REACTIVE
&
FUNCTIONAL
POLYMERS
ELSEVIER
Reactive & Functional Polymers 36 (1998) 217-220
Thermodynamics
of proton dissociation of acridinium ion
in aqueous solution
Maria P. Ros *, Jesus Thomas, Guillermo Crovetto, Juan Llor
Departamento de Quimica-Fhica,
Universidad de Granada, 18071 Granada, Spain
Received 15 December 1996; revised version received 20 January 1997; accepted 6 May 1997
Abstract
Standard thermodynamic function changes (AGO”, AHe’, A&” and Acre”) are reported for proton dissociation
of acridinium cation in aqueous solution. These values have been determined from a series of precise and accurate
acid equilibrium constants, K,. The equilibrium constants have been obtained spectrophotometrically over a range of
temperatures from 2 to 66°C at several ionic strengths (I = 0.005, 0.007, 0.01, 0.05 and 0.25 M). Although the ionic
strength does not affect in general the Gibbs free-energy values, the results suggest a possible influence on entropy and
enthalpy changes. 0 1998 Elsevier Science B.V. All rights reserved.
Keywords: Acridine; Thermodynamic functions; Proton dissociation
1. Introduction
The current interest of the acridine and its
derivatives in pharmaceutical use [l], as labels
in biological assays [2] and DNA probe-based
assays [3] makes determination
of ionization
constants and its corresponding thermodynamic
properties useful.
The dissociation equilibrium constant of acridinium ion (K,) at 20°C and ionic strength 0.01
M has been reported by Albert and Serjeant [4].
In fact, that we found nothing in the literature
studies either about the effect of the temperature
and ionic strength or about the standard thermodynamic function changes in this process proves
that there is, in general, little reliable information
* Corresponding author. Fax: +34 58 272879.
available about the thermodynamic parameters of
the dissociation process. For this reason, most
discussions of the relative acid strengths have
been based only on p K data.
Because the magnitude of a pK value is determined by the relative magnitudes of AH” and
AS” values, a knowledge of these quantities becomes important for an understanding of acid
dissociation. In the same way, relatively little
work [5-71 has been done on the effect of the
ionic strength.
In this article a systematic study of the influence of the ionic strength on the parameters
of the dissociation reaction has been carried out.
The pK, values of acridinium ion in aqueous
solutions have been determined as a function of
ionic strength (ranging from 0.005 to 0.25 M)
and temperature (2 to 66°C). From these values,
the standard thermodynamic function changes of
1381-5148/98/$19.00 0 1998 Elsevier Science B.V. All rights reserved.
PIZ Sl381-5148(97)00084-9
218
MI! Ros et al. /Reactive & Functional Polymers 36 (1998) 217-220
the dissociation process at different experimental
conditions have been calculated.
2. Experimental
Acridine (CtsHgN, IO-azaanthracene)
from
Aldrich was crystallized by sublimation before
being used. Stock solutions of acridine (0.0136
M) were prepared in 0.0284 M hydrochloric acid
solutions. Solutions of acridine at pH around the
pK values were prepared by dilution of the stock
solution of acridine in different buffers. The final
concentration of acridine was 2 x 10e4 M. Citric
acid and dibasic sodium phosphate were used as
buffered solutions. Acid and basic forms of acridine were obtained in HCl and NaOH solutions,
respectively.
pK, values were obtained at different and
strictly controlled ionic strength. The ionic
strength was calculated taking into account the
pH values of the solution and the pK values of
the buffers components and then adjusted to the
desired value with solid KCI. The resulting pH
values slightly change with the addition of solid
KCl. In the case of acid and basic acridine solutions the necessary quantities of HCl and NaOH
were used in order to obtain the desired ionic
strength.
All the solutions used were freshly prepared
with bid&tilled and deionized water obtained
from Millipore MilliQ system: 18 MQ . cm. The
absorbance of the buffered acridine solutions was
recorded on a Perkin-Elmer spectrophotometer.
In each measurement, the temperature was kept
constant (error less than O.l’C) through a flow
system.
The dissociation equilibrium constants of acridinium ion were determined by the general
method of Albert and Sarjeant [4] in buffered
solutions in aqueous media at five ionic strengths
(0.005, 0.007, 0.01, 0.05 and 0.25 M) and temperatures ranging from 2 to 66°C. In the region
380-440 nm the acid and base forms of the
a&dine show an important difference in its UVVisible spectra where the neutral molecule was
found to have only a low absorption. Therefore,
the absorbance of the solutions at 403 nm were
recorded as a function of pH. The buffered solutions did not absorb in this region and in addition
a reference cell was used in all cases.
pK, values have been obtained using the general equation [4]:
=pK,-pH
where Aa, At, and A are the absorbance of acid, basic and buffered solutions, respectively. The leastsquares fits of the log ((A - Ab)/(Aa - A)) vs.
pH were carried out using at least six data points,
for all the temperatures and ionic strengths. The
plots were linear (r2 > 0.99) and the slopes were
around the value - 1. The p K, values were obtained from the intercepts. Typical standard errors
were f0.04 for pKa and the slope.
pH measurements were made using a Radiometer PHM64 pH meter equipped with a
GK2401C combined pH electrode. The combined pH electrode was calibrated with standard
buffered solutions (Crison). The temperature was
controlled with a Minitherm Hi875 1 termopar.
Standard thermodynamic function changes at
the reference temperature 8 = 298.15 K have
been calculated by fitting the pK data to the
Clarke and Glew equation [8]:
PK~=A+;+cx
(1)
- R; B = AH~“/2.303 *
where A = -A&“/2.303
. R.
R; C = -AC&‘/2.303
3. Results and discussion
Fig. 1 shows the dependence of the obtained
pK, values on the temperature at some ionic
strengths. The value given by Albert and Sarjeant
[4] at 20°C and I = 0.01 M agrees very well with
our results. The errors in the determination of the
pK values were in general less than 0.03.
As can be seen in Fig. 1, there is no significant effect of the ionic strength on the pK,
values. In fact, if we compare the pK, values at a
given temperature, these are practically constant
M.P. Ros et al. /Reactive
& Functional
Polymers
219
36 (1998) 217-220
Table 1
Dependence
of the standard thermodynamic
(0 = 298.15 K) of the dissociation process
with the ionic strength.
5.0
t
I
0.29
I
I
0.35
0.32
l/T.
function changes
of acridinium ion
:M)
ASO”
(J/K. mol)
AHs”
(kJ/mol)
AGO”
(kJ/mol)
A$”
(UK . mol)
0.005
0.007
0.01
0.05
0.25
All data
-2Of3
-16f2
-21+4
-26&t
-5f4
-22f2
25.5fO.l
26.7 k 0.7
25 zk 1
23.9 + 0.8
30+ 1
25.2 & 0.6
31.60f0.06
31.33 + 0.05
31.47 f 0.08
31.84 & 0.05
31.97 f 0.05
31.66 f 0.04
360f90
150 f 100
890 z!=90
345 xt 60
I
0.38
The indicated
errors are standard
errors
1Oa
Fig. 1. Dependence of the pK, of acridinium ion with temperature at some ionic strengths (I, . = 0.005; V = 0.01; A = 0.05
M). The solid line is the theoretical line calculated by fitting all
the pK, data (70 points) to Eq. 1. In order to show a clear figure
we have not plotted all the data points.
with the ionic strength taking into account the
experimental error. This behaviour is as expected
considering the fact that there is no net change in
charge during the dissociation of the acridinium
cation.
The Clarke and Glew equation (Eq. 1) is a
Taylor’s series and provides a general representation of pK data as a function of temperature. This
equation contains undetermined constants which
separately define the values for the standard thermodynamic function changes for the reaction at a
chosen temperature 0 = 298.15 K. A regression
analysis has been made increasing successively
the number of terms in Eq. 1 and stopping when
the F-test indicated that a new term was not
statistically significative. The calculated thermodynamic parameters are listed in Table 1.
The pK, series of values (12-14 data points)
obtained at each ionic strength separately were
analyzed. In the case of Z = 0.005, 0.01 and
0.05 M the results of the F-test indicate that
AC,e” values given in Table 1 are statistically
significative, but not in the case of Z = 0.007 and
0.25 M where AC,0 ’ values must be considered
less reliable. In Fig. 2 the plot of pKa vs. l/T
(I = 0.05 M) is clearly curved, indicating that
the contribution of the AC,e” must be calculated.
The values given in Table 1 are good agreement with other published data. For acridine [9],
AS” = -12.58 J/K. mol, and AH” = 34.78
kJ/mol at 15°C. Slightly different values were
obtained for these parameters that could be due
to the contribution of ACpe” parameter and to
different experimental conditions. Other similar
compounds have also a similar behaviour. Thus,
for quinoline [9], AS0 = -16.51 J/K.mol, and
AH” = 22.42 kJ/mol at 25°C; isoquinoline [9],
AS” = -13.79 J/K.mol, and AH” = 24.77
kJ/mol at 25°C; 5,6-benzoquinoline [9], AS’ =
-26.96 J/K.mol, and AH” = 18.98 kJ/mol at
20°C; and pyridine [9], Z = 0, the results ranging for AS” = -30.1 to -32.4 J/K. mol, and
0.29
0.32
0.35
0.38
l/T.lO’
Fig. 2. Plot of pK, vs. l/T at I = 0.05 M. The theoretical lines
has been calculated from Eq. 1 using (1) two terms, (2) three
tWtllS.
220
M.P. Ros et al. /Reactive & Functional Polymers 36 (1998) 217-220
for AH” = 20.6 to 19.5 kJ/mol at 25°C although
otherauthors [lo] giveforpyridine
AS” = -38.8
JKmol, and AH” = 18.3 kJ/mol at 25°C.
The dissociation constants of different classes
of acids show significant differences in their thermodynamic origins. In our case 90% of the free
energy is due to the enthalpy term and only a
small entropy (10%) contribution, following the
typical behaviour of the ionization of ammonium
ions.
From the results given in Table 1 we can see
that there is not an important effect of the ionic
strength on AGe” values and this parameter could
be considered constant within the experimental
error. In this case the errors obtained are small
and therefore, it can be said that this parameter is
really constant at all the ionic strengths.
Due to the fact that AHe” values have been
calculated from the dependence of the equilibrium constants with the temperature these magnitudes are affected of higher errors (Table 1) and
thus we can only affirm that AH@‘, and AS@“,
remain constant with the ionic strength within
the experimental error. In any case, the results
given in Table 1 seem to indicate that the AHs”
and ASo” values could be affected by the ionic
strength. However, we cannot confirm this fact
considering only the present results. In order to
study the possible effect of the ionic strength on
the enthalpy and entropy of the process is necessary to carry out more precise measurements of
the A He” values using calorimetric techniques,
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