P R E P R I N T – ICPWS XV
Berlin, September 8–11, 2008
UV-visible Spectroscopic Study on Nitrophenols Ionization Reactions to 225 oC
Jana Ehlerovaa, Liliana Trevanib, Josef Sedlbauera, Karine Ballerat-Busserollesc and Peter R. Tremaineb
a
Department of Chemistry, Technical University of Liberec, 46117 Liberec, Czech Republic
b
Department of Chemistry, University of Guelph, Guelph, Ontario, Canada N1G 2W1
c
Lab. de Thermodyn. des Solutions et des Polymères, Université Blaise Pascal, 63177 Aubière, France
Email: [email protected]
The UV-visible spectra of aqueous o-, m-, and p-nitrophenol were measured as a function of pH at
temperatures from 50°C to 225°C at a pressure of 7 MPa. These were used to determine equilibrium constants
for the acid ionization reaction of each isomer. The new experimental results, along with data from the
literature, were used to develop a thermodynamic model to describe the dependence of ionization properties on
temperature and pressure. The model yields predictions of the ionization constants for o-, m-, and pnitrophenol, log Ka, to at least 250°C and 20 MPa with an estimated uncertainty in log Ka of less than ± 0.06.
Thermal decomposition studies, carried out by using the system as a stopped-flow reactor, showed that the
three isomers are significantly more stable in HCl acid solutions. At temperatures above 225 oC and alkaline
solutions, the residence time of the solution in the cell and the pre-heater is so long that significant changes in
the spectra of the nitrophenolate species are observed in stopped-flow experiments as an indication of thermal
decomposition reactions.
Introduction
Experimental Methods
Nitrophenols are used as raw materials in the
chemical, pharmaceutical, dyes and herbicides
industries, and are also produced in gas-phase
oxidation of simple aromatic hydrocarbons [1].
Their toxicity is high even at trace levels,
particularly for green plants and for aquatic
organisms in surface waters. Thermodynamic
properties are needed for the design of effective
remediation technologies, for instance, steam
oxidation.
Nitrophenols are well-known colorimetric pH
indicators, due to the differences in the UV-visible
spectra of the protonated and unprotonated species
(190 - 500 nm). However, only the ionization
constants for p-nitrophenol have been determined at
temperatures above 75 °C by colorimetric methods
[2].
The main objective of this work was to
supplement the existing experimental information
on aqueous nitrophenols (see for instance
references in [3]) by determining their acid-base
ionization constants over a wide range of
temperature. A second objective was to develop a
thermodynamic model able to describe these
reactions in a form that is suitable for developing
predictive models, such as functional-group
contribution models, for describing the properties
of structurally similar solutes in water.
Ionization constants were measured by UVvisible spectroscopy in a high temperature-high
pressure platinum flow cell with sapphire windows.
A schematic diagram of the cell and injection
system is shown in Figure 1. Details of the
construction are described elsewhere [4].
Ammonia/ammonium buffer solutions were used to
prepare nitrophenol solutions at fixed and known
pH, with different ratios of protonated to
unprotonated species. To determine the spectra of
the fully protonated and fully unprotonated
indicator, nitrophenol was dissolved in solutions of
dilute HCl(aq), (the “acid extreme”) and NH3(aq) or
Figure 1: Schematic of high temperature-high
pressure cell and injection system.
NaOH(aq) (the “base extreme”). For o- and pnitrophenol at 225 oC and NaOH(aq) (the “base
extreme”) for m-nitrophenol and p-nitrophenol at
temperatures under 200 oC.
Thermal Stability Experiments
Stopped-flow experiments were conducted at
25± and 225±C in order to study the thermal
stability of nitrophenols in the aqueous
ammonia/ammonium buffer solutions and in the
solutions of HCl(aq), NaOH(aq) and NH3(aq) used
to define the two extremes. In these experiments the
flow was stopped for 10 minutes and the
absorbance spectra between 250 and 500 nm were
measured every 30 seconds.
Significant
time-dependant
changes
in
absorbance were observed only for the o- and pnitrophenol solutions in aqueous NaOH at 225 oC,
but not for m-nitrophenol. When NH3 was used
instead of NaOH, the changes were small and
compatible with the time scale of the experiments
(collection times). Typical results are shown in
Figures 2 and 3. No attempt was made to determine
the decomposition mechanism in alkaline media or
to identify the decomposition products.
Although the uncertainty associated with
experiments at 225 oC is somewhat larger than that
at lower temperatures, the experimental results at
this temperature were included in the model
presented below.
Figure 2: Stopped-flow experiment for
o-nitrophenol in alkaline media at 225 °C and 7
MPa : (a) in NaOH(aq) and (b) in aqueous NH3(aq).
Time between spectra, 30 seconds.
Spectroscopic analysis
UV-visible spectra for the solutions of o-, mand p-nitrophenol at constant pressure, 7 MPa, and
temperatures from 100 °C to 225 °C were collected
at wavelength intervals of 0.25 nm from 200 nm to
800 nm and a scan rate of 1200 nm.min-1.
Representative spectra for all three isomers at 175
o
C are shown in Figure 4. The temperature range
was determined by the useful pH range of the
ammonia/ ammonium buffer solutions which
defined the low temperature limit, and the thermal
stability of the indicator. At temperatures over
225°C, the results may include systematic errors
because of thermal decomposition of the indicator.
The deconvolution of the nitrophenol spectra in
the buffer solutions was carried out by linear
combination of the spectra of the acid and the base
forms of the nitrophenol species (Beer’s law). From
the deconvolution of the spectra obtained for each
buffer system, the concentration ratio of ionized
and un-ionized forms of the nitrophenol
Figure 3: Stopped-flow experiments for
m-nitrophenol in NaOH(aq) base extreme and pnitrophenol in NH3 (aq) base extreme at 225 °C and
7 MPa. Time between spectra, 30 seconds.
2
and calculated values, respectively, and σX is the
estimated experimental uncertainty.
(the “indicator ratio”) could be determined at
known pH. These ratios could then be used to
calculate the ionization constant. The resulting
values of the fitted ionization constants for all the
buffer solutions were averaged to determine the
experimental pKa value at each temperature. The
number of buffer solutions at each temperature
ranged from 2 to 6.
2
⎛ ΔH aexp − ΔH acalc ⎞
⎟⎟ +
+ ∑ ⎜⎜
σΔH a
j =1 ⎝
⎠j
P
Results
2
calc
exp
calc
R ⎛
⎛ ΔC exp
⎞
⎞
p,a − ΔC p,a
⎟ + ∑ ⎜ ΔVa − ΔVa ⎟
+ ∑ ⎜⎜
⎜
⎟
⎟ l =1
σΔC p,a
σΔVa
k =1
⎝
⎠l
⎝
⎠k
Q
Average pKa values for the three nitrophenol
isomers are tabulated elsewhere [3], along with
their estimated uncertainties. The uncertainty
estimates are composed of the standard deviation of
the calculated pKas at a given temperature plus the
contribution of possible systematic errors in
solution preparation. The two uncertainty
contributions are usually comparable and do not
exceed 0.03 pKa units each, leading to total
uncertainties of up to 0.06 pKa units.
Thermodynamic model
Our experimental values of pKa can be
combined with other thermodynamic data for
nitrophenol ionization taken from the literature.
These include standard partial molar enthalpies
ΔHΕ, volumes ΔVΕ, and heat capacities
ΔCpΕwhich are known near 25 ΕC. Together with
our results, there are enough experimental results to
develop a model describing the changes of Ka for
each nitrophenol with temperature and pressure.
Several equations of state for the standard
thermodynamic properties of hydrothermal
solutions could be applied for this purpose [e.g. 5, 6,
7]. The model proposed originally by Marshall and
Franck [6] as a correlation tool for the ionization
constant of water
was chosen, because of its simplicity. For Ka the
equation takes the form,
ln K a = − pK a . ln 10 =
=a+
2
⎛ pK aexp − pK acalc ⎞
⎟⎟ +
F = ∑ ⎜⎜
σpK a
i =1 ⎝
⎠i
O
b c
d ⎛
f
g ⎞
+ 2 + 3 + ⎜ e + + 2 ⎟. ln ρ w
T T
T ⎝ T T ⎠
where ρw is the density of water, T is
thermodynamic temperature, and a – g are fitting
parameters. Other ionization properties are obtained
by appropriate derivations of this equation. The
ionization data for each nitrophenol were subjected
to a simultaneous weighted fit with the objective
function where exp and calc stand for experimental
3
2
temperatures arises because with the exception of
the volumetric data at 25°C, the calculated values
are entirely
Figure 4 : Normalized spectra at t = 175 °C, p = 7
MPa for aqueous o-, m-, and p-nitrophenol. The
arrows indicate the direction of increasing pH.
The resulting fitted parameters of the MarshallFranck model for the three nitrophenols are
reported in [3].
Figure 5 provides a graphical comparison of
experimental Ka with predictions of Marshall and
Franck model along the saturation line for water.
The fit of the model to the experimental data
was well within the estimated uncertainties of all
the data used in the regression.
The only
experimental values for comparison with our results
are the ionization constants of p-nitrophenol
reported by Shin et al. [2] at temperatures up to
200°C.
These were obtained by UV-visible
spectroscopy using methods similar to ours in batch
cells, rather than flow cell. The results agree to
within the combined estimated uncertainties,
although Shin et al.’s results for pKa for pnitrophenol are systematically more positive with
increasing temperature below 175 °C. We regard
their measurement at 200 °C as suspect, because of
the possibility of thermal decomposition, which
should be more pronounced in a batch cell due to
the much longer exposure times. Possible reasons
for the systematic difference between the two
studies include the choice of activity coefficients,
and approximations used in data treatment.
Unfortunately, it is not clear from the Shin et al.
paper, how the solutions were prepared and what
was the target ionic strength, so we could not reconstruct their treatment precisely.
Figure 5: Ionization constants of nitrophenols
corrected to psat, plotted as pKa vs. T. Fitted
Marshall and Franck model (full line). pKa data
from this work corrected from p = 7 MPa to psat (>);
o-nitrophenol from Judson and Kilpatrick [8] ( );
m-nitrophenol from Robinson and Peiperl [9] (∼);
p-nitrophenol from Judson and Kilpatrick [8] ( );
p-nitrophenol from Shin et al. [2] (∼)
Solvation Effects
The effects of temperature and pressure on
solvation are reflected in the standard state
thermodynamic properties derived from the model
[6], see Fig.6. In our judgment, the database from
which the model parameters were fitted is too
limited to permit realistic estimates of the
uncertainties associated with these calculated
derivative functions by error propagation analysis.
Instead, approximate uncertainty limits were
estimated from the relative error of model
predictions from experimental pKa data, and the
magnification of uncertainty upon derivation, to
determine the uncertainties in other properties [3].
This leads to estimated uncertainties in ΔHaº of 0.4
- 0.8 kJ·mol-1; ΔSaº, 3 - 6 J·K-1·mol-1; ΔCp,a º, 10 - 40
J·K-1·mol-1; and ΔVaº, 0.4 - 4 cm3·mol-1 where the
lower limits are for 25°C, the upper limits for
200°C. The large uncertainty in ΔVaº at elevated
temperature dependence, rather than the pressure
dependent on parameters determined by fitting the
dependence of pKa.
There are significant differences in the
properties of the three isomers at room temperature.
These have been attributed to the electronwithdrawing effects of the nitro group, which
stabilize the nitrophenolate anion, and to solvation
effects that manifest themselves in the entropy, heat
capacity and volume of ionization. The electronwithdrawing effects are larger for the o- and pisomers than for m-nitrophenol (See, for example,
4
polarization in the near-critical region [15,16].
Because ΔCp,aº and ΔVaº are derivative functions,
Dewick [11]). Fernandez and Hepler [12] have
noted that m-nitrophenol is a weaker acid than o- or
p-nitrophenol at room temperature because its
entropy of ionization is significantly more negative,
and not because of a more endothermic enthalpy of
ionization. Apparent molar volume and heat
capacity measurements on p- and m-nitrophenol
and their sodium salts [13, 14] suggest that this
effect arises primarily because of differences in
solvation for the uncharged species, with Vº(pNphOH) < Vº(m-NphOH), and Cpº(p-NphOH) <
Cpº(m-NphOH).
These observations are all
consistent with solvation effects in which water is
more tightly bound by the o- and p- isomers than by
the m- isomer at 25°C.
Although the relative positions of the Gibbs
energies of ionization of the isomers do not change
with increasing temperature, there are changes in
the contributions of the standard state enthalpy and
entropy. While both ΔHaº and ΔSaº become
increasingly negative with increasing temperature,
the entropic term is the major contribution to the
Gibbs energy of ionization above about 200°C, i.e.
the absolute values of T·ΔSaº >>ΔHaº. This effect
has been observed for a great many acids and bases
[15, 16]. It arises from long-range polarization of
water by the charged anion and hydrated proton, an
effect which becomes much more important at
elevated temperature as the dielectric constant of
water decreases. While the relative order of ΔHaº
and ΔSaº for the three isomers at temperatures
above 200°C is the same as that at room
temperature, ΔSaº(m-NphOH) < ΔSaº(p-NphOH) ≈
ΔSaº(o-NphOH), the relative position of the orthoisomer at intermediate temperature changes
significantly, so that ΔSaº(m-NphOH) < ΔSaº(oNphOH) < ΔSaº(p-NphOH). This effect, which lies
just outside the combined estimates of estimated
uncertainty, may reflect the competing effects of
localized hydrogen bonding about the uncharged
species and solvent polarization effects.
Temperatures in the range t ≈ 100°C to t ≈ 250°C
are typically considered a transition region, in
which water behaves like a “normal” hydrogenbonded solvent, so that both effects are often of
similar magnitude.
Functions of the standard partial molar heat
capacity and volume functions, ΔCp,aº and ΔVaº,
show tendencies that are associated with solvent
polarization effects. There is the transition with
increasing temperature from the distinctly different
experimental values for each isomer at ambient
temperatures, through a maximum, towards the
large negative values associated with solvent
Figure 6: Standard partial molar heat capacity,
ΔCp,aº, and standard Gibbs energy of ionization,
ΔGaº, of nitrophenols as a function of temperature
at psat. Full line: o-nitrophenol; long-dashed line: mnitrophenol; short-dashed line: p-nitrophenol
there is considerable uncertainty in the hightemperature values, so that the results for the three
isomers above ~ 150°C agree to within the rather
large combined error limits. Measurements are
under way to determine the standard partial molar
heat capacity and volume of each isomer and its
conjugate base, in order to confirm whether
significant differences do exist at elevated
temperatures.
Acknowledgement
This work was supported by the International
Association for the Properties of Water and Steam
(IAPWS Young Scientist Fellowship for J.E.); the
Natural Sciences and Engineering Research
Council of Canada (NSERC); and by the Research
Centre of “Advanced Remedial Technologies and
Processes” at the University of Liberec.
Literature
5
[13]H.P. Hopkins, W.C. Duer and F.J. Millero:
Heat capacity changes for the ionization of
aqueous phenols at 25 °C. J. Solution Chem. 5,
263–268 (1976)
[14]C.L. Liotta, A. Abidaud and H.P. Hopkins:
Thermodynamics of acid-base equilibriums. IV.
Influence of solvation factors on acidity.
Volumes of ionization of the m- and p-isomers
of nitrophenol and formylphenol in water at 25
o
C. J. Am. Chem. Soc. 94, 8624 (1972)
[15]R.E. Mesmer, W.L. Marshall, D.A. Palmer,
J.M.
Simonson
and
H.F.
Holmes:
Thermodynamics of aqueous association and
ionization reactions at high temperatures and
pressures. J. Solution Chem. 17, 699–718 (1988)
[16]P.R. Tremaine, P. Bénézeth, C. Xiao and K.
Zhang, in: D.A. Palmer, R. Fernandez-Prini,
A.H. Harvey (eds.) Aqueous Systems at
Elevated Temperatures and Pressures:
Physical Chemistry in Water, Steam and
Aqueous Solutions, Chapt. 13. Elsevier,
Amsterdam (2004)
[1] J. Tremp, P. Mattrel, S. Fingler and W. Giger:
Phenol and nitrophenol as tropospheric
pollutants: emissions from automobile exhaust
and phase transfer in the atmosphere. Water Air
Soil Pollut. 68, 113–123 (1993)
[2] T.-W. Shin, K. Kim and I.-J. Lee:
Spectrophotometric determination of the acid
dissociation constants for cacodylic acid and pnitrophenol at elevated temperatures. J. Sol.
Chem., 26, 379-390 (1997)
[3] J. Ehlerova, L. Trevani, J. Sedlbauer and P.R.
Tremaine: Spectrophotometric determination of
the
ionization
constants
of
aqueous
nitrophenols at temperatures up to 225 ±C. J.
Solution. Chem., 37, 857-874 (2008)
[4] E. Bulemela, L. Trevani and P.R. Tremaine:
Ionization constants of aqueous glycolic acid at
temperatures up to 250 ±C using hydrothermal
pH indicators and UV-visible spectroscopy. J.
Sol. Chem., 34, 769-787 (2005)
[5] J.C. Tanger and H.C. Helgeson: Calculation of
the thermodynamic and transport properties of
aqueous species at high pressures and
temperatures: Revised equations of state for the
standard partial molal properties of ions and
electrolytes. Am. J. Sci. 288, 19–98 (1988)
[6] W.L. Marshall and E.U. Franck: Ion product of
water substance, 0–1000 °C, 1–10,000 bars.
New international formulation and its
background. J. Phys. Chem. Ref. Data, 10,
295–304 (1981)
[7] J. Sedlbauer, J.P. O’Connell and R.H. Wood: A
new equation of state for correlation and
prediction of standard molal thermodynamic
properties of aqueous species at high
temperatures and pressures. Chem. Geol. 163,
43–63 (2000)
[8] C. M. Judson and M. Kilpatrick: The effects of
substituents on the dissociation constants of
substituted
phenols.
I.
Experimental
measurements in aqueous solutions. J. Am.
Chem. Soc. 71, 3110–3115 (1949)
[9] R. A. Robinson and A. Peiper: The ionization
constant of m-nitrophenol from 5 to 50 °C. J.
Phys. Chem. 67, 2860–2861 (1963)
[10]K.S. Pitzer: In: Pitzer, K.S. (ed.) Activity
Coefficients in Electrolyte Solutions, 2nd edn.
CRC Press, Boca Raton (1991)
[11]P.M. Dewick: Essentials of Organic Chemistry.
Wiley, New York (2006). Chap. 4
[12]L.P. Fernandez and L.G. Hepler: Heats and
entropies of ionization of phenol and some
substituted phenols. J. Am. Chem. Soc. 81,
1783–1786 (1959)
6