UV-Vis Spectrophotometric Determination of the Dissociation Constants for Monochlorophenols in Aqueous Solution at Elevated Temperatures

Article abstract of DOI:10.1023/A:1022980614320

The dissociation constants of monochlorophenols (2-, 3-, 4-chlorophenols) were examined using direct UV-vis spectroscopy at temperatures from 25 to 175°C and at saturated vapor pressures in a high-temperature, high-pressure cell. The dissociation constant, K_a increased under experimental temperatures in the order: 2-chlorophenol, 3-chlorophenol, and 4-chlorophenol. The dissociation constant of 4-chlorophenol increased with increasing temperature under experimental conditions, while those of 2-and 3-chlorophenol reached maximum values at around 125°C, and then decreased with further increases in temperature. The slope of ?(log K)/? (1/T) was nonconstant and positive, that is, endothermic, below 150°C. The data on dissociation constants were analyzed by simultaneous regression to yield a five-term equation that described the Van't Hoff isobar. The magnitude of enthalpy ΔH increased at 25°C in the order: 3-chlorophenol, 4-chlorophenol, and 2-chlorophenol. The decrease in enthalpy at the absolute temperature was larger for 3-chlorophenol than for either 2- or 4-chlorophenol. Considering the equilibrium constant K_b for the isocoulombic reaction of monochlorophenol with OH^-, the nearly linear relationship of log K_b vs. 1/T for temperatures between 25 and 175°C indicates that the ΔC_p values for this isocoulombic reaction are low.

Full text of DOI:10.1023/A:1022980614320

C
Journal of Solution Chemistry, Vol. 32, No. 1, January 2003 ( 2003)
UV–Vis Spectrophotometric Determination of the
Dissociation Constants for Monochlorophenols in
Aqueous Solution at Elevated Temperatures
Miho Uchida1, and Akitsugu Okuwaki1
Received July 2, 2002; revised October 10, 2002
The dissociation constants of monochlorophenols (2-, 3-, 4-chlorophenols) were ex-
amined using direct UV–vis spectroscopy at temperatures from 25 to 175 C and at
saturated vapor pressures in a high-temperature, high-pressure cell. The dissociation
constant, Ka increased under experimental temperatures in the order: 2-chlorophenol,
3-chlorophenol, and 4-chlorophenol. The dissociation constant of 4-chlorophenol in-
creased with increasing temperature under experimental conditions, while those of 2-
and 3-chlorophenol reached maximum values at around 125 C, and then decreased
with further increases in temperature. The slope of ∂(log K)/∂ (1/T ) was noncon-
stant and positive, that is, endothermic, below 150 C. The data on dissociation con-
stants were analyzed by simultaneous regression to yield a five-term equation that
described the Van’t Hoff isobar. The magnitude of enthalpy 1H increased at 25 C in
the order: 3-chlorophenol, 4-chlorophenol, and 2-chlorophenol. The decrease in en-
thalpy at the absolute temperature was larger for 3-chlorophenol than for either 2- or
4-chlorophenol. Considering the equilibrium constant Kb for the isocoulombic reaction
of monochlorophenol with OH , the nearly linear relationship of log Kb vs. 1/T for
temperatures between 25 and 175 C indicates that the 1Cp values for this isocoulombic
reaction are low.
KEY WORDS: Monochlorophenol; dissociation constant; UV–vis spectroscopy;
high-temperature aqueous solution.
1. INTRODUCTION
Aromatic compounds, particularly chlorinated phenolic compounds, have
many applications. Chlorophenols are used as pesticides and herbicides and are the
by-products of several industries, being found, for example, in effluents from pulp
bleaching processes. Since chlorophenols have been classified as priority water
pollutants, properties affecting their removal from water merit special attention.
1Department of Applied Chemistry, Graduate School of Engineering, Tohoku University, Aoba 07,
Aramaki, Aoba-ku, Sendai 980-8579, Japan; email; miho@env.che. tohoku.ac.jp
19
C
0095-9782/03/0100-0019/0 2003 Plenum Publishing Corporation

20
Uchida and Okuwaki
The acid dissociation behavior of chlorophenols in aqueous solutions is impor-
tant in predicting the environmental fate of toxic effluents. Because hydrother-
mal conditions are increasingly used in industrial processes, such as wastewater
treatment,(1) it is important to characterize the reaction equilibria and kinetics of
ionic species in sub- and supercritical water. Furthermore, it is useful to determine
the effects of temperature and pH on these properties. Although a few studies have
determined the dissociation constants of monochlorophenols (2-chlorophenol, 3-
chlorophenol, and 4-chlorophenol) at room temperature,(2–14) the effect of tem-
perature on the constant parameters has not been examined except in the study by
Chen and Laidler.(14)
Various methods have been used to study hydrothermal reactions. Spectro-
scopic methods have been utilized widely to study hydrothermal reactions be-
cause they can use to measure reactant concentrations directly. Many readily avail-
able spectral techniques, including IR,(15,16) Raman,(16–19) and UV-visible,(16,20–42)
are being used to characterize reactions, such as equilibria and kinetics, under
hydrothermal conditions. Particularly, in spite of the limitation that measure-
ments of equilibria require species containing appropriate chromophores, UV–
vis spectroscopy have been used in many studies of ionic equilibria of inorganic
compounds,(24–26) complexes,(26–32) and organic compounds,(33–36,42) including
phenols(39) in high-temperature aqueous solutions. The hydrothermal equilibria
of compounds with insufficient absorbance in the UV–vis region have also been
examined using pH indicators that are stable at high temperatures.(37–42)
The purpose of this study was to analyze the acid dissociation of
monochlorophenols in high-temperature aqueous solutions and to characterize
the behavior of hazardous aromatic organochlorinated compounds under different
hydrothermal conditions. The dissociation constants of monochlorophenols at tem-
peratures up to 175 C and at saturated vapor pressures were examined by UV–vis
spectroscopy with a high-temperature, high-pressure cell. Since monochlorophe-
nols have absorption peaks in the UV–vis region, they are appropriate substances
for in situ measurements using UV–vis spectroscopy. The thermodynamic proper-
ties at infinite dilution were determined by regression analysis of the dissociation
constants of each monochlorophenol and variations in the temperature dependence
of acid dissociation equilibria between the different monochlorophenols were
examined.
2. EXPERIMENTAL
2.1. Chemicals and Solutions
2–Chlorophenol (99%) and 3-chlorophenol (99%) were purchased from GL
Sciences Inc.; 4-chlorophenol (99%) was obtained from Aldrich Chemical Co. An-
alytical grade sodium hydroxide was obtained from Kanto Chemical Co. Freshly

Dissociation of Monochlorophenols
21
distilled and deionized water from a Millipore QII system, with a specific resistance
of 18 M -cm, was used to prepare all of the solutions. For spectra measurements,
each monochlorophenol was prepared in aqueous solution to give a sample con-
centration of 5 10 5 M. The NaOH concentration of the solutions was adjusted
using a stock solution of 0.1 M sodium hydroxide.
2.2. Experimental Apparatus and Procedures
The high-pressure, high-temperature cell used for absorption measurements
of solutions is illustrated in Fig. 1. It consists of a cylindrical SUS316 cell with
quartz window that is sealed with PEEK gaskets and has a window aperture of
1 cm. The path length of 1 cm was calibrated by comparisons of the absorbances of
a series of K2Cr2O7/KOH solutions using a 10-mm silica glass cuvette. The overall
experimental apparatus is shown in Fig. 2. The optical cell was mounted in a heater
box and controlled to 1 C of the stated temperatures. Pressure was controlled
with an HPLC pump (Shimadzu LC-10AT VP) and a spring-loaded back-pressure
regulator (TESCOM series 26-1700). During spectrum measurements, the solu-
tions were kept at pressures that were slightly above saturated water vapor pressure
for each temperature. A three-way valve was installed in the feed tubing to allow
alternate pumping of sample solution and water, each with its own nitrogen sparge.
Before taking an absorption spectrum, each solution was deaerated for 30 min by
purging with nitrogen gas. The sample solution was then pumped into the cell at a
delivery rate of 0.05 to 0.2 ml-min 1 for 30 min. All of the spectra were collected
at 0.1-nm intervals using a double-beam spectrophotometer (JASCO V-570). The
spectra were corrected for background absorbance (windows and solution/solvent)
by subtracting the absorbance of a cell filled with pure water at each temperature.
The entire process of flushing the apparatus with a new solution or pure water
and obtaining the spectra took about 40 min for each temperature setting. During
spectral scans, the flow was stopped. The spectra were collected over the range
1
190 to 350 nm at a scan rate of 100 nm-min
.
Fig. 1. Side and cross-sectional views of the SUS316 cell with
quartz windows: (a) 316SS Body; (b) quartz window; (c) solu-
tion inlet; (d) solution outlet; (e) PEEK gasket.

22
Uchida and Okuwaki
Fig. 2. Experimental set-up.
3. RESULTS AND DISCUSSION
3.1. Spectra of Monochlorophenols
The UV–vis spectra of monochlorophenols (2-, 3-, and 4-chlorophenol) did
not change over a 3-h time period under the following conditions:temperature
range 25–175 C; NaOH concentration range 0–2.5 10 2 M; and saturated va-
por pressure. This shows that monochlorophenols do not decompose under these
conditions.

Dissociation of Monochlorophenols
23
The acid dissociation equilibrium for monochlorophenol (Ph-OH) solution is
Ph OH = Ph O + H+
(1)
with pKa at 25 C of 8.3 to 9.4 for 2-chlorophenol, 8.9–9.2 for 3-chlorophenol, and
9.2–9.4 for 4-chlorophenol.(2–14)
Figure 3 shows the spectra of monochlorophenols (5 10 5 M) under acidic
and alkaline conditions. In the range of wavelengths between 200 and 350 nm, each
Fig. 3. Typical spectra of 5 10 5 M monochlorophe-
nols at room temperature and pressure in (a) acidic, and
2
(b) alkaline media: (a) 0 M NaOH; (b) 2.5 10
M
NaOH.

24
Uchida and Okuwaki
monochlorophenol has two peaks and/or shoulders in both acidic and alkaline
aqueous solutions. 2-Chlorophenol and 3-chlorophenol show similar spectra in
acidic media (Fig. 3a), i.e., one shoulder at approximately 215 nm and one peak at
275 nm. In contrast, 4-chlorophenol shows two peaks at 225 and 280 nm that are
slightly red-shifted compared to those of the 2- and 3-chlorophenols. In alkaline
media, each monochlorophenol has two peaks, which are red-shifted compared to
thoseinacidicmedia, andwhicharelocalizedaround240and295nm(Fig. 3b). The
two absorption bands were shifted to progressively longer wavelengths according
to order: 2-, 3-, and 4-chlorophenol.
Figure 4 shows the effect of temperature on the spectrum of a 5 10 5 M
4
3-chlorophenol solution containing 2 10 M NaOH. Increasing temperature
leads to gradual decreases in the two absorption bands at 240 and 295 nm. This
indicates a shift toward the nonionic protonated form of 3-chlorophenol with in-
creasing temperature at constant NaOH concentration. The spectra of the 2- and
4-chlorophenols show similar changes with temperature. At temperatures above
175 C, additional reactions, such as the decomposition of chlorophenol and in-
teractions between NaOH and chlorophenol, occur and the spectra take on dif-
ferent appearances. Therefore, examination of the acid dissociation equilibria
of monochlorophenols was not possible at temperatures above 175 C. The in-
dividual spectra of 5 10 5 M monochlorophenols in aqueous solution without
added NaOH remain constant between 25 and 175 C. This is due to the solutions
being sufficiently acidic to keep the monochlorophenols fully protonated under
experimental conditions. Accordingly, solutions of monochlorophenols without
added NaOH were used to calibrate the absorbance of the pure acidic species of
Fig. 4. Effect of temperature on the absorbance spectra
of 5 10 5 M 3-chlorophenol in 2 10 4 M NaOH.

Dissociation of Monochlorophenols
25
Fig. 5. Effect of NaOH concentration on the absorbance
spectra of 5 10 5 M 3-chlorophenol at 100 C.
monochlorophenols, as described in the next section. The spectra from 240 to 350
nm of monochlorophenols in 2.5 10 2 M NaOH did not change in appearance
with temperature, but a slight red shift was observed.
Figure 5 shows the effect of NaOH concentration on the spectrum of a 5
10 5 M 3-chlorophenol solution at 100 C. There are at least two isosbestic points.
Increasing the NaOH concentration leads to gradual increases in the two absorption
bands at 240 and 295 nm. This indicates a shift toward the ionic deprotonated form
of 3-chlorophenol, the phenolate anion (Ph O ), with increasing NaOH concen-
tration. The spectra of 2- and 4-chlorophenol show similar changes with increasing
NaOH concentration. The absorption maxima of PhOH and PhO , which appear
in the wavelength range from 240 to 330 nm, allow the concentrations of PhOH
and PhO to be monitored by absorption spectroscopy using deconvolution of the
spectra. Therefore, spectral measurement with UV–vis spectroscopy can be used to
investigate the equilibrium behavior of monochlorophenols in aqueous solutions.
3.2. Dissociation Constants Ka of Monochlorophenols
The temperature dependence of the monochlorophenol acid–base equilibrium
is quantified by the dissociation constant at a constant pressure as a function
of temperature. Steady-state absorption spectroscopy is used to determine the
concentrations of monochlorophenols and their phenolate anions as described by
Ryan et al.,(34) and Xing and Johnston.(35)
Monochlorophenols (Ph OH) dissociate a proton and form phenolate anions
(Ph O ), as described Eq. (1).

26
Uchida and Okuwaki
The dissociation constant is expressed as
aPh aH
+
O
Ka =
(2)
aPh
OH
where ai denotes the activity of each species. By defining the activity coefficient
of each species, rearranging Eq. (2) gives
mPh
OH
+
log
=
log Ka + log( mH
)
(3)
mPh
O
where
is the mean activity coefficient of ions and mi is the molal concentration
mPh and (II)
of each species. Therefore, Ka is determined from (I) mPh
/
OH
O
+
+
aH (= mH ) at a given temperature.
First, the ratiomPh OH/mPh isobtained bydeconvolutionofthe absorption
O
spectrum. At a given wavelength i the total absorbance of a solution containing
both the acidic (Ph OH) and basic (Ph O ) species of monochlorophenols, is
expressed as the sum of the absorbances of both species:
xAεA(i) + xBεB(i) = ε(i)
(4)
where xA and xB are the molar concentration fractions of the acidic and basic
species of the monochlorophenols described in Eqs. (5) and (6), and ε (i) is the
absorbance of the mixture solution at a given condition. The εA(i) and εB(i) values
represent the absorbances of the solutions when all the monochlorophenols are
present in the acidic and basic species, respectively.
[Ph OH]
[Ph OH]0
[Ph OH]
xA
=
=
(5)
(6)
[Ph OH] + [Ph O ]
[Ph O ]
[Ph O ]
xB =
=
[Ph OH]0
[Ph OH] + [Ph O ]
[Ph OH]0 denotes the initial molar concentration of monochlorophenol, and
[Ph OH] and [Ph O ] are the molar concentrations of the acidic and basic
species, respectively.
The molar concentrations are converted to molal concentrations. For dilute
electrolyte solutions, the molality of species S, ms, may be obtained from the
molarity as follows:
ms = [S]/
(7)
where is the density of pure water.(43) Thus, Eq. (4) becomes
mPh
mPh
O
OH εA(i) +
εB(i) = ε(i)
(8)
m0
m0
where m0 is the initial molal concentration of the monochlorophenol.
As mentioned previously, a solution of monochlorophenol without any added
NaOH was used to calibrate the absorbance of the pure acidic species and a solution

Dissociation of Monochlorophenols
27
containing 2.5 10 2 M NaOH was used to calibrate the absorbance of the pure
basic species. Over the wavelength range from 240 to 330 nm, each spectrum was
deconvoluted into the acidic and basic contributions by finding the optimum values
of xA and xB using the nonlinear least squares method to estimate the dissociation
constant Ka. The agreement factor U, which is the sum of the least square of
the difference between the measured [ε(i)obs] and calculated absorbance [ε(i)calc
]
from 240 to 330 nm, was calculated using
X
2
U =
[ε(i)obs ε(i)calc
]
i = 240 - 330 nm
(9)
The optimum values of xA and xB, (i.e., those which give the minimum value
of U) were calculated. U did not exceed 0.08 in any curve fitting of spectra. Ex-
amples of typical deconvoluted spectra of 4-chlorophenol in a 2 10 4 MNaOH
solution at 125 C are shown in Fig. 6. The calculated spectrum was estimated as the
sum of the spectra of acidic and basic fractions. The difference in total absorbance
between the observed spectrum and the calculated spectrum was generally less
than 4% at any given wavelength. The ratio mPh OH/mPh
was estimated from
O
the optimum values of xA and xB. The observed ratios of mPh OH/mPh
in all
O
experiments are listed in Tables I–III. The sum of the differences between the total
absorbance in the observed spectrum and the calculated spectrum was plotted as
a function of temperature (Fig. 7a), and initial sodium hydroxide concentration
(Fig. 7b). The sums of the differences exhibit no systematic relationship to low
Fig. 6. Deconvoluted spectra of 4-chlorophenol in 2
10 4 M NaOH solutionat 125 C: (a) observed spectrum;
(b) calculated spectrum; (c) spectrum of acidic fraction;
(d) spectrum of basic fraction; (e) residual spectrum be-
tween (a) and (b).

28
Uchida and Okuwaki
Table I. Initial Concentration of 2-Chlorophenol and Sodium Hydroxide and Observed and
Calculated Values for the mPh OH/mPh
Ratio and Molal Concentrations of Protonsa
O
C
105m0Ph
104m0NaOH
mPh OH/mPh
mPh OH/mPh
O calc
log mH
+
O
obs
OH
25
25
25
25
50
50
50
75
75
5.02
5.02
5.02
5.02
5.06
5.06
5.06
5.13
5.13
5.13
5.13
5.13
5.22
5.22
5.22
5.22
5.22
5.22
5.22
5.32
5.32
5.32
5.32
5.32
5.32
5.45
5.45
5.45
5.45
5.61
5.61
5.61
5.61
5.61
0.702
1.00
1.50
19.8
1.01
1.52
2.02
0.718
1.03
1.54
2.05
5.13
1.04
1.57
2.09
5.22
1.03
2.06
10.3
0.532
0.745
1.07
1.60
2.13
5.32
2.18
5.45
10.9
21.8
0.785
1.12
1.68
2.24
5.61
0.493
0.163
0.111
0.163
0.333
0.220
0.111
1.44
0.695
0.389
0.266
0.124
1.27
0.667
0.471
0.191
1.27
0.493
0.099
5.67
3.35
1.94
1.00
0.639
0.163
1.35
0.539
0.282
0.177
9.00
0.341
0.163
0.111
0.162
0.334
0.220
0.112
1.44
0.699
0.391
0.264
0.123
1.27
0.669
0.472
0.189
1.28
0.494
0.100
5.66
3.35
1.95
1.01
0.640
0.163
1.35
0.536
0.282
0.177
9.02
9.51
9.75
10.0
9.74
9.08
9.32
9.47
8.42
8.57
8.78
8.93
9.38
8.18
8.36
8.50
8.94
8.17
8.50
9.26
7.57
7.71
7.86
8.04
8.17
8.60
7.93
8.35
8.66
8.97
7.29
7.45
7.62
7.75
8.16
75
75
75
100
100
100
100
100
100
100
125
125
125
125
125
125
150
150
150
150
175
175
175
175
175
6.14
3.35
2.57
0.961
6.14
3.35
2.57
0.958
aAll concentrations are molalities.
initial NaOH concentration. On the other hand, at high initial NaOH concentration,
the sums of the differences were distributed toward positive values. Similarly, the
residual values changed from negative to positive with increasing temperature. The
absorbance of the observed spectrum was slightly higher than that of the calculated
spectrum at the measured wavelength. Due to some ion-pairing reaction at high
temperature and high NaOH concentration, the absorbance of the observed spec-
trum was probably too high. Equilibrium constants were estimated from data that

Dissociation of Monochlorophenols
29
Table II. Initial Concentration of 3-Chlorophenol and Sodium Hydroxide and Observed and
Calculated Values for the mPh OH/mPH
Ratio and Molal Concentrations of Protonsa
O
C
105m0Ph
104m0NaOH
mPh OH/mPh
mPh OH/mPh
O calc
log mH
+
O
obs
OH
25
25
25
25
50
50
50
50
75
75
75
75
75
100
100
100
100
100
100
125
125
125
125
125
125
150
150
150
150
150
150
150
175
175
175
175
175
175
5.02
5.02
5.02
5.02
5.06
5.06
5.06
5.06
5.13
5.13
5.13
5.13
5.13
5.22
5.22
5.22
5.22
5.22
5.22
5.32
5.32
5.32
5.32
5.32
5.32
5.45
5.45
5.45
5.45
5.45
5.45
5.45
5.61
5.61
5.61
5.61
5.61
5.61
0.702
1.00
1.50
2.01
0.506
1.01
1.52
2.02
0.513
1.03
1.54
2.05
5.13
0.731
1.04
1.57
2.09
5.22
10.4
0.532
0.745
1.07
2.13
5.32
10.7
0.763
1.09
1.64
2.18
5.45
10.9
21.8
0.785
1.12
2.24
5.61
11.2
22.4
0.852
0.409
0.177
0.136
2.70
0.754
0.429
0.266
4.88
0.854
0.410
0.176
0.135
2.70
0.760
0.429
0.265
4.91
9.63
9.81
10.0
10.2
8.84
9.14
9.34
9.49
8.34
8.62
8.81
8.94
9.38
8.05
8.22
8.40
8.53
8.95
9.26
7.60
7.74
7.89
8.19
8.61
8.92
7.49
7.65
7.82
7.95
8.35
8.66
8.97
7.30
7.46
7.76
8.16
8.47
8.77
1.44
1.44
0.821
0.471
0.136
3.35
2.57
1.56
0.821
0.472
0.136
3.33
2.57
1.56
1.00
1.00
0.370
0.191
11.5
6.14
3.76
0.372
0.192
11.5
6.15
3.78
1.44
1.43
0.409
0.136
10.1
6.69
4.56
3.00
1.08
0.539
0.266
11.5
9.00
4.26
1.78
0.887
0.471
0.407
0.136
10.1
6.70
4.55
3.01
1.08
0.542
0.268
11.5
9.00
4.24
1.78
0.886
0.471
aSee footnote to Table I.
gave a value of log (mPh OH/mPh
) between approximately 1 and 1.2, because
O
values outside this range are sensitive to small errors in the absorbance data.
For singly charged species, the mean ionic activity coefficient of a 1-1 elec-
trolyte was calculated using Pitzer’s formulation of the Debye–Hu¨ckel

30
Uchida and Okuwaki
Table III. Initial Concentration of 4-Chlorophenol and Sodium Hydroxide and Observed and
Calculated Values for the mPh OH/mPH
Ratio and Molal Concentrations of Protonsa
O
C
105m0Ph
104m0NaOH
mPh OH/mPh
mPh OH/mPh
O calc
log mH
+
O
obs
OH
25
25
25
50
50
50
50
50
75
75
75
75
75
5.02
5.02
5.02
5.06
5.06
5.06
5.06
5.06
5.13
5.13
5.13
5.13
5.13
5.22
5.22
5.22
5.22
5.22
5.22
5.22
5.22
5.22
5.32
5.32
5.32
5.32
5.32
5.32
5.32
5.32
5.32
5.32
5.45
5.45
5.45
5.45
5.45
5.45
5.45
5.45
5.45
5.61
5.61
5.61
5.61
5.61
1.00
1.50
2.01
0.709
1.01
1.52
2.02
5.06
0.718
1.03
1.54
2.05
5.13
0.731
1.04
1.57
2.09
5.22
10.4
2.06
10.3
20.6
0.745
1.07
1.60
2.13
5.32
10.7
2.10
10.5
21.0
52.6
1.09
1.64
2.18
5.45
10.9
21.8
2.15
10.8
53.8
1.12
2.24
5.61
11.2
22.4
0.754
0.429
0.299
2.57
0.752
0.428
0.299
2.59
9.85
10.1
10.2
9.03
1.56
1.57
9.19
9.36
9.51
9.94
8.51
8.66
8.84
8.97
9.39
8.09
8.25
8.42
8.55
8.96
9.27
8.54
9.26
9.57
7.76
7.92
8.09
8.22
8.62
8.92
8.21
8.92
9.23
9.63
7.66
7.83
7.96
8.36
8.66
8.97
7.96
8.66
9.36
7.47
7.77
8.17
8.47
8.77
0.754
0.639
0.220
4.88
3.00
1.63
1.22
0.429
8.09
5.67
3.00
0.754
0.639
0.220
4.91
2.99
1.64
1.23
0.430
8.12
5.66
3.00
100
100
100
100
100
100
100
100
100
125
125
125
125
125
125
125
125
125
125
150
150
150
150
150
150
150
150
150
175
175
175
175
175
2.23
2.23
0.754
0.351
2.03
0.370
0.190
13.3
9.00
4.88
4.00
1.27
0.613
3.76
0.695
0.333
0.163
13.3
7.33
7.33
2.03
1.00
0.493
6.69
1.22
0.282
15.7
7.33
2.85
1.50
0.786
0.757
0.350
2.02
0.372
0.189
13.2
8.96
4.90
4.00
1.27
0.612
3.79
0.694
0.333
0.163
13.2
7.31
7.38
2.04
1.00
0.494
6.69
1.22
0.281
15.7
7.31
2.83
1.50
0.785
aSee footnote to Table I.

Dissociation of Monochlorophenols
31
Fig. 7. Deviation plots for the sum of the differ-
ence of total absorbance between calculated and ob-
served spectra as a function of (a) temperature and
(b) initial NaOH concentration. Symbols designate
measurements for solutions of 2-chlorophenol (cir-
cles), 3-chlorophenol (triangles), and 4-chloropenol
(diamonds).
equation:(44)
I1/2
1 + bI1/2
2
ln
=
A
+
ln(1 + bI1/2
)
(10)
b

32
Uchida and Okuwaki
where b = 1.2. The ionic strength I is given by
X
1
1
mizi2 = (mH + mNa + mOH + mPh
)
(11)
+
+
I =
O
2
2
i
where mi and zi are the molalities and charges of all the ionic species in solution,
respectively. The Debye–Hu¨ckel formulation of the osmotic coefficient A is given
by
3/2
1/2
1
3
2 NA
e2
A =
(12)
1000
DkT
where NA is Avogadro’s number, e is the charge of the electron, D is the dielectric
constant of water calculated by Uematsu and Franck,(45) and is the density of
water.(43)
The chlorophenol and Na mass balances and the charge balance are given by
m0Ph
= mPh OH + mPh
(13)
(14)
(15)
O
OH
0
+
mNaOH = mNaOH + mNa
+
+
mH + mNa = mPh
+ mOH
O
where m0Ph OH and mN0 aOH are the initial molal concentration of monochlorophenol
and sodium hydroxide, respectively.
Using the mean activity coefficient, the acid dissociation constant in Eq. (2)
is rewritten as
2
+
mH mPh
O
Ka =
(16)
mPh
OH
(46)
The equilibrium constant for NaOH association Km
and for the ion product of
water Kw(47) are given by Eqs. (17) and (18), respectively.
mNaOH
Km
=
(17)
(18)
2
+
mNa mOH
2
+
Kw = mH mOH
+
The mH is calculated by Eqs. (13–18) using Newton’s iteration method. In
+
order to estimate mH , it is necessary to know I to determine . To determine I,
the molalities of the various ions were determined from the relevant equilibrium
constants, Km and Kw. The molalities of the various ions were then calculated in
the following manner. First, preliminary values of Ka were estimated. Then, an
initial ionic strength was estimated and substituted into Eq. (10) to determine the
mean activity coefficient. The concentration of each ionic species was determined
using the Newton–Raphson method by the simultaneous solution of Eqs. (13–
18) for the hydrogen ions. The ionic strengths of the solutions were calculated

Dissociation of Monochlorophenols
33
from the concentration of each ion and determined by the method of successive
approximations. The molalities of Ph OH (mPh OH) and Ph O (mPh ) were
O
determined to calculate the value of mPh OH/mPh
. Then, the experimental data
O
were analyzed with a nonlinear least squares method to determine the best value
of Ka (i.e., that which minimized the sum of the squares of the difference between
the observed and calculated values of the ratio mPh OH/mPh
). The resulting
O
values of the mean activity coefficient were close to unity.
+
Finally, the calculated value of mH and the mean activity coefficient and
observed value of the ratio mPh OH/mPh
were substituted into Eq. (3) to de-
O
+
termine the dissociation constants, Ka. The calculated values of mH and the ratio
mPh OH/mPh
are shown in Tables I–III.
O
Figure 8 displays the data obtained at four temperatures and plotted according
to Eq. (3). Similar plots constructed for each temperature studied were all linear.
However, contrary to what was expected from Eq. (3), the slope values varied
between 0.86 and 1.55. The values of log Ka were determined from the intercepts
of these plots.
The measured values of the equilibrium constants (Kal, 2-chlorophenol: Ka2,
3-chlorophenol; Ka3, 4-chlorophenol) are shown in Table IV. The measured val-
ues of log Kai for the dissociation constants of monochlorophenols are plotted in
Fig. 9 from 25 to 175 C, at 25 C intervals. The dissociation constants increase
in the order 2-chlorophenol, 3-chlorophenol, and 4-chlorophenol, with increasing
temperature up to 125 C. The dissociation constant of 4-chlorophenol increases
with increasing temperature under experimental conditions, while those of 2- and
Fig. 8. Dependence of the ratio of neutral 4-
chlorophenol and its phenolate anion on molalities
of hydrogen ion at 100 C (circles), 125 C (squares),
150 C (triangles), and 175 C (diamonds).

34
Uchida and Okuwaki
Table IV. Dissociation Constants of 2-Chlorophenol (Kal),
a
3-Chlorophenol (Ka2), and 4-Chlorophenol (Ka3
)
C
log Kal
log Ka2
log Ka3
25
50
75
100
125
150
175
8.96 0.17
8.69 0.21
8.46 0.27
8.23 0.13
8.06 0.55
8.05 0.21
8.12 0.16
9.56 0.38
9.10 0.48
8.76 0.59
8.57 0.12
8.31 0.66
8.41 0.14
8.42 0.09
9.74 0.08
9.34 0.24
9.06 0.22
8.86 0.22
8.72 0.18
8.68 0.20
8.65 0.06
aThe uncertainties indicate three times the standard deviation.
3-chlorophenol have maximum values around 125 C, which then decrease with
further increases in temperature. The slope of ∂(log K)/∂(1/T ) is not constant
over the entire temperature range. Each Ka exhibits a maximum that is simi-
lar to the dissociation constants of -naphthoic acid(28) and 2,5-dinitrophenol(34)
below 200 C. Maxima occur at 100 and 110 C for -naphthoic acid and 2,5-
dinitrophenol, respectively. At low temperature, the energy for deprotonation is
greater than that for ion solvation. With increasing temperature and decreasing den-
sity, the dielectric constant of water decreases. Since the reduction of the dielectric
constant makes the solvation of ionic products much less favorable, the effect
Fig. 9. Correlation of the logarithm of the dissoci-
ation constants of 2-chlorophenol (Ka1, circles), 3-
chlorophenol (Ka2, triangles), and 4-chlorophenol (Ka3
,
diamonds) with the inverse of temperature. Solid lines:
regression equation (Eqs. 20–22) calculated using the
Van’t Hoff equation.

Dissociation of Monochlorophenols
35
of ion solvation becomes much greater. Consequently, the maxima occur under
subcritical temperatures. Since monochlorophenols are weaker acids than either
-naphthoic acid or 2,5-dinitrophenol, the energy for deprotonation is greater. The
energy for ion solvation becomes less important at a lower temperature than in the
case of either -naphthoic acid or 2,5-dinitrophenol. As a result, the maxima oc-
cur at higher temperatures for monochlorophenols than for -naphthoic acid and
2,5-dinitrophenol.
Assuming that the heat capacity change 1Cp,ai for the acid dissociation of
monochlorophenols is represented in Eq. (19), the values of Kai(i = 1–3) as a
function of temperature were fitted to a form of the integrated Van’t Hoff isobar
[∂(ln K)/∂(1/T )]p = 1Ho/R
2
1Cp,ai = a + bT + cT
(19)
ln Kal
ln Ka2
ln Ka3
=
=
=
264 + 44.8 ln T + 4.84 103/T 8.35 10 2T 2.87 105/T 2
(20)
333 + 58.5 ln T + 5.19 103/T 11.9 10 2T 4.00 105/T 2
(21)
240 + 40.8 ln T + 3.65 103/T 7.95 10 2T 3.00 105/T 2
(22)
The resulting regression curves are shown in Fig. 9. The regression curves for
2- and 4-chlorophenol increase monotonically and the increasing rate for disso-
ciation constants decreases with increased temperature. The regression curve for
3-chlorophenol have a maximum around 150 C, which is similar to the observed
values for the dissociation constants.
The thermodynamic quantities 1Hai, 1Sai, 1Cp,ai (i = 1–3) in this equilib-
rium were estimated by differentiation of Eqs. (20–22), and the results are shown
in Table V. For the reaction between a weak acid and a weak base (e.g., the disso-
ciation reaction of monochlorophenols in water), the enthalpy is positive. It is an
endothermic reaction, since the energy needed to break the O-H bond is greater
than that released by solvation of the resulting ions.(28)
The magnitude of the enthalpy 1Hai was in the order of 3-chlorophenol >
4-chlorophenol > 2-chlorophenol, at 25 C. All values of 1Hai decreased with
increasedtemperature. Becausetheenergyofionsolvationbecomesmucheffective
with increasing temperature, the enthalpy of reaction is controlled by ion solvation
rather than the bond energies. The decrease in enthalpy with absolute temperature
is larger for 3-chlorophenol than for 2- or 4-chlorophenol.
The Cl substituent on the aromatic ring acts as an electron-withdrawing group,
thereby producing an inductive effect that increases the acidity of the phenol.
In contrast, halogen substituent groups are ortho, para directors because they

36
Uchida and Okuwaki
Table V. Summary of Thermodynamic Quantities for the Dissociation of
Monochlorophenols in Aqueous Solution at the Saturated Vaper Pressure
C
log Kaa
1Hab
1Sac
1Cpc,a
2-Chlorophenol
25
50
75
100
125
150
175
8.98
8.66
8.41
8.24
8.13
8.08
8.07
25.2
22.5
19.1
15.0
10.1
4.5
2.0
87
96
12
30
71
111
150
188
226
106
118
130
144
159
3-Chlorophenol
25
50
75
100
125
150
175
9.55
9.09
8.76
8.54
8.42
8.37
8.40
36.1
31.1
25.2
18.2
10.1
0.8
9.4
62
78
95
115
136
158
182
30
91
149
206
261
316
369
4-Chlorophenol
25
50
75
100
125
150
175
9.72
9.35
9.07
8.87
8.74
8.66
8.64
28.8
25.7
22.0
17.5
12.4
6.6
90
100
111
123
136
150
165
1
40
80
119
156
193
229
0.0
aCalculated by fitting to a form of the integrated Van’t Hoff isobar.
1
bkJ-mol
.
cJ-mol 1-K
1.
are electron-donating via the resonance effect. Since resonance effects oppose
inductive effects in the ortho, para positions, the electron-withdrawing effect in
those position is smaller than in the meta position. As a result, the energy needed
to break the O-H bond of 3-chlorophenol is greater than that required to break the
O-Hbondofeither2-or4-chlorophenol. Themaximumofthedissociationconstant
occurs at lower temperature for 3-chlorophenol than for 2- or 4-chlorophenols.
Likewise, the enthalpy of the reaction (endothermic) at lower temperatures is
greater for 3-chlorophenol than for 2- or 4-chlorophenol.
3.3. Isocoulombic Reactions of Monochlorophenols with OH
The Kb for the isocoulombic reaction of monochlorophenol with OH was
not measured directly, but was calculated from the measured Ka data, as shown in
Fig. 10. For the reaction

Dissociation of Monochlorophenols
37
Fig. 10. Effect of temperature on the equilib-
rium constants, Kb, for the isocoulombic reaction
of monochlorophenols with OH : 2-chlorophenol
(Kb1; circles), 3-chlorophenol (Kb2; triangles), and 4-
chlorophenol (Kb3, diamonds). Solid lines: linear re-
gression of the data.
Ph OH + OH = Ph O + H2O
the equilibrium constant is expressed as:
(23)
(24)
Ka
aPh
O
Kb =
=
aPh OH aOH
Kw
The most significant difference is that Ka exhibits a maximum, whereas Kb
does not. The nearly linear relationship of log Kb vs. 1/T for temperatures between
25 and 175 C indicates that the 1Cp,bi values for this isocoulombic reaction are
low. Therefore, assuming that the heat capacity change 1Cp,bi is zero (i.e., the
enthalpy 1Hbi has constant value and is not dependent on temperature), the en-
thalpies were calculated as the slope of log Kbi vs. 1/T . The resulting values of
1
1Hbi were 27.8, 23.8, and 25.1 kJ-mol for 2-, 3-, and 4-chlorophenol,
respectively. For the reaction of -naphthol with the strong base OH , the en-
thalpy is always negative.(35,36) Similarly, the reaction of monochlorophenols and
OH is isocoulombic, and so the enthalpy of reaction is controlled primarily by
the bond energies and not by ion solvation. Thus, the reaction described in Eq.
(23) is exothermic. The enthalpy of Eq. (1) 1Hai was recalculated by adding the
enthalpy of dissociation of water, 1Hw to 1Hbi. The obtained values of 1Hai
1
were 27.7, 31.8, and 30.5 kJ-mol for 2-, 3-, and 4-chlorophenol, respectively.
The order of the magnitude of 1Hai was in agreement with that at 25 C. However,
as mentioned above, since 1Hai decreases with temperature and the effect is larger

38
Uchida and Okuwaki
for 3-chlorophenol than for 2- or 4-chlorophenol, the order of the magnitude of
1Hai is reversed with increases in temperature.
4. CONCLUSIONS
Quantitative measurements of the dissociation constants, Ka for monochloro-
phenols (2-, 3-, and 4-chlorophenols) were performed with direct UV–vis spec-
troscopy between 25 and 175 C. Deprotonation of the monochlorophenols be-
comes favorable as the solution temperature is raised from 25 to 125 C. Ka exhibits
endothermic behavior below 150 C for all monochlorophenols, since the energy
needed to break the O-H bond for this relatively weak acid is greater than that re-
leased by solvation of the resulting ions. The reaction changes from endothermic to
exothermic above 150 C for 3-chlorophenol, which is similar to what is observed
with 2,5-dinitrophenol and -naphthoic acid. The decrease in enthalpy with abso-
lute temperature is larger for 3-chlorophenol than for 2- or 4-chlorophenol. The
differences in behavior between 3-chlorophenol and 2- or 4-chlorophenol may
be described by the halogen substituent group. Since halogen substituent groups
are electron-donating by the resonance effect, they act as ortho, para directors
and the electron-withdrawing effect in the o-, p-position is smaller than in the m-
position. As a result, the energy needed to break the O-H bond of 3-chlorophenol
is greater than that needed to break the O-H bond of 2- or 4-chlorophenol.
ACKNOWLEDGMENTS
I thank Dr. Satoshi Uchida for the use of the UV–vis spectrometer system.
ThisresearchwaspartiallysupportedbytheMinistryofEducation, Culture, Sports,
Science and Technology, and by a Grant-in-Aid for the Encouragement of Young
Scientists (H11-11780397).
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