutes, the thermal conductivity reading usually became constant, and equilibrium was assumed. After 10 minutes of additional circulation, the adsorption cell was closed off from the rest of the system, and a sample of the equilibrium vapor was withdrawn for infrared analysis. The amount and composition of adsorbate were established by the difference between moles of each com onent in the feed and in the equilibrium vapor. The ir&ared instrument used was a Perkin-Elmer recordingtype spectrometer that had been calibrated with pure com onents and known mixtures. Several adsorption eaks were u s e t which gave a number of independent checks on t f e analysis. CORRELATION O F DATA
H
Circulating apparatus for vapor-adsorbate equilibrium measurements A. B. C. D.
and n-butane bracket the light hydrocarbons of interest, and an experimental investigation of vapor-adsorbate equilibria for these compounds as pure components and in mixtures was prompted in order to test the adsorption potential relationship more thoroughly as a method of correlating data on pure components and mixtures.
Pure Component Data. The volumes adsorbed were converted to millimoles per gram of charcoal and are plotted against pressure in Figure 2. Corrections were made for the compressibility of the vapors a t the conditions of temperature and pressure in the buret and system vapor space. The n-butane determination at 308 mm. does not lie on the smooth curve through the other points. The deviation is probably due to inadequate time for equilibrium. The desorption points for n-butane are likewise off the smooth curve, which is also possibly due to insufficient time for eqnilib2544
INDUSTRIAL AND ENGINEERING CHEMISTRY
December 1955
0
A ADSORPTION
0
A
2545
rium, even though several hours were allowed betyeen desorption measurements. The first step in establishing an over-all correlation of the data involved only the pure components. The Polanyi adsorption potential as modfiied by Lewis and others (3) was calculated for each component a t each experimental point on the adsorption isotherms. The following form of the adsorption potential term was used:
OEMRPTIOH
RT f " - In 2
0
eo0
400
200
P M Y . Of
Figure 2.
vi j: where R = gas constant T = absolute temperature Vr = molal volume of saturated liquid a t boiling point temperature cor r e s p o n d i n g to adsorption pressure j: = fugacity of pure component as saturated liquid at adsorption temperature f: = fugacity of pure component as vapor a t adsorption temperature and pressure
HG
Adsorption isotherms on Columbia G charcoal a t 60' F.
02
I
I
I
I
I
I
I
4
6
e
I
2
IO
I2
I4
16
6m VI
Figure 3.
1.
A plot of adsorption potential against adsorbate volume, NcVi, where N i is moles of component i adsorbed per gram of charcoal and Vi is as defined above, is shown in Figure 3. Data presented for the light hydrocarbons by Maxwell (6)were used in developing this correlation. A smooth curve can be satisfactorily passed through all points determined a t pressures greater than 12 mm. Below this pressure, the points rapidly diverge downward from the curve, as shown by the dashed lines in Figure 3. The adsorption potential has been found ( 1 ) to be nearly independent of component, temperature, and pressure a t a given total volume of adsorbate. On this basis, an isotherm for any light hydrocarbon a t any temperature and pressure range can be approximated with the adsorption potential from Figure 3 for the particular charcoal used in these tests. Because the data
18
mom CM?
Generalized correlation for light hydrocarbons on charcoal at 60' F.
1.0,
MOLE FRACTION OF LIGHT COMPONEYT IN ADSORBATE
Figure 4.
Vapor-adsorbate equilibria at 1 atm. and 60" F.
MOLE
Figure 5 .
FRACTION
OF
N-BUTANE
IN
ADSORBATE
Vapor-adsorbate equilibria for ternary mixtures at 1 atm. and 60' F.
INDUSTRIAL A N D E N G I N E E R I N G CHEMISTRY
2546
Table I. Component CH4 CsHs CHI
CxHa CHI CaHs CH4 CsHs CHI CIH8 CHI CsHs CH4 CsHs
Adsorption Equilibria Data on Binary and Ternary Mixtures
Yi Nethane-Propane Binary Mixtures 0,297 0.946 0.703 0.054 Xi
__
-
1.000 0.071 0.929
1,000 0.386 0.614
A7i
Component
1.107
CH4 CaHs n-CdHto
2.026 -3.733 0.413 5.432
-
-
__.
1,000 0.523 0.477
1.000 0.985 0.015
5.845 1.437 1.309
__
-
1.000 0.104 0.896
1.000 0.536 0.464
2.746 0.555 4.774
-
-
1.000 0.171 0.829
1.000 0.860 0.140
5.329 0.803 3.899
1,000 0.171 0.829
1.000 0.175 0.825
-
-
1.000 0.870 0.130
4,702 0.798 3.865
1,000 0.860 0.140
4.663 0.845 3.971
__
-
-
1.000
1.000
Propane-n-Butane Binary Mixtures 0.101 0.258 0,899 0.742
-
-
1.000 0.214 0.786
1,000 0,520 0.480
1 000 0.340 0.660
1,000 0.717 0.283
5.641 1.952 3.793
__
-
-
1,000 0.504 0.496
1.000 0.860 0.140
5.745 2.877 2.826
-.
lIO0O 0.729 0.271
--
1,000
-_.
1,000 0.951 0.049 __.
1 I000
-_
CHI CsHs 7l-C4Hlo
1.000 0.452 0.280 0.268
5.125 0.272 1,111 3.878
CHI CaHs n-CaHlo
0.017 n. io8 0.875
0.246 0.138 0.616
0.095 0.600 4.865
CHI CaHs n-C4Hlo CHI CiHs n-CdHia
5.703 4.166 1,546
CHI CaHs n-CdHio
__
-
__
0.209 0.317 0.474
0.010 -
1 000
1.000 0.653 0.128 0.219
0.083 0.119 0.798
-
--
5.560
1.000 0.932 0.058
-
~
.ooo
1,000
I
0.054 0,057 0,889
0.340 0.068 0.592
-
0.992 1,508 2,253
-I
4.753 0.427 0.616 4.119
--
5.162 0.306 0,321 5.026
-
1.000
1,000
5.653
0,030
0.348 0.571 0.081
0.165 2,743 2 252
0.533
0.437 -
1,000 0,034 0.278 0,688
-
1 .ooo
I
5.150
0.223 0.465 0.312
0.190 1.559
-
3,858 __
1 . no0
1.000
5.607
n.ooo n. 138
0.135 0.243 0.622
0.002 0.774 4.823
0.862 1 .ooo
5.712
apply a t one temperature only, RT was arbitrarily set equal to 100 in Figure 3. Mixture Data. Twelve binary x y determinations are shown in Figure 4. Nine ternary x - y determinations for nbutane are shown in Figure 5 ; however, this ternary plot is only a qualitative relationship based on the ratio of methane to propane in the vapor initially charged, rather than on their ratio in the adsorbate. The latter would be more desirable, in that some more definite relationship might be expected, but experimental conditions were such that fixed ratios in the adsorbate could not be readily controlled. Data on the experimental mixture are given in Table I. A useful correlation for mixtures must predict the total amount adsorbed as well as the equilibrium conditions between phases. The correlation should furthermore require a minimum amount of experimental data to characterize the adsorbent. A plot of total volume of pure component adsorbed a t 60’ F. and 760 mm. us. molecular weights shows a nearly linear relationship (Figure 6). By using the average molecular weight of adsorbate, the total volume of mixture adsorbed, zNiV4, was found to be given by the pure component curve of Figure 6 to m-ithin 9% of the experimental value. The empirical correlation used by Lewis and others ( 3 ) to get total moles adsorbed was found less satisfactory as a means of applying data on pure components to the prediction of behavior of the mixtures. However, the method does give an accurate internal correlation of data for each set of binary and ternary data.
-
-_
CHI CJH~ n-CdHio
6.107 1.204 4.437
3.104 0.374 2.280 2.471
1.000 0.052 0.211 0.737
CH4 CsHa n-CdHlo
0.617 5,490
-
1.000
1,000 0.073 0.446 0.481
I
4.816
-
... ... -
-
-
CH4 CaHs n-CdHio
-
__
Xi Yi xi Methane-Propane-n-Butane Ternary Mixtures 0.409 1,000 1.271 0.207 0.643 0.384 1.190
1.000
-
I
CHI CaHs n-C4Hio
-
-
-
Vol. 47, No. 12
-
__
1.000
5.599
The adsorption potential for the pure component was applied to mixtures by the following relationship for each component:
where x = mole fraction in adsorbate y = mole fraction in vapor f a = fugacity of component in adsorbate mixture f a = fugacity of component in vapor mixture y = activity coefficient The fugacity ratio, fa/fu, in Figure 3 corresponds t o f:/$ for for mixtures. pure components and fallv The need for the activity coefficient, y, became apparent when the pure component correlation (Figure 3) was used to calculate using known z - y data. The ratio of as calculated to fi as given by Maxwell (6)was found to vary from unity apparently according to some function of composition. By taking saturated liquid at adsorption temperature and pressure as the standard state, this ratio was treated as an activity coefficient. The most satisfactory parameter for correlating y was found to be the arithmetic difference between the average adsorbate molecular weight, zxiM6, and the molecular weight, Mi, of the component, as shown in Figure 7 . Relative volatilities were determined by means of the final correlation for each of the mixtures runs. Using the experimental total volume of adsorbate per gram, ZNiVi, the value of (lOO/Vc) In fo/fu was determined from the adsorption potential
fi
fi
INDUSTRIAL AND ENGINEERING CHEMISTRY
December 1955
method proposed by Maslan, Altman, and Aberth (4)was tried but was found also relatively unsatisfactory. These authors proposed that the fugacity of the total mixture, rather than that for each component, be found from the pure component correlation and taken equal to the product of fugacity ratios raised t o powers corresponding to the mole fractions in the adsorbate.
6
A
PURE COMPONENTS TERNARIES C,- Cs- NC4 BINARIES CI- C3
0
BINARIES Cj-
X
0
-
5 -
N C,
w ka 1 m u ari
CONCLUSIONS
4 -
2: :E l 0
$
3 -
W
10
I
I
I
I
20
30
43
53
E’xiMi, AVERAGE
MOLECULAR
WElGHi
Figure 6. Relation of total volume of adsorbate at 1 atm. and 60” F. to average molecular weight of adsorbate
1.2
A449 M I X i U R E DATA X METHANE
0 PROPANE A N-BUTANE
5 w -
2541
0.8
0
Y
”
Y
The present correlation will be of use a t other temperatures and pressures and for other Columbia G charcoals only if the activity coefficient does not vary greatly with these factors. Only further study can determine whether this is true. Assuming that the activity coefficient is relatively independent of temperature and pressure and variations among Columbia G charcoals, it is possible to apply the present correlation a t any new conditions without further experimental data. The present adsorption potential for the pure component can be used to estimate a volume of pure component, iZ‘,V,, adsorbed us. molecular weight curve a t the new conditions. With an estimated total volume of mixture adsorbed, an adsorption potential for each component can be obtained and the vapor-adsorbate equilibria estimated. Where more certainty is required, a t least data on pure components would be necessary. Comparison of the pure component correlations shows that for a given adsorption potential the Columbia G charcoal used in the present work held 10 to 15% more adsorbate than the sample used by Lewis and others ( 3 ) . The plot of ZN,Vi us. average molecular weight (Figure 6 ) shows definite effects of composition on the total amounts adsorbed. A small amount of methane increased the quantity of propane adsorbed over that of pure propane and in turn a small amount of propane increased the amount of butane adsorbed. This effect has been observed before and is discussed by Brunauer ( I ) , although no explanation has been found.
0
06
c -> F
NOMENCLATURE
u
’
f’.
0‘
02
I 10
0
I 20
(ZXtMi-Mi
Figure 7.
I
30
I 40
1
Correction for composition of adsorbate
curve for the pure component (Figure 3). Pure component specific volumes, V i , were then used t o calculate fJfU and y was determined from Figure 7 . Equation 1 then gave sly, from which relative volatilities were obtained. The average deviation of calculated relative volatilities from experimental values was 35%. The average algebraic deviation was +9%. The average deviations were somewhat less when the tot a1 volume adsorbed was determined from Figure 6 as a function of average molecular weight. Variatioas in the uae of the adsorption potential method were tried but without improvement over the method presented. Use of specific volumes of saturated liquids a t 60” F. for propane and n-butane and the supercritical volume for methane a t 60’ F. and a pressure corresponding to the extrapolated fugacity gave poorer results. The same was true when average specific volumes for the mixtures were used in the adsorption potential term. The
x y
fugacity of pure component as saturated liquid at adsorption temperature f: = fugacity of pure component as vapor a t adsorption temperature and pressure fa = fugacity of component in adsorbate mixture fD = fugacity of component in vapor mixture M = molecular weight N = moles of adsorbate per gram of charcoal V = molal volume of saturated liquid a t boiling point temperature corresponding to adsorption pressure = mole fraction of component in liquid = mole fraction of component in vapor =
LITERATURE CITED
(1) Brunauer, S., “Adsorption of Gases and Vapors,” vol. I, “Physi-
cal Adsorption,” Princeton University Press, Princeton, N. J., 1945.
(2) Lew-is, W.K.,Gilliland, E. R., Chertow, B., and Cadogan, W. P., IXD,ENG.CHEW,42,1319(1950). (3) Ibid., p. 1326. (4) Maslan, F. D., Altman, hl., and Aberth, E. R . , J. Phys. Chem., 57,106 (1953). (5) Maxwell, J. B.,“Data Book on Hydrocarbons,” Van Nostrand, New York, 1950. RECEIVED for reiriew September 20, 1954.
ACCEPTED April 2F, lY55.
