Aqueous solubility, HenryХs law constants and air

Chemosphere 57 (2004) 1543–1551
www.elsevier.com/locate/chemosphere
Aqueous solubility, HenryÕs law constants and
air/water partition coefficients of n-octane and
two halogenated octanes
S. Sarraute, H. Delepine, M.F. Costa Gomes *, V. Majer
Laboratoire de Thermodynamique des Solutions et des Polyme`res, UMR 6003, CNRS/Univ. Blaise Pascal,
24, avenue des Landais, 63177 Aubie`re Cedex, France
Received 16 July 2003; received in revised form 19 July 2004; accepted 22 July 2004
Abstract
New data on the aqueous solubility of n-octane, 1-chlorooctane and 1-bromooctane are reported between 1 C and
45 C. HenryÕs law constants, KH, and air/water partition coefficients, KAW, were calculated by associating the measured
solubility values to vapor pressures taken from literature. The mole fraction aqueous solubility varies between (1.13–
1.60) · 107 for n-octane with a minimum at approximately 23 C, (3.99–5.07) · 107 for 1-chlorooctane increasing
monotonically with temperature and (1.60–3.44) · 107 for 1-bromooctane with a minimum near 18 C. The calculated
air–water partition coefficients increase with temperature and are two orders of magnitude lower for the halogenated
derivatives compared to octane. The precision of the results, taken as the average absolute deviations of the aqueous
solubility, the HenryÕs law constants, or the air/water partition coefficients, from appropriate smoothing equations as a
function of temperature is of 3% for n-octane and of 2% and 4% for 1-chlorooctane and 1-bromooctane, respectively. A
new apparatus based on the dynamic saturation column method was used for the solubility measurements. Test measurements with n-octane indicated the capability of measuring solubilities between 106 and 1010 in mole fraction, with
an estimated accuracy better than ±10%. A thorough thermodynamic analysis of converting measured data to air/water
partition coefficients is presented.
2004 Elsevier Ltd. All rights reserved.
Keywords: Aqueous solubility; HenryÕs law constant; Air/water partition coefficient; n-Octane; 1-Chlorooctane; 1-Bromooctane
1. Introduction
The prediction of the transport and fate of hydrophobic organic chemicals in the environment requires
*
Corresponding author. Fax: +33 4 73407185.
E-mail address: margarida.c.gomes@univ-bpclermont.fr
(M.F. Costa Gomes).
the knowledge of their physical and chemical properties.
An important thermodynamic property is the solubility
in water as it is essential in waste minimization and
waste remediation calculations. Furthermore, when
combined with vapor pressure data, it permits the calculation of the HenryÕs law constant and from there,
several partition coefficients such as the air/water partition coefficient or the octanol/water coefficient (Mackay,
1991; Boethling and Mackay, 2000).
0045-6535/$ - see front matter 2004 Elsevier Ltd. All rights reserved.
doi:10.1016/j.chemosphere.2004.07.046
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S. Sarraute et al. / Chemosphere 57 (2004) 1543–1551
The solubility of a pollutant in water constitutes
per se a physical property of capital importance for environmental processes. It can affect adsorption and
desorption in soils and volatility of chemicals from
aquatic systems and it has also impact in possible transformations by hydrolysis, photolysis, oxidation, reduction and biodegradation. Other specialized pathways,
such as washout from atmosphere by rain, are also influenced by the extent of solubility (Reza et al., 1999).
The air–water partition coefficient, KAW, relating
atmospheric and aqueous concentrations of a substance
is in direct relation with HenryÕs law constant, KH,
which can be calculated using solubility data. Apart
from their need in the construction of transport models
that attempt to describe the movement of pollutants in
the environment, these data are also required in a number of technological applications (e.g. air-stripping processes for the renovation of waters contaminated by
organic solvents (Gosset, 1987)).
Different methods, extensively reviewed by Mackay
and Shiu (1981) were described for the determination
of HenryÕs law constants. More recently, several other
techniques have been proposed and their respective merits and deficiencies discussed (Sanemasa et al., 1982;
Gosset, 1987; Dohnal and Hovorka, 1999).
Halogenated paraffins, in particular chlorinated alkanes are used as plasticizers, lubricants and flame retardants (Sijm and Sinnige, 1995). Very little information is
available on the physical–chemical properties of these
persistent hydrophobic chemicals such as aqueous solubility, octanol/water partition coefficient, HenryÕs law
constant, relative bioavailability, while more information is available on their acute toxicity and bioaccumulation (Drouillard et al., 1998a). The complexity of
technical mixtures of chlorinated paraffins has precluded
the determination of physical properties of individual
compounds or of simpler synthesized mixtures to determine key physical properties that may be used in assessing the behavior of halogenated paraffins in the
environment (Drouillard et al., 1998a), namely as a prerequisite for many thermodynamics-based environmental fate models in use today (Drouillard et al., 1998b).
In this work, the solubilities of n-octane, 1-chlorooctane and 1-bromooctane in water were determined
experimentally as a function of temperature using a dynamic saturation column method. The data are associated with the pure solute vapor pressures for the
calculation of the HenryÕs law constant and air/water
partition coefficient. The thermodynamic background
and simplifying assumptions used in the conversion of
the experimental solubility data and in the relevant coefficients are discussed and possible sources of error are
outlined.
Several data have been reported in the literature concerning the aqueous solubility of n-octane (Polak and
Lu, 1953; McAuliffe, 1963, 1966; Price, 1976; Jo¨nsson
et al., 1982; Tewari et al., 1982; Heidman et al., 1985;
Reza et al., 1999; Miller and Hawthorne, 2000) so this
system was chosen to test the experimental technique
and also as a reference for examining the effect of the
halogen atoms on the aliphatic chain. For 1-chlorooctane no experimental values for the aqueous solubility
were found in the literature. For 1-bromooctane two
identical experimental values for the aqueous solubility
at 25 C, determined using a generator column procedure, are reported independently by Tewari et al.
(1982) and Miller and Wasik (1985). These results are
presented as a reliable recommended value (xs =
1.56 · 107) in the compilation by Horvath and Getzen
(1999) and are also referred in the review paper of Abraham (1984) on solution thermodynamics of organic
compounds.
Values for the air/water partition coefficients (the so
called dimensionless HenryÕs law constants) of the three
studied solutes are also reported by Yaffe et al. (2003)
for testing several prediction methods. We could not,
however, trace the direct experimental determinations
of the aqueous solubility or other properties that served
as a basis for the calculation of these data.
2. Experimental section
2.1. Materials
All the solutes used were obtained from Acros
Organics with stated purities of at least 99 mol% and
were used with no further purification. Distilled water
was used as solvent.
2.2. Apparatus and operation
The experimental technique used in this work is
based on a dynamic saturation column method and, in
its essence, was already described in the literature for
use at ambient (May and Wasik, 1978; DeVoe et al.,
1981; Owens et al., 1986) and superambient conditions
(Miller and Hawthorne, 1998; Bergin, 2002). It is based
on the saturation of a known quantity of water flowing
through a saturation cell which is filled with an inert stationary phase impregnated with the organic solute. The
solute is subsequently trapped in a specific extraction
column and finally quantitatively removed by an appropriate solvent. This method was originally proposed for
solids and later adapted for liquid solutes (Owens et al.,
1986). The analysis of the final solution are carried out
either using liquid (May and Wasik, 1978; DeVoe
et al., 1981; Owens et al., 1986) or gas chromatography
(Miller and Hawthorne, 1998; Bergin, 2002).
The apparatus used is schematically represented in
Fig. 1. The heated components are kept inside a water
bath maintained at constant temperature to within
S. Sarraute et al. / Chemosphere 57 (2004) 1543–1551
1545
Fig. 1. Solubility apparatus used in this work. VP, isocratic pump; M, manometer; SC, saturation cell preceded by a pre-heating coil;
EC, extraction cell; T, liquid thermostat.
±0.05 C by means of a PID temperature controller
(TRONAC Inc., model PTC-40). A pump operating
in a constant-flow mode (Gilson, model 306 or Perkin
Elmer, model 410 were employed successively) was used
during the measurements to supply a regular stream of
water required for the solubility determinations. The solvent is pumped during 10-12 h, at flow rates between 0.1
and 0.2 ml/min, through a pre-heating coil (of approximately 2 ml), into the saturation cell (a 2.0 ml, 0.46 cm
i.d. · 10 cm analytical column from Modulo-cart) containing the solute immobilized in a dry and clean solid
stationary phase (Gaz chrom R 60/80 by Alltech). The
saturated solution then flows through an inverse phase
analytical column (C18 5lm, 0.46 cm i.d. · 10 cm, from
Kromacil) where the solute is quantitatively retained.
The pressure, read at the pump, varies from 40 to
70 bar (and decreases 20–30 bar along the saturation
cell), depending on the temperature and on the degree
of the column packing (different fillings of the saturation
cell corresponding to different pressure drops in relation
with the packing of the solid sorbent). The organic compound is subsequently removed by back flushing this
extraction column with a good solvent (5-15 ml of methanol in our case). The quantity of water passed through
the saturation (and extraction) cell, typically 100–200 ml,
is determined gravimetrically.
The analyses of the final solutions were performed
using a Shimadzu model CG14A gas chromatograph
equipped with FID and a split/splitless injector. Chromatographic separations were accomplished with a 30 m
capillary column (from SGE, 0.32 mm i.d., 0.25 lm film
thickness). The quantification of the studied organic
compounds was carried out by an internal standard technique using n-nonane, n-decane and n-dodecane for the
analysis of n-octane and of the chloro and bromo octane,
respectively.
To determine whether saturation of the water was
achieved and to verify that the solute was not mechanically carried away out of the saturation cell, determinations were performed at different flow rates (0.1–0.3 ml/
min) for all the solutes studied. Identical results were
obtained for each flow rate, demonstrating that the
residence time of the solvent in the saturation cell
was sufficient to saturate the water with the organic
solutes. Two extraction columns of different lengths
(5 and 10 cm) were used in order to test if the solutes
are quantitatively captured. The quantity of methanol
used during the back flush of the column was also
optimised to assure the complete extraction of the organic solute.
2.3. Thermodynamic data analysis
The solubility of the organic solutes in water, expressed as molar fraction of the solute or in quantity
of solute per volume of solvent, is calculated from the
experimental quantities: the amount of water passed
through the saturation column and the amount of solute
retained in the specific extraction column, determined by
gas chromatography. As mentioned above, the resistance offered by the packing of the saturation cell implies
that the saturation of the water stream by a solute is
done at an elevated pressure, which decreases along
the saturation cell. Since the solubility data of interest
to environmental studies is close to ambient pressure
(p0 = 1 bar) it will be examined to what extent the measured values of aqueous solubility and HenryÕs law constant are affected by the pressure attained during the
experiments. To our knowledge, this issue has not been
addressed by earlier investigators who have also used the
saturation column method.
A detailed discussion of the thermodynamic relations
used in the data analysis presented below can be found
in reference text books of chemical thermodynamics
(Smith et al., 1996; Prausnitz et al., 1999; Sandler,
1999). Specifically, the thermodynamic treatment of
the HenryÕs law constant in relation with the aqueous
solubility of organic compounds was recently addressed
by Sedlbauer et al. (2002).
The dependence of the mole fraction solubility xsol
s
with pressure can be approximated by a differential
equation:
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S. Sarraute et al. / Chemosphere 57 (2004) 1543–1551
o ln xsol
V 1 V s
s
ffi s
op
RT
T
ð1Þ
where V 1
s and V s are the partial molar volume of the
solute in water at infinite dilution and the molar volume
of the pure solute, respectively. This equation is exact in
the limit of infinite dilution where the HenryÕs activity
coefficient equals unity and no water is dissolved in the
organic phase. In the temperature range of the present
experiments, both V 1
s and V s do not vary significantly
with pressure and the relative error dr xsol
(expressed in
s
percentage), due to neglecting the pressure dependence
of the mole fraction solubility, can be estimated as:
dr xsol
s ¼
dxsol
s
100
sol
xs ðp0 Þ
ðV 1 V s Þðp p0 Þ
1 100
ffi exp s
RT
ð2Þ
0
0
where xsol
s ðp Þ is the mole fraction solubility at p = 1 bar.
The difference between the partial molar volume at
infinite dilution and the molar volume of a pure liquid
hydrophobic solute is generally negative near 25 C.
We suppose, by analogy with accurate data available
for benzene and toluene (Degranges, 1998) (for which
1
3
V1
s V s ¼ 10 cm mol ), that for the solutes treated,
1
V s does not exceed –15% of V s and that the pressure attained at the end of the column, will never surpass
40 bar. The systematic error in the solubility determination, due to the pressure during the measurements, will
then be positive and less than 3% of the mole fraction.
The exact thermodynamic definition of the HenryÕs
law constant, KH, is:
fs ðp; T ; xs Þ
ð3Þ
K H ðp; T Þ ¼ lim
xs !0
xs
where fs is the fugacity of the organic solute in the aqueous phase and xs is the corresponding mole fraction.
HenryÕs law constant has always dimension of pressure;
it should not be considered as a coefficient characterizing
phase equilibrium but rather as a proportionality constant between the fugacity of a solute and its concentration in the same phase in the limit of infinite dilution.
For hydrophobic substances, sparingly soluble in
water, the limit in Eq. (3) can be approximated by the
ratio fs =xsol
s . In this case, the fugacity of the solute in
the aqueous phase is equal to that in the organic phase
as we are in a situation of thermodynamic equilibrium.
For the organic liquids investigated in this study, the
fugacity of the solute in the organic phase can be
approximated by the fugacity of the pure solute, fs ,
which is then calculated in the usual way
V s ðp psat
sat sat
sol
s Þ
fs ðp; T ; xs Þ ffi fs ðp; T Þ ffi /s ps exp
ð4Þ
RT
where psat
s and V s are the vapor pressure and the molar
volume of the pure solute, respectively. The exponential
term in Eq. (4) is the so called Poynting correction which
expresses the change in fugacity due to a variation of
pressure between psat
s and p (e.g. the pressure of the solubility measurement), provided that V s can be considered constant. The fugacity coefficient of the solute is
obtained from (at temperatures below or around the
normal boiling point of the solute):
sat Bs ps
ð5Þ
/sat
¼
exp
s
RT
where Bs is the second virial coefficient for the pure
solute.
The HenryÕs law constant at the experimental pressure can then be calculated from the measured mole
fraction solubilities, xsol
s , and the thermodynamic properties of the pure solute:
V s ðp psat
s Þ
sat
/sat
s p s exp
RT
ð6Þ
K H ðp; T Þ ffi
xsol
s
HenryÕs law constant is in direct relation with the
1
ig 0
Gibbs free energy of hydration DG1
hyd ¼ Gs Gs ðp Þ ¼
0
RT lnðK H =p Þ which corresponds to the difference in
the Gibbs free energy of the solute in the standard state
of infinite dilution and as an ideal gas at the pressure p0.
The pressure dependence of KH is then expressed by
o ln K H
V1
ð7Þ
¼ s
op
RT
T
Assuming that the partial molar volume at infinite dilution of the solute, V 1
s , is independent of pressure, the
HenryÕs law constant at the reference pressure p0 can
be calculated by
V s p psat
s
sat
/sat
p
exp
s
s
RT
1
ð8Þ
K H ðp0 ; T Þ ffi
V
ð
p
p0 Þ
s
xsol
exp
s
RT
For the temperatures studied, the vapor pressure of the
solutes is low (<0.1 bar), the fugacity coefficient is close
to unity and the HenryÕs law constant can be approximated as
K H ðp0 ; T Þ ffi
psat
s
xsol
s
ð9Þ
the comparison of Eqs. (2) and (8)
Since p > p0 psat
s
suggests that neglecting both exponential terms in Eq.
(8) introduces an uncertainty similar to that estimated
above for the solubility. This systematic error is, however, of opposite sign, i.e. the HenryÕs law constant calculated using Eq. (9) will be lower by less than 3% due to
the relatively high pressure in the saturation cell.
HenryÕs law constant is often used to express the air–
water partition coefficient which is defined as (Mackay,
1991)
S. Sarraute et al. / Chemosphere 57 (2004) 1543–1551
K AW ¼ lim
aq
C s !0
C air
s
C aq
s
ð10Þ
i.e. the limiting ratio of the concentration of a solute, expressed in molarity, in the atmospheric and aqueous
phases. In the majority of the conditions relevant to
environmental studies, the two coexisting phases are in
equilibrium, the organic solute is very dilute in the aqueous phase and air can be considered as an ideal gas (the
fugacity of the solute is then equal to its partial pressure
ps). In these conditions,
sol
C air
s ffi p s =RT ffi K H xs =RT
C aq
s ffi
xsol
s
V w
ð11Þ
The limit in Eq. (10) can then be replaced by the ratio of
concentrations expressed in molarity leading to the correct relation between the air/water partition coefficient
and HenryÕs law constant:
K AW ¼
K HM w
RT qw
ð12Þ
where KH is the HenryÕs law constant expressed in Pa, R
is the gas constant expressed in J mol1 K1, Mw and qw
are the molar mass of water in kg mol1 and its density
in kg m3, respectively. The uncertainty in the calculated
KAW will be similar to that estimated for KH as the properties of pure water are very well known.
Considering the thermodynamic treatment presented
and the significance of the HenryÕs law constant which
cannot be considered as a partition coefficient, the
authors recommend that the air/water partition coefficient should not be associated with the term HenryÕs
law constant.
3. Results and discussion
The present experimental method was tested by
measuring the solubility of n-octane in water from 10
to 45 C. The experimental results, expressed in molarity
and mole fraction, are recorded in Table 1 together with
the calculated HenryÕs law constants and air–water partition coefficients. For the vapor pressure of n-octane
was used the recommended value of Ruzicka and Majer
(1994), the relative atomic masses were taken from the
IUPAC tables (IUPAC, 1995) and the orthobaric properties of water from the Wagner and Pruss equation of
state (Wagner and Pruss, 1993).
The mole fraction aqueous solubilities obtained in
this work for n-octane are represented in Fig. 2 together
with literature data reported in the same temperature
range. The values published by different authors exhibit
a large scatter, typical of the difficulties encountered in
the measurement of such small values of solubility. A
1547
solubility minimum is observed as expected for an alkane of medium molar mass. The exact value of the temperature corresponding to that minimum is, however,
difficult to locate due to the scatter of the experimental
values. From the analysis of all these data as well as a
careful study of the sources of the systematic errors during our experiments, it is believed that the values for the
mole fraction aqueous solubility reported in this work
are accurate to within 10%. The results obtained for
the aqueous solubility of 1-chlorooctane and 1-bromooctane between 5 and 40 C and 1 and 40 C, respectively, are depicted in Fig. 3 and recorded in Table 1
together with the calculated HenryÕs law coefficients
and air–water partition coefficients.
The vapor pressures for the halogenated alkanes necessary for the calculation of the HenryÕs law constants,
were obtained from the Antoine equation (Smith et al.,
1996) with parameters taken from the work of Li and
Rossini (1961) (which reports Antoine constants above
65 C for 1-chlorooctane and above 70 C for 1-bromooctane). The validity of the extrapolation of the Antoine
equation towards lower temperatures was verified by
comparing the enthalpies at 25 C calculated from the
Clapeyron equation with the ones determined from
calorimetric experiments (Majer and Svoboda, 1985).
It was observed that the derived enthalpies of vaporization did not differ more than 1.7% from the experimental
values which is considered as satisfactory.
The temperature dependence of the aqueous solubility of the three organic solutes was fitted to polynomial
equations of the general form:
ln xsol
s ¼
n
X
Ai ðT Þi
ð13Þ
i¼0
Several empirical methods were reported in the literature
to represent the temperature dependence of the HenryÕs
law constant. In the present case, KH is fitted to a linear
function of the inverse temperature (Krause and Benson, 1989):
ln K H ¼
n
X
Bi ð1=T Þi
ð14Þ
i¼0
and equivalently for the air/water partition coefficient:
ln K AW ¼
n
X
C i ð1=T Þi
ð15Þ
i¼0
The coefficients Ai, Bi and Ci for Eqs. (13)–(15) as well as
the average absolute per-cent deviations of the fit, are
listed in Table 2. Three parameter equations were used
as is current practice in the literature. It is observed that
the errors associated with the parameters of the fit never
exceed the absolute value of the parameter. Furthermore, the correlation factors are always acceptable.
The use of more than three parameters in such limited
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S. Sarraute et al. / Chemosphere 57 (2004) 1543–1551
Table 1
Aqueous solubility (expressed in molarity and mole fraction), vapor pressure, HenryÕs law constant and air–water partition coefficient
of n-octane, chlorooctane and bromooctane
T/C
n-Octane
10.3
24.8
24.9
24.9
24.9
24.9
24.9
24.9
30.1
34.9
35.0
39.9
39.9
44.7
44.8
44.8
Cs/106 mol dm3
7.55
6.87
6.84
7.31
7.14
6.99
7.20
7.49
6.24
8.13
7.34
8.04
8.20
8.67
8.64
8.79
7
xsol
s =10
psat/Paa
KH/108 Pa
KAW
1.36
1.24
1.24
1.32
1.29
1.26
1.30
1.35
1.13
1.47
1.33
1.46
1.49
1.58
1.57
1.60
762.0
1841
1846
1846
1846
1846
1846
1846
2472
3185
3210
4114
4114
5224
5237
5249
56.0
148
149
140
143
146
142
136
219
216
241
282
276
331
333
328
42.8
108
109
102
104
107
103
99.5
157
153
171
197
193
228
229
226
1-Chlorooctane
5.0
9.9
9.9
9.9
10.0
19.1
25.0
25.1
25.2
30.0
34.8
35.1
35.1
40.0
22.7
23.7
23.5
22.6
22.1
25.4
24.7
26.2
25.8
27.1
27.0
26.9
28.0
27.5
4.09
4.27
4.23
4.08
3.99
4.59
4.46
4.73
4.67
4.90
4.90
4.87
5.07
4.99
26.30
39.55
39.71
39.71
40.04
81.99
126.6
127.1
128.0
179.8
248.4
252.5
252.5
347.5
0.51
0.75
0.76
0.79
0.82
1.51
2.46
2.33
2.37
3.23
4.53
4.63
4.45
6.29
0.40
0.58
0.58
0.61
0.62
1.12
1.79
1.70
1.73
2.32
3.20
3.28
3.15
4.39
1-Bromooctane
1.1
5.0
5.2
9.9
9.9
14.9
24.9
29.9
34.9
40.0
40.1
13.8
11.5
12.5
9.11
9.18
9.41
8.88
11.3
14.4
18.9
18.0
2.48
2.07
2.26
1.64
1.66
1.70
1.60
2.05
2.60
3.44
3.27
6.560
9.431
9.606
14.68
14.68
22.60
50.53
73.61
105.5
150.1
150.6
0.27
0.45
0.43
0.89
0.89
1.33
3.15
3.59
4.06
4.37
4.60
0.21
0.35
0.33
0.68
0.68
1.00
2.30
2.58
2.87
3.05
3.21
a
From Ruzicka and Majer (1994) for n-octane and from Li and Rossini (1961) for 1-chlorooctane and 1-bromooctane. The values
in italics are extrapolations using Antoine equation as explained in the text.
temperature range leads always to overfitting without
decreasing significantly the average absolute deviation
of the fit. Consequently, very large errors in the parameters of Eqs. (13)–(15) were observed. The average
absolute deviations listed characterize the precision of
the results obtained and vary from 2% to 4% for the
three solutes studied.
By taking the temperature derivative of Eq. (9), one
obtains:
R
d ln K H
d ln xs
dp
¼ RT 2
RT 2 sat
dð1=T Þ
dT
dT
ð16Þ
When the vapor phase is considered as ideal, this equation is approximately equivalent to
S. Sarraute et al. / Chemosphere 57 (2004) 1543–1551
2.0
1549
6
1.5
x /10-7
xoctane /10-7
5
4
3
1.0
2
0.5
10
20
30
40
50
1
0
T /°C
20
30
40
50
T /°C
Fig. 2. Mole fraction solubility of n-octane in water. (d) Mole
fraction solubility, this work; other symbols stand for mole
fraction solubilities from literature: (h) McAuliffe (1966); (,)
Polak and Lu (1953); () Price (1976); ( ) Tewari et al. (1982);
(n) Jo¨nsson et al. (1982); (m) Heidman et al. (1985); (.) Reza
et al. (1999); () Miller and Hawthorne (2000).
Dhyd H 1 ¼ Dsol H 1 Dvap H ig
10
ð17Þ
where the term on the left hand side is the enthalpy of
hydration which corresponds to the transfer of the solute from an ideal gas phase to the infinitely dilute solution at the same temperature. The term DsolH1 concerns
the transfer of the pure solute to the infinitely dilute
solution and DvapHig is the enthalpy of vaporization of
the pure solute to an ideal gas state. By taking the temperature derivatives of Eqs. (13) and (14), the values ob-
Fig. 3. Mole fraction aqueous solubilities. (j) Chlorooctane,
this work; (d) bromooctane, this work; (s) bromooctane,
Horvath and Getzen (1999).
tained for the solubility and the HenryÕs law constant
can be confronted by means of Eq. (17) with the tabulated calorimetric values of DvapHig at 25 C (Majer
and Svoboda, 1985). This comparison is possible because, as shown above, the enthalpy of vaporization calculated from vapor pressures using the Clapeyron
equation is close to the calorimetric value. Thus, by
using Eq. (17) the validity of the temperature correlation
of the present values of KH and xs are tested. This is
shown in Table 3 where the last column represents the
percentage deviation of DvapHig calculated from Eq.
(17) from the calorimetric values. It is apparent that in
Table 2
Coefficients Ai for Eq. (13) Bi for Eq. (14) and Ci for Eq. (15), for mole fraction solubility, HenryÕs law constants and air/water partition
coefficients, respectively, and average per cent absolute deviations (AADa) from fits
n-Octane
A0
A1
A2
AAD
+25.526
2.7975 · 101
+4.7251 · 104
3.1%
B0
B1
B2
AAD
7.0475
+2.2996 · 104
4.1490 · 106
3.3%
C0
C1
C2
AAD
29.139
+2.4687 · 104
4.3548 · 106
3.1%
1-Chlorooctane
A0
A1
A2
AAD
21.845
+4.2442 · 102
6.0619 · 105
2.1%
B0
B1
B2
AAD
+33.573
2.3949 · 103
5.5673 · 105
2.1%
C0
C1
C2
AAD
+13.927
2.1286 · 103
5.5538 · 105
2.1%
1-Bromooctane
A0
A1
A2
AAD
+122.13
9.4611 · 101
1.6240 · 103
4.4%
B0
B1
B2
AAD
110.56
+8.2195 · 104
1.2945107
4.4%
C0
C1
C2
AAD
130.20
+8.2456 · 104
1.2943 · 107
4.3%
a
Calculated as
P
AAD ¼
jðX exp X calc =X exp j
n
where Xexp is the experimental value, Xcalc is the calculated number and n is the number of points fitted.
1550
S. Sarraute et al. / Chemosphere 57 (2004) 1543–1551
Table 3
Enthalpy of hydration, DhydH1, enthalpy of solution, DsolH1 and ideal enthalpy of vaporization, DvapHig of the three solutes
Substance
DhydH1/kJ mol1
DsolH1/kJ mol1
DvapHig/kJ mol1a
Deviation/%
n-Octane
1-Chlorooctane
1-Bromooctane
40.2
51.0
38.6
+1.5
+4.7
+16.5
+41.5
+52.4
+55.8
+0.5
+6.2
1.3
The deviation concerns the difference between the values of the enthalpy of vaporization calculated in this work (Eq. (17)) and the
tabulated calorimetric values from the literature.
a
Calorimetric values from Majer and Svoboda (1985).
two cases (n-octane and bromooctane) the agreement is
excellent and only in the case of chlorooctane the difference is more important but still considered acceptable.
The aqueous solubility of C8H17Cl is circa four times
higher than that of n-octane and exhibits a monotonous
behavior with temperature in the range studied. This results seems to confirm the tendency observed for the
aqueous solubility of chloroalkanes. From the data reported by Horvath and Getzen (1999), the aqueous solubility of 1-chloropentane still exhibits a minimum
around 15 C that disappears for temperatures ranging
from 3 to 25 C for 1-chlorohexane. No value for
the aqueous solubility of chlorooctane was found in
the literature for direct comparison.
The solubility of C8H17Br in water is higher than that
of n-octane (aqueous mole fraction solubilities half of
those for C8H17Cl) and exhibits a minimum at lower
temperature compared to n-octane. The only value
found in the literature for comparison, at 25 C (Horvath and Getzen, 1999), is in reasonable agreement with
the result obtained from Eq. (13) with parameters from
Table 2 (our value is 9% higher). It is also possible to
compare the values of KAW calculated from Eq. (15) at
25 C with the air/water partition coefficients reported
for the three solutes in the data base of Yaffe et al.
(2003). The literature values of KAW are for n-octane
120, for chlorooctane 1.55 and for bromooctane 2.40.
Our values (see below) are about 12% lower for n-octane
and for bromooctane and 10% higher for chlorooctane.
These differences do not express, however, the quality of
our values due to uncertain origin of the literature data.
values show large scatter, specially around the minimum
of solubility, the overall agreement with our results is
satisfactory. We conclude that the experimental values
for the aqueous solubility and related coefficients (KH
and KAW) exhibit an overall uncertainty below 10%.
The equations used for data fitting represent xs, KH
and KAW within this error range.
According to Eq. (13) with the parameters in Table 2,
the aqueous solubility exhibits minimum values for n-octane at 23 C (xs = 1.27 · 107) and for 1-bromooctane
at 18 C (xs = 1.57 · 107). It is estimated that the uncertainty associated with the position of the minimum in
solubility is of ±1 C, it can be somewhat larger in the
case of n-octane for which only scanty data are available
at temperatures below 25 C. In the case of chlorooctane, the solubility is four times higher than n-octane
and varies monotonously with temperature in the range
covered. The values at 25 C are as follows: for n-octane,
xs = 1.27 · 107, KH = 147 · 108 Pa and KAW: 107;
for bromooctane, xs = 1.70 · 107, KH = 2.97 · 108 Pa,
KAW = 2.15; and for chlorooctane xs = 4.67 · 107,
KH = 2.35 · 108 Pa and KAW = 1.72.
It is believed that the original experimental data reported in this work improve knowledge about the air–
water partitioning of halocarbons, substances considered as environmental chemicals. Furthermore a rigorous and compact thermodynamic analysis of the
solubility data was performed. It was also shown how
these data can be coupled with the vapor pressures of
the pure solutes in order to calculate partition coefficients of environmental relevance.
4. Conclusions
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