Illlil ll
(.gun]galA
ELSEVIER
Fluid Phase Equilibria 135 (1997) 137-144
Solubility of methyl fluoride in some alcohols
C . S . O . Silva, I.M.A. F o n s e c a ' , L.Q. L o b o
I)~7)artamento de Engenharia Quhnica, Unicersidade de Coimhra, 3000 Coimhra, Portugal
Received 23 September 1996; accepted 2 February 1997
Abstract
The solubilities of methyl fluoride in some polar solvents (methanol, ethanol, propanol and n-butanol) have
been measured at temperatures ranging from about 280 to 300 K, and at atmospheric pressure. The solubility is
the lowest in methanol, increasing with the C-content of the alcohols. H-bonding factors, based on the ideal gas
solubilities and the solubilities in the alcohols, appear to be linearly dependent on the H-bonding factor in
methanol. The molar Gibbs energy, enthalpy and entropy of solution were calculated from the experimental
results (at I atm partial pressure of the gas and 298 K). © 1997 Elsevier Science B.V.
Kevword.~." Experiments; Data; Solubility of gas in liquid; Methyl lluoride: Alcohols
I. I n t r o d u c t i o n
The solubility of gases in liquids is an excellent tool to investigate solute-solvent intermolecular
forces in the liquid state. In this work the solubility of methyl fluoride in some polar solvents
(methanol, ethanol, propanol and n-butanol) has been measured at temperatures ranging from about
280 to 300 K, using a modification ot" the experimental technique of Kaszonyi et al. (1992). Methyl
fluoride is a molecule with particular theoretical interest as it possesses a high dipole moment
( #t = 1.85 × 10 i, esu) and almost no quadrupole ( Q = 0.04 × 10 2~' esu). As far as we are aware
solubilities for this substance were reported only for water and heavy water (Wilhelm et al., 1977). Bo
et al. (1993) have also determined the solubilities of several nonpolar gases in normal l-alkanols.
It has been flmnd that the solubilities of gases in polar solvents are usually much lower than the
corresponding ideal solubilities, except in a few cases (Hayduk and Laudie, 1973). This reduction in
• Correslmmdin g author.
()378-3812/97/$17.(X) ,,c) 1997 Elsevier Science B.V. All rights reserved
PII S 0 3 7 8 - 3 8 1 2 ( 9 7 ) 0 0 0 5 7 - 5
138
('.5. O. Sih a ct al. / / ' h d d I'lm~e l:'quilihria 135 (1997) 137-144
solubility has been attributed to hydrogen-bonding in the solvent which can be represented by an
H-bonding factor (Hayduk and Laudie, 1973):
A
(-l(~'~'l"c"t-- A"id
(1)
defined as the ratio of the actual solubility of the gas expressed in mole fraction ( x ) to its ideal
solubility ( x +d) in a particular solvent. Hydrogen-bonding factors have been useful in relating gas
solubilities in one hydrogen-bonding solvent tt> those in other (similar) solvents (Short et al., 1983).
Wilhelm and Battino (1973) defined A H ' , the standard enthalpy change on solution, commonly
called enthalpy of solution, as the change in enthalpy for the process:
Y(gas, hyp 1 atm) --* Y ( s o l u t i o n h y p , x = 1)
for one mole of gaseous Y (1 arm, ideal gas), where .v is the mole fraction of Y in a hypothetical
liquid solution at infinite dilution. AG" and AS" are defined similarly. These functions are obtained
by straightforward thermodynamics.
2. Experimental
The experimental apparatus for the solubilit 5, measurements is represented in Fig. I. The principle
of the method is to bring a measured volume of liquid into contact with a known w~lume of gas at a
given temperature and pressure. After equilibrium has been attained the change in the gas volume
yields the amount of gas dissolved in the liquid and hence the solubility.
The experimental technique used in this work is similar to that reported by Kaszonyi et al. (1992)
with the following modifications: (i) the temperature of the liquid sample is measured with a platinum
resistance thermometer (calibrated against a precision mercury thermometer graduated in 0.01°C
Ga~
V acUtlll|
2
•
.
3
4
'
I
[
•
•
XIj
Walcr
Fig. I. Stflubility apparatus: (I), (2) and (3) ~,tt~pc~cks:(4) liquid sample in.lector; (5) platinum resistance therm~mleter;(6)
double-~,alled absorption ves.,,el;(7) magnetic slirrer: (8) mercury manonleter: (9) mercury reservoir.
C.S.O. Siha et al. / Huid Phase l-quilibria 135 (1997) 137-144
139
certified by NPL, UK); and (ii) the mercury levels in the manometer are measured with a precision
cathetometer (to 0.01 mm). After the whole apparatus has been evacuated the gas is introduced into
the system. The working pressure is adjusted to the atmospheric pressure, and a sample of the alcohol
is injected into the absorption vessel. After mechanical shaking for 10 to 15 min, the mercury is
brought to level in the three tubing branches. The change in the mercury level in the right hand tube
of the manometer is measured with a cathetometer later (at least 8 h) to make sure that equilibrium
has been attained.
The accuracy of the experimental method was checked by measuring the solubility of carbon
dioxide and of nitrous oxide in water, being found to be about 3%. The gases used in this work,
obtained from Matheson, were of the highest purity available (99.5 mol%). The alcohols were all
purchased from Riedel-de Haen with specified minimum purities, in mol%, as follows: methanol,
99.63; ethanol, 99.6; propanol, 99.97; and n-butanol, 99.97. The method used for degassing the
solvents was a combination of the techniques described by Gibbs and Van Ness (1972) and Bell et ai.
(1968). No further purifications were made.
3. Calculations
To simplify the treatment of the raw data certain assumptions are necessary: (i) the volume change
of the liquid sample during saturation is negligible; (ii) CH 3F is an ideal gas at the conditions used:
(iii) Raoult's law is valid for the solvent in the mixture; and (iv) the system follows Henry's law.
The volume of the absorbed gas (V 2) is given by:
V~ = V, - AV
(2)
where V I is the volume of the pure liquid sample injected in the absorption vessel and AV is the
measured volume change in the gas alter equilibrium has been attained. The quantity (in mol) of gas
absorbed in the liquid is obtained by
n2 -
P2V2
RT
(3)
where
P~=P-(I
- x 2)P,"
(4)
and P represents the equilibrium pressure, T the equilibrium temperature, and x~ the solubility, in
mole fraction, of the gas in the liquid solution. P~* is the vapour pressure of the pure solvent
calculated using Antoine equation with the constants given in Reid et al. (1987). All the solubilities
were corrected to 1 atm partial pressure using Henry's law. The Ostwald coefficients L2. ~, have been
determined at the experimental equilibrium pressure P = 1 atm, using the following expression:
L2,1 = V 2 / V I
(5)
The calculation of the ideal solubilities of CH3F to find the H-bonding-factors, Eq. (1), was made
using Wagner's vapour pressure equation (Fonseca and Lobo, 1994). The dependence of the solubility
of CH ~F on temperature ( l , in K) has been represented by the equation:
Rlnx~ = A + B / T + ClnT
with the parameters fitted to the data by a least-squares method.
(6)
C.S.O. Siha et al. / kluid I'hase Equilihria 135 ¢19971 I,¢7- 144
140
Table 1
Solubility of C H s F in some alcohols exprcssed as mole fraction, x, at a partial pressure P, = 101 325 Pa, Ostwald
coefficient, L_,.~ at P = 101 325 Pa and Henry coefficient, tt,~
Solvent
T (K)
r , x 10 ~
L.~.I
H.,.i (MPa)
Methanol
288.82
291.15
293.48
295.29
298.66
3(F0.21
1.535
1.228
0.843
0.559
0.312
0.116
0.979
0.796
0.557
0.375
0.216
0.082
66.(1)
82.53
120.2
18 I. I
324.9
872. I
Ethanol
283.90
285.71
288.30
290.37
292.70
293.74
294.52
297.1 I
298.92
301.51
2.076
1.981
1.999
1.948
2.243
2.249
2.133
1.841
1.453
I. 124
0.870
0.837
0.856
0.843
0.984
0.993
0.946
0.830
0.663
0.523
48.80
5 I. 14
50.70
52.(12
45.18
45.05
47.5 I
55.05
69.74
90.14
Propanol
284.94
286.49
288.55
291.41
292.45
293.48
296.07
298.66
300.21
3.038
2.742
2.819
2.778
2.927
2.883
2.893
2.663
2.622
0.975
0.884
0.915
0.911
0.96-t0.953
0.965
0.898
0.890
33.35
36.95
35.94
36.47
34.61
35.14
35.02
38.05
38.64
n-Butanol
281.07
283.38
286.23
288.56
291.41
292.7(I
293.22
295.81
298.66
300.47
3.136
3.567
3.176
3.309
3.386
3.724
3.537
3.115
2.932
2.903
0.807
0.925
0.831
0.872
0.901
0.945
(I.947
0.841
0.799
0.797
32.31
28.40
31.91
3(I.62
29.92
27.2 I
28.65
32.53
34.56
34.91
The
thermodynamic
functions
of
the
solution
havc
becn
obtained
from
Eq.
(6)
by
standard
expressions:
_
AH~'=
AS~'=
RT[ Olnx2 ]
[ 0-~nTlr ' = -B
[~ O--7~7-fnT] p , + I n x
+ CT
z
=(a
(7)
+C)
+ C'lnT
(8)
141
C.S.O. Silt a et al. / Fluid Phase Equilibria 135 (1997) 137 144
10"2
i0 "3
,2
10"4
F zhar~l
10-5
2 44
2 45
2 46
2 47
2 48
h~KT/K)
Fig. 2. S o l u b i l i t y o f m e t h y l l l u o r i d e , x 2, in m o l e f r a c t i o n in v a r i o u s a l c o h o l s .
and
AG o = A H ~ ) - TASk)
(9)
4. Results and discussion
The solubilities of methyl fluoride in the alcohols are reported in terms of the Ostwald coefficient,
the Henry coefficient, and of mole fraction in Table i. Measurements made at 303 K for the solubility
of CH3F in methanol, ethanol and propanol were not included in the table since at this temperature
the values obtained for the solubilities in the three solvents are out of the reliability range of the
apparatus.
The solubility data are plotted in Fig. 2, showing that among the solvents the solubility of CH3F is
the lowest in methanol, increasing with the C-content in the alcohol. This can be explained assuming
that H-bonding in the solvent has the effect of 'excluding' solute molecules and hence of reducing the
solubility below what it would be without such molecular interaction (Hayduk and Laudie, 1973).
This figure also shows the presence of a maximum in the value of solubility at about 293 K.
To represent the temperature dependence of the mole fraction solubilities Eq. (6) was fitted to the
corrected x 2 values. The optimized parameters of Eq. (6) and the percentage standard deviation of
x 2, or(= ( I / M ) [ ~ ] x 2 ( e x p ) - x e ( c a l c ) ] / x 2 ( e x p ) × 100] where M is the number of experimental
points), are listed in Table 2.
Table 2
P a r a m e t e r s in the e q u a t i o n R I n x~ = A + B ~ T + C I n T
System
A(JK
CH
CH
CH
CH
146 187.8
4 5 413.1
5030.9
13 6 1 9 . 5
3l ' : / M e t h a n o l
~F / E t h a n o l
3F/Propanol
:~F/n-Butanol
J mol-I)
B×lO~(Jmol- 6302.640
- 1975.268
- 219.230
- 592. I (X)
i)
C'(JK
i mol-I)
- 21 9 5 9 . 7 5
- 6817.18
- 762.55
- 2050.283
tr ( % )
3
7
3
4
142
C.S.O. Sib : ('t cd. / I " h , d I'ho'.(. t-quilil?,.'io 135 (1997) 1.¢7- 144
Table 3
Molar Gibbs energy of solution AG(~. enthalpy of solution ..3tt~, ) and entrop'~
°
. of solution ..IS 2,
at
290;
K and
I aim partial
p r e s s u r e o f C H ~1:
System
..IH~° (J tool ) )
CH 0:/Methanol
- 244659
CH 3F/Ethanol
CH 3F/Propanol
CH 3F/n-Butanol
CH3F/H20
"
A,½"~) (J l]'lol I K I)
20620
-- 5 7 2 7 4
- 245.62
15 95F;
- 8124
- 76.35
- 19 192
(-
AG~,) (J mol i )
- 889.75
14 6 4 0
-112.49
10; 1 2 9 )
(-
14347
117.74)
(16975)
"' Values given by Wilhelm ct al. (1977).
The experimental Gibbs energies, enthalpics and entropies of solution at 298 K. calculated from
Eqs. ( 7 ) - ( 9 ) for the systcrns studied in this work are recorded in Table 3 together with those for the
C H 3 F / H ~ _ O system (Wilhelm et al., 1977). The solubility of C H 3 F in water at 298 K is greater than
that in methanol but less than the solubilities in the other alcohols; the same happens lk}r the
corresponding values o f AG(., ).
Fig. 3 represents the or-factors for ethanol, propanol and n-butanol as a function of O<,,,~,h~,,,,j for
C H 3 F and lk)r other gases for which solubility data have been selected from the literature (Gerrard
and Fogg, 1991; Boyer and Bircher, 1960). There is a linear relation between hydrogen-bonding
factors in methanol and those for the other alcohols. Although it is difficult to predict solubilities
accurately from graphs such as that in Fig. 3, the relation is at least semiquantitative.
Wilhelm and Battino (1973) suggested that solubilities of gases in a particular solvent can be
correlated by representing log .r, as a function of ( E / k ) , the Lennard-Jones energy parameter for the
1o !
t
)
I•
:@
.-Butanol
•
tithatlul
I%panol
!
I
&
OA
•
I"!
,.:tJ
001 [~11 ! .
o oI
¢, H4
0 1
:,{':
c:ll:
~.]1,
H: ~
I
10
GI
me,hazel
Fig. 3. Hydrogen-bonding factors ~x, for ethanol, propanol and n-butanol as
methanol, {.Vn,,.th,,n,,i,at 298 K.
a |'unction of
hydrogen-bonding factor for
C.S.O. Sih aet al. / t"luid Phase Equilibria 135 ~1997J 137-144
143
-o 5
C
-1
HS
"~ -15 . \ 2
,,
C
-25
2(~)
250
300
340
e/~ (K)
Fig. 4. Solubilityof some gases in methanol, x~, as a functiono f the Lennard-Jonesenergy parameter E/k, at 298 K.
solute. Using values of the energy parameter given by Casparian and Cole (1974) for CH3F and by
Mourits and Rummens (1977) for the other gases, we have plotted in Fig. 4 the solubilities of several
polar gases in methanol. This plot shows that such a correlation can be established, at least in
principle.
5. Conclusions
The solubility of methyl fluoride in methanol, ethanol, propanol, and n-butanol increases as the
C-content of the alcohol increases. This is related to H-bonding or association in the alcohol; i.e.
solvents with smmg H-bonding tendencies dissolve less of the same gas than those with weaker
H-bonding tendencies.
The results show that there is a semiquantitative correlation between H-bonding factors of
methanol and those of the other associated solvents studied.
6. List of symbols
A
B
C
H2.i
L2.t
I12
P
Pl"
P~
parameter in Eq. (6)
parameter in Eq. (6)
parameter in Eq. (6)
Henry coefficient
Ostwald coefficient
amount of substance of gas absorbed in liquid (mol)
pressure
vapour pressure of the pure solvent
partial pressure of the solute gas
144
R
T
Vi
V,
.It" .:,
xTa
(.\S.O. Sib a at al. /1"Tuid Phase Equilihria I.¢5 ( 19971 137-144
gas constant (J mol ~ K ~)
temperature (K)
volume of the pure liquid (cm ~)
volume of the absorbed gas (cm ~)
mole fraction solubility
ideal solubility
6.1. Greek letters
AGO
AH~ )
AS(, )
AV
O"
H-bonding factor in the solvent
molar Gibbs energy of solution (J tool t)
molar enthalpy of solution (J tool ~)
molar entropy of solution (J tool ~ K ~)
measured volume change in the gas (cm ~)
standard deviation of x, (%)
References
Bell, T.N., Cussler, E.L., Harris. K.R., Pepela. C.N.. l)unlop. P.J.. 1968. An apparatus tor degassing liquids by vacuum
sublimation. J. Phys. Chem. 72. 4693--4695.
Bo. S.. Battino. R.. Wilhelm, E., 1993. Solubility of gases m liquids. 19. Solubility of Fie, Ne. Ar. Kr. Xe. N,, O~. CH~,
CI~ and SF¢, in nomlal 1-alkanols n-Cil-lel ~ ~()H (I <_ I _< II) at 298.15 K. J Chem. Eng. Data 38, 611-616.
Boyer, F.L.. Bircher, L.J., 1960. The solubility of nitrogen, argon, methane, ethylene and ethane m normal primary alcohols.
J. Phys. Chem. 64, 1330-133 I.
Casparian, A.S., Cole. R.H., 1974. Viscosities of polar ga,,es by relaxation of capillary tlow. J. Chem. Phys. 60, 1106-1109.
Fonseca, I.M.A., Lobo, LQ.. 1994. The triple point of methyl fluoride. J. Chem. Thermodyn. 26, 671-672.
Gerrard, W., Fogg, P.G.T., 1991. Solubility ot +Gases in l,iquids. Wiley, New York.
Gibbs, R.E., Van Ness, H.C., 1972. Vapor-liqt.id equilibria from total-pressure measurements. A new apparatus. Ind. Eng.
Chem. Fundam. I I, 410-413.
Hayduk, W., Laudie, H., 1973. Solubilities of gases in water and other associated solvents. AIChE J. 19. 1233-1238.
Kaszonyi, A., Harustiak, M., Kizlink, J., 1992. Solubility of acetylene in vinyl acetate and in a mixture of vinyl acetate and
acetic acid. J. Chem. Eng. Data 37, 37-38.
Reid, R.C., Prausnitz. J.M., Poling. B.E., 1987. The Properties of Gases in l,iquids, 4th ed. McGraw-Hill, New York.
Mourits, F.M., Rummens, F,H.A., 1977. A critical evaluation of Lennard-Jones and Stockmayer [x)tential parameters and of
some correlation methods. Can. J. Chem. 55, 3007 3020.
Short, I., Sahgal, A.. Hayduk. W., 1983. Solubility of ammonia and hydrogen sulphide in several polar solvents. J. Chem.
Eng. Data 28, 63-66.
Wilhelm, E., Battino, R., 1973. Thermodynamic functions of the solubilities of gases in liquids at 25°C. Chem. Rev. 73.
1-9.
Wilhelm, E., Battino, R.. Wilcock, R.J., 1977. Lo~,-pressure solubility of gases in liquid water. Chem. Rev. 77, 219- 262.