# Protonation Equilibria of Biologically Active Ligands in Mixed Aqueous Organic Solvents

Authors: Ahmed A. El-Sherif; Mohamed M. Shoukry; Abeer T. Abd Elkarim; Mohammad H. Barakat
Journal: Bioinorganic Chemistry and Applications
 (2014)

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## Abstract

The review is mainly concerned with the protonation equilibria of biologically active ligands like amino acids, peptides, DNA constituents, and amino acid esters in nonaqueous media. Equilibrium concentrations of proton-ligand formation as a function of pH were investigated. Also, thermodynamics associated with protonation equilibria were also discussed.


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## Body

## 1. Introduction

There is a growing interest in studying the properties of biologically active ligands like amino acids, peptides, DNA, and amino acid esters. Amino acids and peptides have special importance among the other chemical groups since they are the foundation stones of the living organisms. They are not only components of tissues, but also reactive organic compounds which are important regulators of biological processes. It is obvious that one has to know the physical properties of amino acids and peptides in order to explain the behavior and synthesis of proteins and enzymes in the organisms. The study of amino acids, peptides, or DNA units has been the subject of increasing research efforts, which have revealed the role of hydrogen ion at molecular level. These compounds are usually considered as good model systems to attain a better insight into the characteristics of naturally occurring copper metalloproteins [1]. The data related to the protonation constants of biologically active ligands in various media will be valuable in further understanding of their chemistry in biological systems. Also, analytical chemists are supposed to know the related constants of the species present in the medium to determine the accuracy and most suitable medium for their analysis.The major reasons for the determination of protonation constants can be summarized as follows: one can calculate the pH and the ratio of different forms of a certain substance by the use of its protonation constants. Due to the fact that different forms of different substances have different UV spectra, by choosing a suitable pH value one can carry out spectrophotometric quantitative analyses. The choice of the pH value requires knowledge of protonation constants. Protonation constants are important in preparative chemistry as well. If the protonation constants of a certain substance are known, it is possible to isolate it with a maximum yield by finding the pH range where the compounds show minimum ionization. Protonation of a newly synthesized compound can also give supportive information about its structure. If theoretically calculated protonation constants are in good accordance with the experimental values, it is possible that the proposed structure could be correct. It is necessary that the protonation constants be known in order to prepare buffer solutions at different pH values [2]. There are various techniques such as conductometry, spectrophotometry, and potentiometry [3–5] that are used in the determination of protonation constants. In addition, for the calculations of stability constants of the complex formation of biologically active ligands with metal ions, their protonation constants are used [6]. Knowledge of the equilibrium constants of some compounds is necessary for the calculation of the concentration of each ionized species at any pH, which is important for the complete understanding of the physicochemical behavior of such molecules [6].Mostly, physical constants of bioligands are studied in aqueous media. Despite this, little is known about the chemistry of biomolecules in the mixtures of organic solvents and water, in regard to either protonation constants or synthetic applications. One reason for this dearth of knowledge is thatin vivo reactions take place in aqueous media, so that interest in bioligands properties in aqueous solution has predominated. However, in the literature it has been shown that water is not an ideal model forin vivo reactions. In enzymes, membranes, and other biologically important media, the p
K
a values are far different from those in water, as these media tend to be lipophilic rather than hydrophilic [7, 8]. Studies in media other than water should provide some understanding of the chemistry of bioligands in living systems. With this in mind and in continuation of our research program directed to study the solution equilibria of biologically active compounds [9–15], the present review discusses the valuable information existing in the literature for some bioactive ligands like amino acids, amino acid esters, peptides, and DNA constituents in nonaqueous media.


## 2. Determination of Protonation Constants of Biologically Active Ligands

The determination of equilibrium constants is an important process for many branches of chemistry [16]. Developments in the field of computation of equilibrium constants from experimental data were reviewed by Legget [17] and Meloum et al. [18]. Since that time, many more programs have been published, mainly so as to be able to use microcomputers for the computations. The most commonly used programs for solution equilibrium constant determination are PKPOT [19], PKAS [20], BEST [21], MINIQUAD [22], MINIQUAD-75 [23], SUPERQUAD [24], PSEQUAD [25], and HYPERQUAD [26]. All of these programs use least-square refinements to reduce the differences between calculated and experimental data to get the best model that gives the best fit. The sum of square of residuals between experimental and calculated values is normally very small; it is typically between 10−6 and 10−9.


## 3. Protonation Equilibria in Nonaqueous Solutions

Vast data are available on the protonation and stability constants of the bioligands in aqueous solutions [27–40]. But the knowledge which may be obtained from mixed solvents could complement the vast amount of data collected from studies on the chemistry of bioligands in aqueous media. Water-organic solvent mixtures have attracted much interest because of their frequent use and wide field of applications. These solvent mixtures are also used as reaction media for a variety of organic and analytical processes such as synthesis, titrations, or liquid chromatographic separations. Solvent molecules act as proton acceptors and therefore the solvent plays a crucial role in chemical and biochemical reactions. The acid-base behavior of substances in these mixtures is of considerable interest and importance. Data related to protonation constants will be valuable in the further understanding of biological systems. Hence, studies in media other than water should provide some understanding of the chemistry of bioligands in living systems.


### 3.1. Autoprotolysis Constants (K
ap)

The autoprotolysis constant (K
ap) of a solvent is an important parameter in understanding acid-base equilibria in mixed solvents. It also determines the extreme limits of acidity and basicity in a given solvent medium. The dielectric constant and the intrinsic acidic and basic strength of the solvent will influence the self-ionization process in these media. The recent IUPAC document on the criteria for standardization of pH measurements in organic solvent and water-organic solvent mixtures underlines the importance of knowing the autoprotolysis constant (in fact, -
log
⁡
10
⁡
K
ap
= normal scale length of pH) pertaining to each solvent considered, which can be achieved by appropriate thermodynamic methods [41]. Various methods have been used to obtain values of the autoprotolysis constants in aqueous organic solvent mixtures by potentiometry [42]. These experimental methods are difficult, and those methods that make use of the hydrogen electrode can only be applied to systems in which there is no complication due to reduction at the hydrogen electrode. Furthermore, Rosés et al. [43, 44] derived equations to calculate autoprotolysis constants of aqueous organic solvent mixtures and they expressed that all of the autoprotolysis processes are determined by the polarity and hydrogen bounding capability of solvents. Kiliç and Aslan [45] derived a convenient and rapid potentiometric technique using a combined glass-pH electrode for the determination of autoprotolysis constants in a variety of aqueous organic mixed solvents. For different compositions of organic solvent-water mixtures, reproducible values of autoprotolysis constants were calculated from several separate series of [H
+] and [OH
-] measurements as given in Tables 1 and 2 [45–51].Table 1
The autoprotolysis constants (K
ap) obtained for various water-solvent mixtures.
Medium
p
K
ap
Water
13.78a, 13.69b, 13.77c
20% ethanol—80% water
14.03d
30% ethanol—70% water
14.17a, 14.16b
40% ethanol—60% water
14.30d
50% ethanol—50% water
14.40a, 14.28b
60% ethanol—40% water
14.48d
70% ethanol—30% water
14.67a, 14.59b
12.5% dioxane—87.5% water
14.17e
20% dioxane—80% water
14.31f, 14.22b
25% dioxane—75% water
14.37e
37.5% dioxane—62.5% water
14.50
40% dioxane—60% water
14.89f, 14.86b
50% dioxane—50% water
15.44e
60% dioxane—40% water
15.54f, 15.62b
62.5% dioxane—37.5% water
15.70e
70% dioxane—30% water
15.75e
50% DMSO—50% water
15.48g
aData taken from reference [38], bdata taken from reference [37],cdata taken from reference [39], ddata taken from reference [40], edata taken from reference [41], fdata taken from reference [42], and gdata taken from reference [43].Table 2
The autoprotolysis constants (K
ap) obtained for various water-organic solvent mixtures measured in an ionic medium of 0.1 mol·dm−3 NaClO4 at 25°Ca.
Organic component
Organic component %
0
10
20
30
40
50
60
70
80
Methanol
13.69
13.75
13.78
13.70
13.73
13.71
13.72
13.77
13.94
Ethanol
13.69
13.82
13.99
14.16
14.24
14.28
14.39
14.59
14.77
1-Propanol
13.69
13.78
13.98
14.04
14.15
14.31
14.46
14.61
14.82
2-Propanol
13.69
13.76
13.95
14.05
14.26
14.34
14.56
14.79
—
Dioxane
13.69
13.91
14.22
14.51
14.86
15.24
15.62
—
—
Acetonitrile
13.69
13.96
14.19
14.44
14.74
15.04
—
—
—
aData taken from reference [37].
The % is expressed as (volume/Volume).


### 3.2. Protonation Equilibria in Mixed Ethanol-Water Medium

Among the organic solvents, ethanol is closest to water in structure and properties and therefore the behaviour of dissociation or protonation in ethanol is similar to that of aqueous solution. The water-ethanol mixture is a very interesting binary mixture. One reason is that ethanol can dissolve a majority of organic acids and bases more effectively than water. However, it has been suggested that solvents such as water-ethanol mixtures provide a better model forin vivo reactions [52, 53] because the mixtures simultaneously show a low polar character and partially aqueous contents, as do all biological systems. In addition, ethanol-water mixtures are a suitable solvent for the determination of equilibrium constants [53–56]. In addition, there has been increasing interest during the past few years in the properties of mixed-solvent systems such as polar aprotic solvents and water. Water and organic solvents such as methanol, ethanol, acetone, and acetonitrile are completely miscible. Their mixtures are macroscopically homogeneous, but it has been reported that the water and organic solvent molecules are not homogeneously dispersed microscopically, because of the hydrogen-bonding network formation and hydrophobic interactions [46]. Consequently, the molecular composition of the solvation layer around a solute molecule is not the same as that of the bulk mixing ratio of water and organic solvent. The hydrophobic and hydrophilic properties of a solute may be reflected by preferential solvation in such mixed solvents. Therefore, the influence which the solvent exerts on the log
⁡
10
⁡
K values depends upon the extent and nature of the solute-solvent interaction, which involves species participating in the acid-base equilibrium. Also, analytical chemists are supposed to know the related constants of the species present in the medium to determine the accuracy and most suitable medium for their analysis. Protonation constants of some α-amino acids at different water-ethanol mixtures are given in Tables 3–7 [45, 55, 57–63].Table 3
Protonation constants of someα-amino acids in water and water-ethanol mixtures [Temp. = 25°C and I
=
0.1 mol·dm−3 NaCl].
AA
Water
30% E—70% W
50% E—50% W
70% E—30% W
log
⁡
K
1
log
⁡
K
2
log
⁡
K
3
log
⁡
K
1
log
⁡
K
2
log
⁡
K
3
log
⁡
K
1
log
⁡
K
2
log
⁡
K
3
log
⁡
K
1
log
⁡
K
2
log
⁡
K
3
Cys
10.29
8.36
1.71
10.25
8.23
2.60
10.81
7.06
2.93
12.07
6.33
2.82
Tyr
10.14
9.03
2.17
10.46
9.03
2.48
10.90
9.06
3.25
10.70
8.80
2.69
Trp
9.33
2.35
—
9.17
2.83
—
9.05
3.21
—
8.78
2.77
—
Lys
10.69
9.08
2.04
10.20
8.77
1.71
10.07
8.76
2.68
9.64
8.48
2.19
His
9.08
6.02
1.70
8.86
6.21
3.00
8.53
5.85
2.93
8.32
5.73
2.99
Note. E: ethanol, W: water, Cys = L-cysteine, Tyr = L-tyrosine, Trp = L-tryptophane, Lys = L-lysine, and His = L-histidine. aData taken from reference [50]; log
⁡
K
1: corresponds to 11 species (i.e., L−  +  H+
⇌ LH); log
⁡
K
2 corresponds to 12 species (i.e.,  LH  +  H+
⇌
L
H
2
+). The % is expressed as (volume/Volume).Table 4
Protonation constants of someα-amino acids in water and water-ethanol mixtures [Temp. = 25°C and I
=
0.1 mol·dm−3 NaClO4
]
a.
AA
30% E—70% W
40% E—60% W
50% E—50% W
60% E—40% W
70% E—30% W
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
Gly
9.54
2.81
9.39
2.89
9.35
2.90
9.20
3.05
9.05
3.06
Ala
9.70
2.88
9.55
2.97
9.45
3.05
9.18
3.06
9.12
3.08
Val
9.50
2.85
9.41
2.95
9.30
3.15
9.07
3.16
9.02
3.17
Leu
9.60
2.87
9.40
3.01
9.30
3.06
9.07
3.17
9.05
3.25
Ile
9.55
2.92
9.47
2.98
9.31
3.15
9.20
3.17
9.10
3.24
Phe
9.12
2.75
8.95
2.95
8.79
2.98
8.70
3.05
8.60
3.00
Ser
9.18
2.88
9.00
2.87
8.77
2.89
8.68
2.91
8.66
2.90
Met
9.05
2.67
8.99
2.85
8.85
2.88
8.67
3.00
8.62
3.02
Note. E: ethanol, W: water, Val = L-valine, Ala = L-alanine, Leu = L-leucine, Ile = L-isoleucine, Phe = L-phenylalanine, Ser = L-serine, and Thr = L-threonine; adata taken from reference [49]. log
⁡
K
1 corresponds to 11 species (i.e., L−  +  H+
⇌  LH); log
⁡
K
2 corresponds to 12 species (i.e.,  LH  +  H+
⇌
L
H
2
+). The % is expressed as (volume/Volume).Table 5
Protonation constants of dipeptides in water and water-ethanol mixture[Temp. = 25°C and I
=
0.1 mol·dm−3 NaCl].
Dipeptides
Water
20% E—80% Wc
40% E—60% Wc
60% E—40% Wc
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
Gly-Gly
8.15a
3.11b
7.66
3.33
6.92
3.54
6.98
3.81
GlyVal
8.23d
3.15d
7.67
3.52
7.06
3.91
6.92
4.18
GlyLeu
7.91e
3.15e
7.60
3.38
7.17
3.94
7.02
4.22
GlyPhe
7.98d
3.03d
7.62
3.33
6.96
3.65
6.98
3.91
GlyThr
8.50c
3.20c
7.74
3.19
7.07
3.37
7.05
3.67
GlyMet
8.22f
3.11f
7.64
3.32
7.02
3.62
6.96
3.91
Note. E: ethanol; W: water; adata taken from reference [51]; bdata taken from reference [52]; cdata taken from reference [42]; ddata taken from reference [53]; edata taken from reference [54]; fdata taken from reference [55]. log
⁡
K
1 corresponds to 11 species (i.e., L−  +  H+
⇌   LH); log
⁡
K
2 corresponds to 12 species (i.e.,  LH  +  H+
⇌
L
H
2
+). The % is expressed as (volume/Volume).Table 6
Protonation constants ofα-amino acid esters in water-ethanol mixtures [Temp. = 25°C and I
=
0.1 mol·dm−3 NaCl]
a.
Esters
Water
30% E—70% W
50% E—50% W
70% E—30% W
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
CysOMe
9.17
6.38
9.23
6.09
9.34
5.88
10.82
6.24
CysOEt
9.36
6.54
9.37
6.17
9.48
5.97
11.13
6.28
TyrOMe
9.73
7.04
10.10
6.73
10.49
6.42
10.62
6.17
TyrOEt
9.71
7.05
10.13
6.75
10.52
6.45
10.63
6.19
TrpOMe
10.66
7.10
11.27
6.78
11.56
6.51
11.39
6.29
TrpOEt
10.79
7.10
11.28
6.98
11.70
6.80
11.77
6.45
LysOMe
9.99
6.98
9.80
6.90
9.52
6.63
9.12
6.46
LysOEt
10.32
7.18
9.70
6.81
9.45
6.60
9.00
6.35
HisOMe
7.10
4.96
6.72
4.78
6.54
4.61
6.34
4.39
Note. E = ethanol; W = water;  adata taken from reference [56]. log
⁡
K
1 corresponds to 11 species (i.e., L−  +  H+
⇌  LH); log
⁡
K
2 corresponds to 12 species (i.e.,  LH  +  H+
⇌
L
H
2
+). The % is expressed as (volume/Volume).Table 7
Protonation constants ofα-amino acid-t-butyl and benzyl esters at 25°C for different water-EtOH mixtures [
I
=
0.1 mol·dm−3 NaCl]
a.
Esters
20% E—80% W
X
=
0.0331
30% E—70% W
X
=
0.1167
40% E—60% W
X
=
0.1740
50% E—50% W
X
=
0.2356
60% E—40% W
X
=
0.3161
70% E—30% W
X
=
0.4183
80% E—0% W
X
=
0.5521
GlyO-t-Bu
7.66
7.51
7.40
7.30
7.20
7.05
6.64
AlaO-t-Bu
9.01
8.86
8.75
8.55
8.39
8.16
8.14
ValO-t-Bu
7.74
7.56
7.15
6.95
6.85
6.51
6.50
LeuO-t-Bu
7.79
7.63
7.38
7.21
7.02
6.92
6.75
PheO-t-Bu
7.15
7.00
6.67
6.56
6.50
6.40
6.33
IleO-t-Bu
7.75
7.65
7.35
7.18
7.12
6.89
6.64
GlyO-Bz
7.27
—
7.07
—
6.94
—
6.85
AlaO-Bz
7.35
—
7.17
—
7.01
—
7.00
ValO-Bz
7.16
—
6.88
—
6.77
—
6.68
SerO-Bz
6.75
—
6.60
—
6.51
—
6.54
Note. E = ethanol; W = water; adata taken from ref. [57]. log
⁡
K
1 corresponds to 11 species (i.e., L−  +  H+
⇌  LH); log
⁡
K
2 corresponds to 12 species (i.e.,  LH  +  H+
⇌
L
H
2
+). The % is expressed as (volume/Volume).


#### 3.2.1. Protonation Equilibria of Amino Acids in Mixed Water-Ethanol Medium

(1) Protonation Equilibria of L-Cysteine and L-Tyrosine in Mixed Water-Ethanol Medium (Figure1). The protonation constants of L-cysteine and L-tyrosine given in Table 3 correspond to the following equilibria:
(1)
L
2
-
+
H
+
⇌
HL
-
K
1
=
[
H
L
-
]
[
L
2
-
]
[
H
+
]
H
L
-
+
H
+
⇌
H
2
L
K
2
=
[
H
2
L
]
[
H
L
-
]
[
H
+
]
H
2
L
+
H
+
⇌
H
3
L
+
K
3
=
[
H
3
L
+
]
[
H
2
L
]
[
H
+
]Figure 1Thelog
⁡
10
⁡
K
1 values are related to the attachment of H+ to the phenolic oxygen in L-tyrosine and attachment of H
+ to the –NH2 group in L-cysteine. The log
⁡
10
⁡
K
2 value corresponds to the attachment of a proton to the –NH2 and sulfur thiol groups in L-tyrosine and L-cysteine, respectively. The log
⁡
10
⁡
K
3 values are the smallest and are thought to correspond to the protonation of carboxyl groups. The decrease in the log
⁡
10
⁡
K
2 value with increasing ethanol concentration can be explained by the fact that the dipolar ionic form is more strongly solvated than the anionic form in aqueous media.(2) Protonation Equilibria of L-Tryptophane in Mixed Water-Ethanol Medium (Figure2). Two protonation equilibria were found for L-tryptophan. The first protonation constant was attributed to the protonation of –NH2 and the second one was related to protonation of carboxyl group. No protonation constants were found related to the protonation of the nitrogen atom of the indole ring of this amino acid. The decrease in log
⁡
10
⁡
K
1 values of L-tryptophane was observed with the increase of ethanol concentration. This can be explained by better solvation of the dipolar ionic form.Figure 2(3) Protonation Equilibria of L-Lysine and Histidine in Mixed Water-Ethanol Medium (Figure3). L-lysine contains –NH2 group in the side chain. Due to the fact that this amino group is far from the carboxyl groups, resulting in a lesser degree of inductive effect, it can be said that log
⁡
10
⁡
K
1 is related to the protonation of this amino group. The second and third protonation equilibria belong to the α-amino and carboxyl groups, respectively. For L-histidine, the log
⁡
10
⁡
K
1 and log
⁡
10
⁡
K
2 values are related to the protonation of –NH2 and the nitrogen in imidazole groups, respectively. The third constant (log
⁡
10
⁡
K
3) belongs to the protonation of the carboxylate group.Figure 3Thelog
⁡
10
⁡
K
1 and log
⁡
10
⁡
K
2 values of L-lysine decrease with increasing ethanol concentration in the solvent mixture. Although, in the first equilibrium involving K
1, HL seems to be an uncharged species, it belongs to the dipolar form of amino acid. If we consider the solvation of these species, the fact that both K
1 and K
2 decrease with increasing ethanol concentration seems highly reasonable. The fact that the log
⁡
10
⁡
K
1 value of L-histidine decreases with increasing ethanol concentration can be explained by the fact that dipolar HL is better solvated in water than anionic form (L
-). The reason for the decrease in log
⁡
10
⁡
K
2 is that solvation of the multicharged species formed by the attachment of a proton to the imidazole ring is very limited in ethanol.(4) Protonation Equilibria of Someα-Amino Acids in Mixed Water-Ethanol Medium. The numerical log
⁡
10
⁡
K
1 and log
⁡
10
⁡
K
2 values of eight α-amino acids (glycine, L-alanine, L-valine, L-leucine, L-isoleucine, L-phenylalanine, L-serine, and L-methionine) determined in water-ethanol mixtures are given in Table 4. When the change of log
⁡
10
⁡
K
1 given in Table 4 with the solvent composition is examined, it is observed that these values decrease with the increase in the mole fraction of ethanol. This can be explained by the structural changes in amino acids as the medium changed toward the ethanol-rich direction. The neutral form amino acid HL is subjected to higher solvation compared with L
- forms in ethanol rich media. In water-rich media, however, the reverse behavior will be the case. This shows that log
⁡
10
⁡
K
1 values are expected to decrease with an increase in ethanol ratio. Therefore, one can conclude that the dipolar form of amino acids, HL
±, dominates in these media. It is the only way to explain the decrease in log
⁡
10
⁡
K
1 as the amount of ethanol in the media is increased since L
- is expected to solvate more than HL
± in media rich in ethanol.The fact thatlog
⁡
10
⁡
K
2 values are observed to increase as the mole fraction of ethanol is increased in all media investigated supports that dipolar ion is predominant in water-ethanol mixtures as well as in water. This is due to the fact that ethanol solvated H
2
L better than HL
±. If the neutral form were the predominant species in water-ethanol mixture, log
⁡
10
⁡
K
2 values would decrease with increasing amount of ethanol since ethanol solvated HL better than H
2
L
+ [64, 65].


#### 3.2.2. Protonation Equilibria of Dipeptides in Mixed Water-Ethanol Medium

Two protonation equilibria were found for all glycine-dipeptides (Figure4) mentioned in Table 5.Figure 4The first protonation constant was attributed to the protonation of amino group and the second one was related to the protonation of carboxylate group:(2)
L
-
+
H
+
⇌
HL
K
1
=
[
HL
]
[
L
-
]
[
H
+
]
HL
+
H
+
⇌
H
2
L
+
K
2
=
[
H
2
L
+
]
[
HL
]
[
H
+
]
where L is dipeptide monoanion; HL is total zwitterion and neutral forms of dipeptide.Chattopadhyay and Lahiri [66] examined the effect of a change of solvent composition on BH
+ ionization and the related Gibbs transfer energies in mixed solvent. They emphasized that the electrostatic charge effects due to changes in the dielectric constants with changes in solvent composition are of minor importance in explaining solvent effects, and that solute-solvent interactions have greater significance in interpretation of solvent effects.These protonation constants have been considered in some detail to gain more information about the effect of solvent composition on the corresponding equilibria. For this purpose, by examining the protonation equilibria of dipeptides with solvent composition, it is generally observed thatlog
⁡
10
⁡
K
1 values decrease with the increase in the mole fraction of ethanol, whereas log
⁡
10
⁡
K
2 values increase as the composition of ethanol increases in water-ethanol mixture. The decrease in log
⁡
10
⁡
K
1 values and increase in log
⁡
10
⁡
K
2 values of all dipeptides with the increasing ethanol concentration were the same as for the corresponding free amino acids [48].The fact that the variation oflog
⁡
10
⁡
K
1 and log
⁡
10
⁡
K
2 of dipeptides with solvent composition is similar to those of the corresponding free amino acids and can be concluded from the domination of the zwitterionic form of dipeptides in water-ethanol mixtures. In water-rich media, however, the reverse behavior will be the case. This shows that log
⁡
10
⁡
K
1 values are expected to increase with an increase in the ethanol ratio. This is the only way which explains the decrease in log
⁡
10
⁡
K
1 as the amount of ethanol in the media is increased because L
- is expected to solvate more than the zwitterionic form of dipeptides in the media rich in ethanol. The fact that log
⁡
10
⁡
K
2 values are observed to increase as the mole fraction of ethanol is increased in all media investigated supports that dipolar ion is predominant in water-ethanol mixtures as well as in water. This is due to the fact that ethanol solvated H
2
L better than zwitterionic form. If the neutral form were the predominant species in water-ethanol mixture, log
⁡
10
⁡
K
2 values would decrease with increasing amount of ethanol because ethanol solvated HL better than H
2
L
+. Thus, we can say that the log
⁡
10
⁡
K
2 values for these dipeptides increase as the mole fraction of ethanol increases.In order to compare thelog
⁡
10
⁡
K
1 and log
⁡
10
⁡
K
2 dipeptides with the log
⁡
10
⁡
K
1 value of glycine and log
⁡
10
⁡
K
2 value of the other amino acid forming the dipeptide the protonation constants of the corresponding amino acids are given in Table 5. According to the data given in Table 5, the log
⁡
10
⁡
K
2 values of amino acids increase and the log
⁡
10
⁡
K
1 values of glycine decrease due to peptide formation. For instance log
⁡
10
⁡
K
1 and log
⁡
10
⁡
K
2 values for glycine in 60% water-40% ethanol medium were 9.39 and 2.89 [55, 67], respectively, while same values were found as 6.92 and 3.54 for Gly-Gly. Similarly log
⁡
10
⁡
K
1 for glycine was 9.20 [55] and log
⁡
10
⁡
K
2 for phe was 3.05 [57] in 40% water-60% ethanol medium while these values were observed to be 6.98 and 3.91 for GlyPhe in the same medium. This may be attributed to the fact that the electron density of the amino group in dipeptide is lower than that of glycine, and the electron density of the carboxylate group is higher than that of the corresponding amino acid due to the formation of a peptide bond.Estimation of equilibrium concentrations of proton-ligand formation as a function of pH provides a useful picture of proton-ligand binding in media [16]. Figures 5, 6, 7, and 8 show the change of different species of dipeptides with pH. It is seen from these figures that the most predominant species is HL (total concentration of the zwitterion and the neutral form) between pH 5 and 7 in all ethanol-water media.Figure 5
Species distribution diagram (25°C,I
=
0.1 M NaCl) for Gly-Gly system as a function of pH in water (L
=
1.5
×
10
-
3 M).Figure 6
Species distribution diagram (25°C,I
=
0.1 M NaCl) for Gly-Gly system as a function of pH in 20% ethanol-80% water (L
=
1.5
×
10
-
3 M).Figure 7
Species distribution diagram (25°C,I
=
0.1 M NaCl) for Gly-Gly system as a function of pH in 40% ethanol-60% water (L
=
1.5
×
10
-
3 M).Figure 8
Species distribution diagram (25°C,I
=
0.1 M NaCl) for Gly-Gly system as a function of pH in 60% ethanol-40% water (L
=
1.5
×
10
-
3 M).


#### 3.2.3. Protonation Equilibria of Amino Acid Esters in Ethanol-Water Medium

The changes of the protonation constants of methyl and ethyl esters of L-tyrosine, L-tryptophane, L-cysteine, L-lysine, and methyl ester of L-histidine with increasing ethanol percentage are given in Table6. It can be seen that log
⁡
10
⁡
K
1 values of L-tyrosine, L-tryptophane, and L-cysteine esters increase with increasing percentage of ethanol. On the other hand, the log
⁡
10
⁡
K
2 values show a decrease with increasing ethanol concentration. The log
⁡
10
⁡
K
1 and log
⁡
10
⁡
K
2 values of L-lysine and L-histidine esters both decrease with increasing ethanol concentration in the solvent mixture.The protonation constants of methyl and ethyl esters of L-cysteine, L-tyrosine, and L-tryptophane are given in Table6. The log
⁡
10
⁡
K
1 values belong to equilibria related to the attachment of H
+ to the sulfur atom in the thiol group of cysteine esters, to the nitrogen atom in the indole ring of tryptophane, and to the phenolic oxygen group in tyrosine esters. The log
⁡
10
⁡
K
2 values are related to the protonation of –NH2 groups in the esters. The increase in the ethanol concentration decreases the dielectric constant of the medium. The amino acid ester (HL) is subjected to stronger solvation compared with the (L
-) and (H
2
L
+) forms in ethanol-rich media. In water-rich media, however, the reverse behavior will be the case. This shows that log
⁡
10
⁡
K
1 values are expected to increase and log
⁡
10
⁡
K
2 values to decrease with an increase in the ethanol ratio.Thelog
⁡
10
⁡
K
1 and log
⁡
10
⁡
K
2 values of L-lysine esters are related to the protonation of amino groups in side chain and α-amino groups, respectively. For L-histidine esters, these values correspond to the protonation of nitrogen in the imidazole and α-amino group in L-histidine. Contrary to the other esters investigated both log
⁡
10
⁡
K
1 and log
⁡
10
⁡
K
2 values of lysine methyl and ethyl esters and histidine methyl ester decrease with an increase in the ratio of ethanol in the solvent mixture. Since H
2
L
2
+ is better solvated than HL
+ and the latter is better solvated than L-molecular species in water-rich media, therefore, results obtained for these esters can be explained by the specific solvation effect.The stoichiometric protonation constants for t-butyl esters and benzyl esters ofα-amino acids determined in ethanol-water mixtures (20–80 volume % ethanol) are given in Table 7. These values are related to this equilibrium reaction L
+
H
+
⇌
L
H
+, where L and LH
⁡
+ show α-amino acid esters and their protonated species, respectively. The protonation constants given in Table 7 are considered in more detail in order to gain more information about the effect of solvent composition and specific effects of substituents on the basicities of the esters in solvent mixtures. The numerical log
⁡
10
⁡
K values for t-butyl esters of α-amino acids determined in ethanol-water mixtures decrease with increasing ethanol content in the solvent mixture. It is observed that a nearly linear relationship exists between the aforementioned protonation constants and the mole fraction of ethanol from 0.0331 to 0.4183 for all α-amino acid esters investigated. However, log
⁡
⁡
β values at a mole fraction of ethanol of 0.4183 are slightly higher than those expected from the linear trend. The linear variation for all α-amino acid esters that are given in Table 7 is very similar to that found for ammonia aliphatic alkyl amines, pyridine, and salicylideneanilines [68, 69]. The dissociation constants of charged acids in water-ethanol mixtures vary with solvent composition in a manner that is not completely understood. It is suggested that electrostatic charging effects resulting from the change in dielectric constant with solvent effects and the solute-solvent interactions have greater significance in the interpretation of solvent effects. Thus, results obtained for α-amino acid esters in water-ethanol medium can be explained by specific solvation effects. Since ethanol would solvate L better than LH
⁡
+, the log
⁡
⁡
β values, which are related to the formation of LH
⁡
+, would decrease upon addition of ethanol. The derivations of linearity in 80% ethanol may result from the preferential solvation of solute by one of the components of the solvent mixture that could change the effective dielectric constant value in the cybotactic region [70]. Furthermore, another factor which explains the increase in the log
⁡
⁡
β values of all α-amino acid esters in ethanol-rich regions is the differences in the solvent stabilization of the ionic species (H
+ and LH
⁡
+), brought about by changing the percentage of ethanol [5, 71]. Using the protonation constants obtained in this paper, the effects of the type of ester groups on the basicity of the amine groups of α-amino acid esters have been discussed. The most important factor that affects the basicity and therefore the protonation constant of a compound is the structural effect. Table 7 showed that the basicity of t-butyl esters of glycine, L-alanine, and L-valine is higher than that of the corresponding benzyl esters of the same amino acids in 20%–60% ethanol-water mixtures. This effect can be explained by taking the electronic effect of the t-butyl and benzyl groups investigated into account.


### 3.3. Protonation Equilibria in Mixed Water-Dioxane Medium

Simple heterocyclic compounds, such as 1,4-dioxane, in pesticides are very important to living organisms and environment. Also, it is well established that the “effective” or “equivalent solution” dielectric constants in protein [72, 73] or active site cavities of enzymes [74] are small compared to that in bulk water. Estimates for the dielectric constants in such locations range from 30 to 70 [73, 75]. Hence by using aqueous solutions containing ~10–50% dioxane, one may expect to simulate to some degree the situation in active site cavities [76] and hence to extrapolate the data to physiological conditions. Mixed solvents such as water-dioxane mixtures provide a better model forin vivo reactions because the mixtures simultaneously show a lesser polar character and are partially aqueous contents, as do all biological systems [2]. On the other hand, water-1,4-dioxane mixtures are a very interesting binary solvent system. These mixtures are a favorite mixed solvent system in which to study association and mobilities of ions, because the dielectric constant can be varied over a large range. Changes in protonation constants upon addition of 1,4-dioxane to aqueous solutions are due to increasing ion-ion interactions resulting from the decreasing dielectric constant and from changes in solvent-ion and solvent-solvent interactions [77–79]. Also, the protonation constant of amino acid esters will be helpful in the determination of the microscopic equilibrium constants of the corresponding amino acids. Protonation constants of some α-amino acids at different water-dioxane mixtures are given in Tables 8 and 9 [49, 67].Table 8
Protonation constants of someα-amino acids in water and water-dioxane mixtures [Temp. = 25°C and I
=
0.1 mol·dm−3 NaClO4
]
a.
AA
Water
10% D—90% W
20% D—80% W
30% D—70% W
40% D—60% W
50% D—50% W
60% D—40% W
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
Gly
9.58
2.32
9.55
2.41
9.56
2.70
9.58
2.87
9.58
3.15
9.60
3.48
9.62
3.60
Ala
9.70
2.40
9.66
2.52
9.72
2.64
9.68
2.89
9.73
3.10
9.75
3.39
9.78
3.63
Val
9.61
2.38
9.56
2.53
9.60
2.68
9.58
2.86
9.60
3.08
9.62
3.43
9.66
3.65
Leu
9.75
2.50
9.59
2.64
9.64
2.80
9.66
2.99
9.70
3.31
9.72
3.55
9.75
3.75
Ile
9.76
2.47
9.72
2.58
9.63
2.74
9.67
2.97
9.65
3.26
9.70
3.55
9.71
3.74
Phe
9.20
2.43
9.15
2.57
9.10
2.64
9.12
2.81
9.15
3.10
9.16
3.36
9.15
3.50
Ser
9.15
2.42
9.07
2.57
9.10
2.62
9.13
2.82
9.16
3.04
9.20
3.17
9.17
3.35
Thr
9.04
2.45
9.02
2.53
9.06
2.60
9.05
2.75
9.07
2.96
9.10
3.18
9.17
3.40
Note. D = dioxane, W = water, Val = DL-valine, Ala = DL-alanine, Leu = L-leucine, Ile = L-isoleucine,  Phe = DL-phenylalanine, Ser = L-serine, and Thr = L-threonine.  aData taken from reference [61]. log
⁡
K
1 corresponds to 11 species (i.e., L−  +  H+
⇌   LH); log
⁡
K
2  corresponds to 12 species (i.e.,  LH  +  H+
⇌
L
H
2
+). The % is expressed as (volume/Volume).Table 9
Protonation constants ofα-amino acid esters in water-1,4-dioxane  mixtures [Temp. = 25°C and I
=
0.1 mol·dm−3 NaCl]
a.
Esters
p
q
a
Water
20% D—80% W(X
=
0.05)
40% D—60% W(X
=
0.12)
60% D—40% W(X
=
0.24)
log
⁡
K
log
⁡
K
log
⁡
K
log
⁡
K
GlyO-Me
1
1
7.67
7.47
7.18
6.85
GlyO-t-Bu
1
1
8.38
8.21
7.91
6.97
ValO-Me
1
1
7.53
7.38
7.09
6.08
ValO-Et
1
1
8.86
8.01
7.15
6.58
ValO-t-Bu
1
1
9.05
8.54
7.71
7.13
SerO-Me
1
1
7.10
6.99
6.69
6.61
SerO-Et
1
1
7.46
7.25
6.84
6.79
LeuO-Me
1
1
7.66
7.47
7.19
6.32
LeuO-Et
1
1
7.75
7.78
7.31
6.73
LeuO-t-Bu
1
1
7.89
8.04
7.60
6.95
PheO-Me
1
1
7.11
6.66
6.42
6.01
PheO-Et
1
1
8.12
7.24
6.81
6.24
PheO-t-Bu
1
1
8.26
7.39
6.96
6.64
Note. D = dioxane, W = water; adata taken from ref. [44].  log
⁡
K
1 corresponds to 11 species (i.e., L−  +  H+
⇌    LH).


#### 3.3.1. Protonation Equilibria ofα-Amino Acids in Mixed Dioxane-Water Medium

The stoichiometric protonation constants of eightα-amino acids determined in various water-dioxane mixtures are listed in Table 8. The log
⁡
10
⁡
K
1 and log
⁡
10
⁡
K
2 values given in Table 8 refer to the protonation equilibria shown in Figure 9, respectively.Figure 9Many studies have shown that the equilibrium constant is linearly related to the fraction of organic solvent [80–83]. The results obtained in this study confirm this observation. Regarding the variation of log
⁡
10
⁡
K
1 values with the solvent composition, one can postulate that the zwitterionic to neutral form ratio decreases as the dioxane content increases. This can be inferred from considerations on specific solute-solvent interactions and structural changes in amino acids from water to dioxane-water media. In the related equilibrium, the HL form would be more solvated than the L
- and the HL
± forms in solvents containing a higher concentration of dioxane, whereas the opposite forms would hold for a solvent rich in water. If the zwitterionic to neutral form ratio (K
Z) was unaltered with increasing percentage of dioxane, the log
⁡
10
⁡
K
1 values would be expected to decrease, which is not observed in the present case. If the neutral form became the predominant species when dioxane content is increased, the log
⁡
10
⁡
K
1 values would increase as in the case of a neutral monobasic acid, whereas they are not altered [84]. Thus, even with the limitations, we can conclude that the K
Z values for all amino acids studied decrease as the concentration of dioxane increases. The fact that the log
⁡
10
⁡
K
2 values increase with increasing the dioxane content for all the amino acids suggests that the zwitterion is the predominant species in dioxane-water mixtures as well as in water. This is based on the assumption that dioxane solvates H
2
L
+ better than HL
±. If the neutral form was the predominant species in water-dioxane mixtures, the log
⁡
10
⁡
K
2 values would decrease on addition of dioxane since dioxane would solvate HL better than H
2
L
+.


#### 3.3.2. Protonation Equilibria of Amino Acid Esters in Mixed Water-Dioxane Medium

The stoichiometric protonation constants of theα-amino acid esters and the effect of the solvents upon the values of these protonation constants of α-amino acid esters were discussed in water-1,4-dioxane mixtures and the results are tabulated in Table 9. These values are related to the protonation of –NH2 group of the esters, as follows:
(3)
where B and BH
+ denote the α-amino acid esters and their protonated species, respectively. The protonation constants given in Table 9 were considered in more detail in order to gain more information about the effect of solvent composition and specific effects of substituents on the basicities of the esters in solvent mixtures. When it comes to the variation of log
⁡
⁡
β with the solvent composition, it can be observed that these constants decrease as the concentration of 1,4-dioxane increases in going from water to 60% in dioxane-water mixtures. It is also observed that a nearly linear relationship exists between the protonation constants and the mole fraction of 1,4-dioxane between 0.05 and 0.24 for all α-amino acid esters investigated. Many studies have shown that the equilibrium constant is linearly related to the fraction of organic solvent [80–83]. Although the effect of solvent composition upon the protonation constants of various compounds has been the subject of extensive investigation, the variation of charged acid-base equilibrium constants with solvent composition is yet to be fully elucidated. The solvation of a solute in a mixed solvent is much more complicated than solvation in a single pure solvent, and in the literature there are several different explanations. In organic solvent-water mixtures, acid-base equilibria can be estimated from two effects: an electrostatic one that can be explained by the Barron equation and a nonelectrostatic one that includes specific solute-solvent interaction [85]. Bates [71] and Chattopadhyay and Lahiri [66] have examined the effect of changing solvent composition on ionization of BH+ and the related Gibbs energies of transfer in mixed solvents, and they found that the dielectric constant alone cannot serve for a quantitative explanation for the solvent effect [66, 71]. However, the solvent mixtures are characterized by dramatic changes in the physical constants of the solvent (i.e., their melting and boiling point, dielectric constants, etc.) and macroscopic solvent parameters (mole fraction of cosolvent, etc.), upon changes of their compositions that influence many solute properties of these mixtures. Also, according to Takamuku et al. [86], the 1,4-dioxane-water mixtures can be divided into three mole fraction ranges, that is, 0.3
≤
X
dioxane
≤
0.9, X
dioxane, 0.1
≤
X
dioxane
≤
0.2, and 0
≤
X
dioxane
≤
0.07. When 0.3
≤
X
dioxane
≤
0.9, the structure observed for pure dioxane obviously remains the same as it is mixed with 0.1
≤
X
dioxane
≤
0.2, small binary clusters consisting of one or two 1,4-dioxane molecules and several water molecules are formed by hydrogen bonding. When 0
≤
X
dioxane
≤
0.07, first-, second-, and third-neighbor O
⋯
O interactions are observed due to the hydrogen-bonded network of water. This suggests that the addition of a small amount of 1,4-dioxane breaks down the hydrogen-bonded network of water. In dioxane-rich mixtures the protonated base (BH
+) is mostly solvated by the less basic dioxane molecules, whereas the proton is more solvated in the water-dioxane mixtures than in water and dioxane, because the water-dioxane mixture is more basic than either pure water or dioxane. Also, the molecular composition of the solvation layer around a solute molecule is not the same as that of the bulk mixture ratio of water to organic solvent. The hydrophobic and hydrophilic properties of a solute may be reflected in preferential solvation in such mixed solvents. Therefore, the influence that the solvent exerts on the p
K values depends upon the extent and nature of the solute-solvent interactions, which involves species participating in the acid-base equilibrium [72]. From the theoretical explanation mentioned above, we can explain results for all α-amino acid esters in terms of specific solvation effects. A study of specific solvation effects on the α-amino acid esters showed that the log
⁡
10
⁡
K values show a decrease upon increase of the 1,4-dioxane ratio for all α-amino acid esters investigated. Because 1,4-dioxane should solvate B better than BH
+, the log
⁡
10
⁡
K values, which are related to the formation of BH
+, should decrease upon addition of 1,4-dioxane. We can say that the deviations from the trend with the log
⁡
⁡
β values for all amino acid esters appeared in 1,4-dioxane-rich regions (x
=
0.24), resulting from differences in solvent stabilization of the ionic species (H
+ and BH
+) being brought about by changing the percentage of 1,4-dioxane. The results show that each methyl ester is generally a slightly weaker base than its corresponding ethyl and t-butyl analogues. In addition, the alanine methyl ester is a weaker base than its benzyl ester. The acidity or basicity of a compound in a given medium is influenced by both the electronic effect of substituents and the solvent effect of the medium. Moreover, it is sometimes extremely difficult to assess how much each effect contributes to the acidity or basicity. In the literature it has also been shown that the methyl derivatives of aliphatic amines are weaker bases than their ethyl derivatives in water-methanol mixtures [49]. We can say that the basicity of a compound is a result of various factors such as (i)
the solvent effect (solvation power, the tendency of forming hydrogen bonds, selective solvation, dielectric constant, and the composition of the solution in the first solvation layer) in the case of mixed solvents,
(ii)
structural effects, electronic effects, and steric effects.


#### 3.3.3. Protonation Equilibria of DNA in Mixed Water-Dioxane Medium

The solvent effect on the acid dissociation constants of a ligand [87] can be summarized as follows.(i)
As the solvent dielectric constant decreases, thelog
⁡
10
⁡
K of the ligand increases and vice versa.
(ii)
On decreasing the extent of hydrogen bonding in water by an organic solvent, the proton-accepting properties of the water increase, and consequently thelog
⁡
10
⁡
K of the ligand decreases.
(iii)
Increasing proton solvation by an organic solvent is accompanied by a decrease in thelog
⁡
10
⁡
K of ligand.The protonation constants of inosine and inosine 5′-monophosphate (Figure10) as representative examples of DNA constituents in different compositions of dioxane-water mixtures reveal the following features (Table 10) [88]: p
K
a (N1H) of inosine increases linearly with increasing percentage of organic solvent in the medium, Figure 11. This may be correlated with the ability of a solvent of relatively low dielectric constant to increase the electrostatic forces between the proton and the ligand anion and consequently the log
⁡
10
⁡
K value increases. The same trend was observed for the log
⁡
10
⁡
K’s of (N1H) and phosphoric acid (P-OH) groups of IMP, Figure 11. The log
⁡
10
⁡
K of the phosphate group in IMP is more affected by replacement of water molecules with dioxane. This is in agreement with the above finding, namely, the existence of hydrogen bonds between water molecules and phosphate ions. Thus, the phosphate ions are easily protonated in going from water to dioxane.Table 10
Protonation constants of inosine and IMP in dioxane-water media at 25°C andI
=
0.1 mol·dm−3 NaCl.
System
% dioxane
p
q
a
log
⁡
β
b
Inosine IMPc
0%
1
1
8.55
1
1
9.09
1
2
15.07
Inosine IMP
25%
1
1
8.85
1
1
9.12
1
2
15.83
Inosine IMP
37.5%
1
1
9.04
1
1
9.24
1
2
16.28
Inosine IMP
50%
1
1
9.14
1
1
9.35
1
2
16.63
Inosine IMP
62.5%
1
1
9.36
1
1
9.39
1
2
16.72
a
p and q are stoichiometric coefficients corresponding to ligand and H+, respectively; bdata taken from reference [83], cIMP =  inosine 5′-monophosphate. The % is expressed as (volume/Volume). Species 11 refers to L− + H+
⇌   LH; species 12 refers to L + 2H+
⇌
L
H
2
+.Figure 10Figure 11
Effect of dioxane on thep
K
a of inosine and inosine 5′-monophosphate (IMP).


### 3.4. Protonation Equilibria in Mixed Water-DMSO Medium

The use of this mixed solvent has some advantages over pure DMSO. Thus, pure DMSO is very hygroscopic and controlling its water content is difficult [89]. This fact would affect reproducibility of experimental results. However, water-DMSO 50% : 50% mixture has only small hygroscopic character. A further advantage is its compatibility with the standard glass electrode, so that the pH measurements may be carried out in a similar way to that employed in a purely aqueous solution. In contrast, the use of pure DMSO is not recommended for potentiometry. Another advantage of the water-DMSO 50% : 50% mixture is its large acidity range (p
K
w
=
15.50) which allows the investigation of deprotonation equilibria of weak acids which could be hardly studied in water [90, 91]. Protonation constants of some α-amino acids at different water-DMSO mixtures and different temperatures are given in Tables 11 and 12 [92, 93].Table 11
Protonation constants of someα-amino acids in water and DMSO-water media at 25°C and  I
=
0.1 mol·dm−3 Na
NO
3
a.
System
p
q
b
30% DMSO
50% DMSO
70% DMSO
Gly
1
1
9.51 ± 0.01
9.79 ± 0.01
10.29 ± 0.01
1
2
12.92 ± 0.01
13.85 ± 0.02
15.20 ± 0.01
L-Ala
1
1
9.65 ± 0.009
9.85 ± 0.02
10.48 ± 0.02
1
2
13.07 ± 0.01
13.96 ± 0.03
15.36 ± 0.03
L-Val
1
1
9.34 ± 0.01
9.57 ± 0.01
9.99 ± 0.01
1
2
12.61 ± 0.02
13.61 ± 0.02
14.88 ± 0.03
L-Ser
1
1
8.95 ± 0.01
9.12 ± 0.02
9.89 ± 0.01
1
2
12.16 ± 0.02
12.95 ± 0.03
14.71 ± 0.03
L-Isl
1
1
9.31 ± 0.01
9.51 ± 0.02
9.91 ± 0.02
1
2
12.75 ± 0.01
13.56 ± 0.03
15.02 ± 0.03
L-Leu
1
1
9.46 ± 0.01
9.77 ± 0.05
10.18 ± 0.01
1
2
12.80 ± 0.02
13.81 ± 0.05
14.98 ± 0.02
L-Phe
1
1
9.06d
9.35c
—
1
2
12.44
13.43
L-Met
1
1
9.07d
9.48c
—
1
2
12.69
13.48
L-Thr
1
1
8.95d
9.32c
—
1
2
12.30
13.38
aData taken from reference [87]; b
p and q are stoichiometric coefficients corresponding to ligand and H+, respectively; cdata taken from reference [88]. Species 11 refers to L− + H+
⇌  LH; species 12 refers to L + 2H+
⇌
L
H
2
+; the % is expressed as (volume/Volume).Table 12
Protonation constants (log
10
⁡
β
q
,
r) of some α-amino acids in solution containing 50% DMSO at different temperatures and I
=
0.1 mol·dm−3 NaNO3 and their thermodynamic parametersb.
System
p
q
a
20°C
25°C
30°C
35°C
Thermodynamicsd
Δ
H
°
Δ
S
°
Δ
G
°
Gly
1
1
10.04
9.79
9.53
9.30
−85.71
−100.14
−55.87
1
2
14.14
13.85
13.52
13.22
−21.04
6.87
−23.09
L-Ala
1
1
10.02
9.85
9.69
9.46
−63.51
−24.63
−56.17
1
2
14.19
13.96
13.75
13.46
−19.34
13.80
−23.46
L-Val
1
1
9.79
9.57
9.49
9.26
−57.69
−9.63
−54.82
1
2
13.89
13.61
13.47
13.17
−21.75
4.31
−23.03
L-Ser
1
1
9.42
9.12
8.82
8.52
−103.66
−173.88
−52.02
1
2
13.30
12.95
12.60
12.25
−17.28
15.34
−21.85
L-Isl
1
1
9.72
9.51
9.39
9.20
−58.07
−12.28
−55.74
1
2
13.81
13.56
13.42
13.19
−11.06
40.55
−23.01
L-Leu
1
1
10.09
9.77
9.45
9.13
−110.57
−184.0
−55.74
1
2
14.18
13.81
13.42
13.04
−21.07
6.51
−32.01
aDMSO% (vol/vol), bdata taken from reference [87], c
q and r are stoichiometric coefficients corresponding to ligand and H+ respectively; dWith respect to the thermodynamic parameters (Δ
H
°, Δ
S
° and Δ
G
°) these parameters referred to these reactions (L−  +  H+
⇌  LH±, corresponding to –NH2 group) and (LH± + H+
⇌
L
H
2
+, corresponding to carboxylate group); The % is expressed as (volume/Volume).


#### 3.4.1. Protonation Equilibria of Amino Acids in Water-DMSO Medium

The stoichiometric protonation constants of glycine, L-alanine, L-phenylalanine, L-threonine, and L-methionine in water-DMSO mixtures are given in Table11. The protonation constants K
1 and K
2 are related to the protonation of the amino nitrogen and the carboxyl oxygen, respectively. When the change of log
⁡
10
⁡
K
1 with the solvent composition, given in Table 11, is examined for α-amino acids, it is observed that these values increase with an increasing percent of dimethyl sulfoxide. The linear relation is given in Figure 12, for glycine as a representative example of α-amino acids. This can be explained by structural changes in the amino acids as the medium becomes more like dimethyl sulfoxide. The neutral form of the amino acids (HL) is subjected to a larger amount of solvation, compared to the anionic form (L
-) in dimethyl sulfoxide rich media. In water-rich media, however, the reverse will be the case.Figure 12
The variation of protonation constants of glycine with the percentage of DMSO.Also, thelog
⁡
10
⁡
K
2 values increase for each α-amino acid as the percentage of dimethyl sulfoxide in the solvent mixtures increases. Many studies have shown that the equilibrium constant is linearly related to the fraction of the organic solvent [80, 81, 83]. In general, the protonation of compounds containing O–H increases with increasing organic content of the solvent, due to the decrease in the dielectric constant of the bulk solvents. As the dielectric constant decreases, the ion-ion interaction involving the proton and the anionic oxygen donor of the ligand increases to a greater extent than the ion-dipole interaction between the proton and the solvent molecule. Thus the log
⁡
10
⁡
K of carboxylic acid (COO−H+) groups of α-amino acids increases linearly with increasing percentage of organic solvent in the medium. The same trend was observed for the p
K
a’s of phosphoric acid (P-OH) groups of IMP [88].Serine isα-amino acid containing three functional groups (–OH, –NH2, and COOH), but only two protonation equilibria are determined corresponding to the protonation of amino and carboxylate groups, respectively. This can be attributed to the fact that, as the relative pH in most experiments did not extend to pH > 12, the protonation constant for the aliphatic OH group in serine could not be measured. Also, a qualitative explanation could be based on the formation of an intramolecular hydrogen bond as depicted in Figure 13. Assume that the proton on this –OH group is more difficult to remove in the amino acid form.Figure 13
Intramolecular H-bond of serine.Also, thelog
⁡
10
⁡
K
2 values increase for each α-amino acid as the percentage of dimethyl sulfoxide in the solvent mixtures increases. Many studies have shown that the equilibrium constant is linearly related to the fraction of the organic solvent [80, 81, 83].


### 3.5. Thermodynamic Quantities Associated with the Protonation Equilibria

Changes in the temperature may have some significant influences on the equilibrium properties and cause spontaneous changes in aquatic systems. Solution equilibria may be shifted to the right or to the left by an increase in temperature. The direction and magnitude of the shift are given by the sign and magnitude of the heat of the reaction as a function of temperature.The thermodynamic parameters are useful tools for studying these interactions with bioligands and understanding the relative stability of the formed complexes.The enthalpy changes accompanying complex formation can be obtained directly through calorimetric measurements, or indirectly from the temperature dependence of the stability constants via the Van’t Hoff equation. The accuracy ofΔ
H values based on calorimetric measurements is generally higher than the accuracy of those based on the temperature dependence of log
⁡
10
⁡
K values. However, many parameters influence the accuracy of calorimetric data, for example, the enthalpy of dissociation of water, the heat of dilution, and so forth [93].Application of the temperature dependence oflog
⁡
10
⁡
K generally requires the determination of very accurate stability constants over as wide range in temperature as possible. However, the wider the range employed, the greater the uncertainty in the calculated values, because of the temperature dependence of the enthalpy values.A high electrolyte concentration was used in order to keep variations of the activity coefficients at a minimum. Precise thermodynamic data can only be obtained provided that an inert electrolyte of fairly high concentration (≥0.1 mol·dm−3) is used. The main differences between different procedures for the study of ionic equilibria in aqueous and nonaqueous solutions are due to the activity coefficients. As in most equilibria in nonaqueous medium, a background electrolyte is added to maintain constant ionic strength ranging from about 0.1 to 0.3 mol·dm−3. This is allowed in some water-ethanol or water-dioxane mixtures, but not in solvents of low dielectric constants where the solubility of electrolyte is very low. In rigorous thermodynamic calculations, the equilibrium constant should be expressed in terms of activities of the component ions at equilibrium. However, for convenience, concentrations are generally used, where, for a particular species i, a
i
=
(
C
i
)
F
i, where C
i is the concentration of the ion i, a
i is its activity, and F
i is its activity coefficient. The activity coefficients are maintained constant by working in a medium of constant ionic strength.The results are summarized in Tables12 and 13 and can be interpreted as follows.(1)
The protonation reactions (L
-
+
H
+
⇌
L
H
±, corresponding to –NH2 group) and (L
H
±
+
H
+
⇌
L
H
2
+, corresponding to carboxylate group) in Table 12 of the investigated amino acids are exothermic (Figure 14). Three factors affect these protonation reactions:
(i)
the neutralization reaction, which is an exothermic process,
(ii)
desolvation of ions, which is an endothermic process,
(iii)
the change in the configuration and the arrangement of the hydrogen bonds around the free and protonated ligands.
(2)
Thelog
⁡
10
⁡
K values decrease with increasing temperature revealing that the acidity increases with increasing temperature (Tables 12 and 13).Table 13
Effect of temperature and thermodynamic parameters of protonation equilibria of inosine and IMPa.
System
T (°C)
p
q
b
log
⁡
β
Equilibrium
Δ
H
°  (kJ mol−1)
Δ
S
°   (JK−1 mol−1)
Δ
G
°  (kJ mol−1)
Inosine
15
1
1
8.77
Inosine
−35.0
46.3
−48.8
IMP
15
1
1
9.33
(1) L−  +  H+
⇌  LH
1
2
15.28
Inosine
20
1
1
8.67
IMP
20
1
1
9.17
1
2
15.17
Inosine
25
1
1
8.55
IMP
−33.0
64.6
−52.2
IMP
25
1
1
9.09
(2) L3−+  H+
⇌   LH2−
0.32
114.6
−33.8
1
2
15.07
(3) LH2− + H+
⇌
L
H
2
-
Inosine
30
1
1
8.44
IMP
30
1
1
8.95
1
2
14.99
a
p and q are stoichiometric coefficients corresponding to ligand and H+, respectively; bdata taken from reference [83]; L denotes inosine and inosine 5′-monophosphate.Figure 14
Van’t Hoff plot oflog
⁡
⁡
K of glycine in 50% DMSO-50% H2O mixture (log
⁡
⁡
K
1 denotes protonation of amino group and log
⁡
⁡
K
2 denotes protonation of carboxylate group).The protonation reactions (1) and (2) in Table13 [88] of the N1 site of inosine and inosine 5′-monophosphate are exothermic.The phosphate ion in IMP is more solvated than its protonated form. Consequently it contributes more to the endothermic process upon protonation and the netΔ
H
° is only 0.32 kJ mol−1. Also, the phosphate ions form ordered hydrogen bonds with water molecules, which is confirmed by the large entropy change Δ
S
°
=
114.6 JK−1 mol−1. This contributes to a negative free energy change Δ
G
°
=
33.8 kJ mol−1. The positive entropies due to the release of ordered water molecules and the breaking of hydrogen bonds were found by Kramer-Schnabel and Linder [73] to give positive Δ
H
°, large Δ
S
°, and net negative Δ
G
° for the protonation and complexation reactions of organic monophosphates and copper ions. The log
⁡
10
⁡
K values decrease with increasing temperature revealing that its acidity increases with increasing temperature.


## 3.1. Autoprotolysis Constants (K
ap)

The autoprotolysis constant (K
ap) of a solvent is an important parameter in understanding acid-base equilibria in mixed solvents. It also determines the extreme limits of acidity and basicity in a given solvent medium. The dielectric constant and the intrinsic acidic and basic strength of the solvent will influence the self-ionization process in these media. The recent IUPAC document on the criteria for standardization of pH measurements in organic solvent and water-organic solvent mixtures underlines the importance of knowing the autoprotolysis constant (in fact, -
log
⁡
10
⁡
K
ap
= normal scale length of pH) pertaining to each solvent considered, which can be achieved by appropriate thermodynamic methods [41]. Various methods have been used to obtain values of the autoprotolysis constants in aqueous organic solvent mixtures by potentiometry [42]. These experimental methods are difficult, and those methods that make use of the hydrogen electrode can only be applied to systems in which there is no complication due to reduction at the hydrogen electrode. Furthermore, Rosés et al. [43, 44] derived equations to calculate autoprotolysis constants of aqueous organic solvent mixtures and they expressed that all of the autoprotolysis processes are determined by the polarity and hydrogen bounding capability of solvents. Kiliç and Aslan [45] derived a convenient and rapid potentiometric technique using a combined glass-pH electrode for the determination of autoprotolysis constants in a variety of aqueous organic mixed solvents. For different compositions of organic solvent-water mixtures, reproducible values of autoprotolysis constants were calculated from several separate series of [H
+] and [OH
-] measurements as given in Tables 1 and 2 [45–51].Table 1
The autoprotolysis constants (K
ap) obtained for various water-solvent mixtures.
Medium
p
K
ap
Water
13.78a, 13.69b, 13.77c
20% ethanol—80% water
14.03d
30% ethanol—70% water
14.17a, 14.16b
40% ethanol—60% water
14.30d
50% ethanol—50% water
14.40a, 14.28b
60% ethanol—40% water
14.48d
70% ethanol—30% water
14.67a, 14.59b
12.5% dioxane—87.5% water
14.17e
20% dioxane—80% water
14.31f, 14.22b
25% dioxane—75% water
14.37e
37.5% dioxane—62.5% water
14.50
40% dioxane—60% water
14.89f, 14.86b
50% dioxane—50% water
15.44e
60% dioxane—40% water
15.54f, 15.62b
62.5% dioxane—37.5% water
15.70e
70% dioxane—30% water
15.75e
50% DMSO—50% water
15.48g
aData taken from reference [38], bdata taken from reference [37],cdata taken from reference [39], ddata taken from reference [40], edata taken from reference [41], fdata taken from reference [42], and gdata taken from reference [43].Table 2
The autoprotolysis constants (K
ap) obtained for various water-organic solvent mixtures measured in an ionic medium of 0.1 mol·dm−3 NaClO4 at 25°Ca.
Organic component
Organic component %
0
10
20
30
40
50
60
70
80
Methanol
13.69
13.75
13.78
13.70
13.73
13.71
13.72
13.77
13.94
Ethanol
13.69
13.82
13.99
14.16
14.24
14.28
14.39
14.59
14.77
1-Propanol
13.69
13.78
13.98
14.04
14.15
14.31
14.46
14.61
14.82
2-Propanol
13.69
13.76
13.95
14.05
14.26
14.34
14.56
14.79
—
Dioxane
13.69
13.91
14.22
14.51
14.86
15.24
15.62
—
—
Acetonitrile
13.69
13.96
14.19
14.44
14.74
15.04
—
—
—
aData taken from reference [37].
The % is expressed as (volume/Volume).


## 3.2. Protonation Equilibria in Mixed Ethanol-Water Medium

Among the organic solvents, ethanol is closest to water in structure and properties and therefore the behaviour of dissociation or protonation in ethanol is similar to that of aqueous solution. The water-ethanol mixture is a very interesting binary mixture. One reason is that ethanol can dissolve a majority of organic acids and bases more effectively than water. However, it has been suggested that solvents such as water-ethanol mixtures provide a better model forin vivo reactions [52, 53] because the mixtures simultaneously show a low polar character and partially aqueous contents, as do all biological systems. In addition, ethanol-water mixtures are a suitable solvent for the determination of equilibrium constants [53–56]. In addition, there has been increasing interest during the past few years in the properties of mixed-solvent systems such as polar aprotic solvents and water. Water and organic solvents such as methanol, ethanol, acetone, and acetonitrile are completely miscible. Their mixtures are macroscopically homogeneous, but it has been reported that the water and organic solvent molecules are not homogeneously dispersed microscopically, because of the hydrogen-bonding network formation and hydrophobic interactions [46]. Consequently, the molecular composition of the solvation layer around a solute molecule is not the same as that of the bulk mixing ratio of water and organic solvent. The hydrophobic and hydrophilic properties of a solute may be reflected by preferential solvation in such mixed solvents. Therefore, the influence which the solvent exerts on the log
⁡
10
⁡
K values depends upon the extent and nature of the solute-solvent interaction, which involves species participating in the acid-base equilibrium. Also, analytical chemists are supposed to know the related constants of the species present in the medium to determine the accuracy and most suitable medium for their analysis. Protonation constants of some α-amino acids at different water-ethanol mixtures are given in Tables 3–7 [45, 55, 57–63].Table 3
Protonation constants of someα-amino acids in water and water-ethanol mixtures [Temp. = 25°C and I
=
0.1 mol·dm−3 NaCl].
AA
Water
30% E—70% W
50% E—50% W
70% E—30% W
log
⁡
K
1
log
⁡
K
2
log
⁡
K
3
log
⁡
K
1
log
⁡
K
2
log
⁡
K
3
log
⁡
K
1
log
⁡
K
2
log
⁡
K
3
log
⁡
K
1
log
⁡
K
2
log
⁡
K
3
Cys
10.29
8.36
1.71
10.25
8.23
2.60
10.81
7.06
2.93
12.07
6.33
2.82
Tyr
10.14
9.03
2.17
10.46
9.03
2.48
10.90
9.06
3.25
10.70
8.80
2.69
Trp
9.33
2.35
—
9.17
2.83
—
9.05
3.21
—
8.78
2.77
—
Lys
10.69
9.08
2.04
10.20
8.77
1.71
10.07
8.76
2.68
9.64
8.48
2.19
His
9.08
6.02
1.70
8.86
6.21
3.00
8.53
5.85
2.93
8.32
5.73
2.99
Note. E: ethanol, W: water, Cys = L-cysteine, Tyr = L-tyrosine, Trp = L-tryptophane, Lys = L-lysine, and His = L-histidine. aData taken from reference [50]; log
⁡
K
1: corresponds to 11 species (i.e., L−  +  H+
⇌ LH); log
⁡
K
2 corresponds to 12 species (i.e.,  LH  +  H+
⇌
L
H
2
+). The % is expressed as (volume/Volume).Table 4
Protonation constants of someα-amino acids in water and water-ethanol mixtures [Temp. = 25°C and I
=
0.1 mol·dm−3 NaClO4
]
a.
AA
30% E—70% W
40% E—60% W
50% E—50% W
60% E—40% W
70% E—30% W
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
Gly
9.54
2.81
9.39
2.89
9.35
2.90
9.20
3.05
9.05
3.06
Ala
9.70
2.88
9.55
2.97
9.45
3.05
9.18
3.06
9.12
3.08
Val
9.50
2.85
9.41
2.95
9.30
3.15
9.07
3.16
9.02
3.17
Leu
9.60
2.87
9.40
3.01
9.30
3.06
9.07
3.17
9.05
3.25
Ile
9.55
2.92
9.47
2.98
9.31
3.15
9.20
3.17
9.10
3.24
Phe
9.12
2.75
8.95
2.95
8.79
2.98
8.70
3.05
8.60
3.00
Ser
9.18
2.88
9.00
2.87
8.77
2.89
8.68
2.91
8.66
2.90
Met
9.05
2.67
8.99
2.85
8.85
2.88
8.67
3.00
8.62
3.02
Note. E: ethanol, W: water, Val = L-valine, Ala = L-alanine, Leu = L-leucine, Ile = L-isoleucine, Phe = L-phenylalanine, Ser = L-serine, and Thr = L-threonine; adata taken from reference [49]. log
⁡
K
1 corresponds to 11 species (i.e., L−  +  H+
⇌  LH); log
⁡
K
2 corresponds to 12 species (i.e.,  LH  +  H+
⇌
L
H
2
+). The % is expressed as (volume/Volume).Table 5
Protonation constants of dipeptides in water and water-ethanol mixture[Temp. = 25°C and I
=
0.1 mol·dm−3 NaCl].
Dipeptides
Water
20% E—80% Wc
40% E—60% Wc
60% E—40% Wc
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
Gly-Gly
8.15a
3.11b
7.66
3.33
6.92
3.54
6.98
3.81
GlyVal
8.23d
3.15d
7.67
3.52
7.06
3.91
6.92
4.18
GlyLeu
7.91e
3.15e
7.60
3.38
7.17
3.94
7.02
4.22
GlyPhe
7.98d
3.03d
7.62
3.33
6.96
3.65
6.98
3.91
GlyThr
8.50c
3.20c
7.74
3.19
7.07
3.37
7.05
3.67
GlyMet
8.22f
3.11f
7.64
3.32
7.02
3.62
6.96
3.91
Note. E: ethanol; W: water; adata taken from reference [51]; bdata taken from reference [52]; cdata taken from reference [42]; ddata taken from reference [53]; edata taken from reference [54]; fdata taken from reference [55]. log
⁡
K
1 corresponds to 11 species (i.e., L−  +  H+
⇌   LH); log
⁡
K
2 corresponds to 12 species (i.e.,  LH  +  H+
⇌
L
H
2
+). The % is expressed as (volume/Volume).Table 6
Protonation constants ofα-amino acid esters in water-ethanol mixtures [Temp. = 25°C and I
=
0.1 mol·dm−3 NaCl]
a.
Esters
Water
30% E—70% W
50% E—50% W
70% E—30% W
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
CysOMe
9.17
6.38
9.23
6.09
9.34
5.88
10.82
6.24
CysOEt
9.36
6.54
9.37
6.17
9.48
5.97
11.13
6.28
TyrOMe
9.73
7.04
10.10
6.73
10.49
6.42
10.62
6.17
TyrOEt
9.71
7.05
10.13
6.75
10.52
6.45
10.63
6.19
TrpOMe
10.66
7.10
11.27
6.78
11.56
6.51
11.39
6.29
TrpOEt
10.79
7.10
11.28
6.98
11.70
6.80
11.77
6.45
LysOMe
9.99
6.98
9.80
6.90
9.52
6.63
9.12
6.46
LysOEt
10.32
7.18
9.70
6.81
9.45
6.60
9.00
6.35
HisOMe
7.10
4.96
6.72
4.78
6.54
4.61
6.34
4.39
Note. E = ethanol; W = water;  adata taken from reference [56]. log
⁡
K
1 corresponds to 11 species (i.e., L−  +  H+
⇌  LH); log
⁡
K
2 corresponds to 12 species (i.e.,  LH  +  H+
⇌
L
H
2
+). The % is expressed as (volume/Volume).Table 7
Protonation constants ofα-amino acid-t-butyl and benzyl esters at 25°C for different water-EtOH mixtures [
I
=
0.1 mol·dm−3 NaCl]
a.
Esters
20% E—80% W
X
=
0.0331
30% E—70% W
X
=
0.1167
40% E—60% W
X
=
0.1740
50% E—50% W
X
=
0.2356
60% E—40% W
X
=
0.3161
70% E—30% W
X
=
0.4183
80% E—0% W
X
=
0.5521
GlyO-t-Bu
7.66
7.51
7.40
7.30
7.20
7.05
6.64
AlaO-t-Bu
9.01
8.86
8.75
8.55
8.39
8.16
8.14
ValO-t-Bu
7.74
7.56
7.15
6.95
6.85
6.51
6.50
LeuO-t-Bu
7.79
7.63
7.38
7.21
7.02
6.92
6.75
PheO-t-Bu
7.15
7.00
6.67
6.56
6.50
6.40
6.33
IleO-t-Bu
7.75
7.65
7.35
7.18
7.12
6.89
6.64
GlyO-Bz
7.27
—
7.07
—
6.94
—
6.85
AlaO-Bz
7.35
—
7.17
—
7.01
—
7.00
ValO-Bz
7.16
—
6.88
—
6.77
—
6.68
SerO-Bz
6.75
—
6.60
—
6.51
—
6.54
Note. E = ethanol; W = water; adata taken from ref. [57]. log
⁡
K
1 corresponds to 11 species (i.e., L−  +  H+
⇌  LH); log
⁡
K
2 corresponds to 12 species (i.e.,  LH  +  H+
⇌
L
H
2
+). The % is expressed as (volume/Volume).


### 3.2.1. Protonation Equilibria of Amino Acids in Mixed Water-Ethanol Medium

(1) Protonation Equilibria of L-Cysteine and L-Tyrosine in Mixed Water-Ethanol Medium (Figure1). The protonation constants of L-cysteine and L-tyrosine given in Table 3 correspond to the following equilibria:
(1)
L
2
-
+
H
+
⇌
HL
-
K
1
=
[
H
L
-
]
[
L
2
-
]
[
H
+
]
H
L
-
+
H
+
⇌
H
2
L
K
2
=
[
H
2
L
]
[
H
L
-
]
[
H
+
]
H
2
L
+
H
+
⇌
H
3
L
+
K
3
=
[
H
3
L
+
]
[
H
2
L
]
[
H
+
]Figure 1Thelog
⁡
10
⁡
K
1 values are related to the attachment of H+ to the phenolic oxygen in L-tyrosine and attachment of H
+ to the –NH2 group in L-cysteine. The log
⁡
10
⁡
K
2 value corresponds to the attachment of a proton to the –NH2 and sulfur thiol groups in L-tyrosine and L-cysteine, respectively. The log
⁡
10
⁡
K
3 values are the smallest and are thought to correspond to the protonation of carboxyl groups. The decrease in the log
⁡
10
⁡
K
2 value with increasing ethanol concentration can be explained by the fact that the dipolar ionic form is more strongly solvated than the anionic form in aqueous media.(2) Protonation Equilibria of L-Tryptophane in Mixed Water-Ethanol Medium (Figure2). Two protonation equilibria were found for L-tryptophan. The first protonation constant was attributed to the protonation of –NH2 and the second one was related to protonation of carboxyl group. No protonation constants were found related to the protonation of the nitrogen atom of the indole ring of this amino acid. The decrease in log
⁡
10
⁡
K
1 values of L-tryptophane was observed with the increase of ethanol concentration. This can be explained by better solvation of the dipolar ionic form.Figure 2(3) Protonation Equilibria of L-Lysine and Histidine in Mixed Water-Ethanol Medium (Figure3). L-lysine contains –NH2 group in the side chain. Due to the fact that this amino group is far from the carboxyl groups, resulting in a lesser degree of inductive effect, it can be said that log
⁡
10
⁡
K
1 is related to the protonation of this amino group. The second and third protonation equilibria belong to the α-amino and carboxyl groups, respectively. For L-histidine, the log
⁡
10
⁡
K
1 and log
⁡
10
⁡
K
2 values are related to the protonation of –NH2 and the nitrogen in imidazole groups, respectively. The third constant (log
⁡
10
⁡
K
3) belongs to the protonation of the carboxylate group.Figure 3Thelog
⁡
10
⁡
K
1 and log
⁡
10
⁡
K
2 values of L-lysine decrease with increasing ethanol concentration in the solvent mixture. Although, in the first equilibrium involving K
1, HL seems to be an uncharged species, it belongs to the dipolar form of amino acid. If we consider the solvation of these species, the fact that both K
1 and K
2 decrease with increasing ethanol concentration seems highly reasonable. The fact that the log
⁡
10
⁡
K
1 value of L-histidine decreases with increasing ethanol concentration can be explained by the fact that dipolar HL is better solvated in water than anionic form (L
-). The reason for the decrease in log
⁡
10
⁡
K
2 is that solvation of the multicharged species formed by the attachment of a proton to the imidazole ring is very limited in ethanol.(4) Protonation Equilibria of Someα-Amino Acids in Mixed Water-Ethanol Medium. The numerical log
⁡
10
⁡
K
1 and log
⁡
10
⁡
K
2 values of eight α-amino acids (glycine, L-alanine, L-valine, L-leucine, L-isoleucine, L-phenylalanine, L-serine, and L-methionine) determined in water-ethanol mixtures are given in Table 4. When the change of log
⁡
10
⁡
K
1 given in Table 4 with the solvent composition is examined, it is observed that these values decrease with the increase in the mole fraction of ethanol. This can be explained by the structural changes in amino acids as the medium changed toward the ethanol-rich direction. The neutral form amino acid HL is subjected to higher solvation compared with L
- forms in ethanol rich media. In water-rich media, however, the reverse behavior will be the case. This shows that log
⁡
10
⁡
K
1 values are expected to decrease with an increase in ethanol ratio. Therefore, one can conclude that the dipolar form of amino acids, HL
±, dominates in these media. It is the only way to explain the decrease in log
⁡
10
⁡
K
1 as the amount of ethanol in the media is increased since L
- is expected to solvate more than HL
± in media rich in ethanol.The fact thatlog
⁡
10
⁡
K
2 values are observed to increase as the mole fraction of ethanol is increased in all media investigated supports that dipolar ion is predominant in water-ethanol mixtures as well as in water. This is due to the fact that ethanol solvated H
2
L better than HL
±. If the neutral form were the predominant species in water-ethanol mixture, log
⁡
10
⁡
K
2 values would decrease with increasing amount of ethanol since ethanol solvated HL better than H
2
L
+ [64, 65].


### 3.2.2. Protonation Equilibria of Dipeptides in Mixed Water-Ethanol Medium

Two protonation equilibria were found for all glycine-dipeptides (Figure4) mentioned in Table 5.Figure 4The first protonation constant was attributed to the protonation of amino group and the second one was related to the protonation of carboxylate group:(2)
L
-
+
H
+
⇌
HL
K
1
=
[
HL
]
[
L
-
]
[
H
+
]
HL
+
H
+
⇌
H
2
L
+
K
2
=
[
H
2
L
+
]
[
HL
]
[
H
+
]
where L is dipeptide monoanion; HL is total zwitterion and neutral forms of dipeptide.Chattopadhyay and Lahiri [66] examined the effect of a change of solvent composition on BH
+ ionization and the related Gibbs transfer energies in mixed solvent. They emphasized that the electrostatic charge effects due to changes in the dielectric constants with changes in solvent composition are of minor importance in explaining solvent effects, and that solute-solvent interactions have greater significance in interpretation of solvent effects.These protonation constants have been considered in some detail to gain more information about the effect of solvent composition on the corresponding equilibria. For this purpose, by examining the protonation equilibria of dipeptides with solvent composition, it is generally observed thatlog
⁡
10
⁡
K
1 values decrease with the increase in the mole fraction of ethanol, whereas log
⁡
10
⁡
K
2 values increase as the composition of ethanol increases in water-ethanol mixture. The decrease in log
⁡
10
⁡
K
1 values and increase in log
⁡
10
⁡
K
2 values of all dipeptides with the increasing ethanol concentration were the same as for the corresponding free amino acids [48].The fact that the variation oflog
⁡
10
⁡
K
1 and log
⁡
10
⁡
K
2 of dipeptides with solvent composition is similar to those of the corresponding free amino acids and can be concluded from the domination of the zwitterionic form of dipeptides in water-ethanol mixtures. In water-rich media, however, the reverse behavior will be the case. This shows that log
⁡
10
⁡
K
1 values are expected to increase with an increase in the ethanol ratio. This is the only way which explains the decrease in log
⁡
10
⁡
K
1 as the amount of ethanol in the media is increased because L
- is expected to solvate more than the zwitterionic form of dipeptides in the media rich in ethanol. The fact that log
⁡
10
⁡
K
2 values are observed to increase as the mole fraction of ethanol is increased in all media investigated supports that dipolar ion is predominant in water-ethanol mixtures as well as in water. This is due to the fact that ethanol solvated H
2
L better than zwitterionic form. If the neutral form were the predominant species in water-ethanol mixture, log
⁡
10
⁡
K
2 values would decrease with increasing amount of ethanol because ethanol solvated HL better than H
2
L
+. Thus, we can say that the log
⁡
10
⁡
K
2 values for these dipeptides increase as the mole fraction of ethanol increases.In order to compare thelog
⁡
10
⁡
K
1 and log
⁡
10
⁡
K
2 dipeptides with the log
⁡
10
⁡
K
1 value of glycine and log
⁡
10
⁡
K
2 value of the other amino acid forming the dipeptide the protonation constants of the corresponding amino acids are given in Table 5. According to the data given in Table 5, the log
⁡
10
⁡
K
2 values of amino acids increase and the log
⁡
10
⁡
K
1 values of glycine decrease due to peptide formation. For instance log
⁡
10
⁡
K
1 and log
⁡
10
⁡
K
2 values for glycine in 60% water-40% ethanol medium were 9.39 and 2.89 [55, 67], respectively, while same values were found as 6.92 and 3.54 for Gly-Gly. Similarly log
⁡
10
⁡
K
1 for glycine was 9.20 [55] and log
⁡
10
⁡
K
2 for phe was 3.05 [57] in 40% water-60% ethanol medium while these values were observed to be 6.98 and 3.91 for GlyPhe in the same medium. This may be attributed to the fact that the electron density of the amino group in dipeptide is lower than that of glycine, and the electron density of the carboxylate group is higher than that of the corresponding amino acid due to the formation of a peptide bond.Estimation of equilibrium concentrations of proton-ligand formation as a function of pH provides a useful picture of proton-ligand binding in media [16]. Figures 5, 6, 7, and 8 show the change of different species of dipeptides with pH. It is seen from these figures that the most predominant species is HL (total concentration of the zwitterion and the neutral form) between pH 5 and 7 in all ethanol-water media.Figure 5
Species distribution diagram (25°C,I
=
0.1 M NaCl) for Gly-Gly system as a function of pH in water (L
=
1.5
×
10
-
3 M).Figure 6
Species distribution diagram (25°C,I
=
0.1 M NaCl) for Gly-Gly system as a function of pH in 20% ethanol-80% water (L
=
1.5
×
10
-
3 M).Figure 7
Species distribution diagram (25°C,I
=
0.1 M NaCl) for Gly-Gly system as a function of pH in 40% ethanol-60% water (L
=
1.5
×
10
-
3 M).Figure 8
Species distribution diagram (25°C,I
=
0.1 M NaCl) for Gly-Gly system as a function of pH in 60% ethanol-40% water (L
=
1.5
×
10
-
3 M).


### 3.2.3. Protonation Equilibria of Amino Acid Esters in Ethanol-Water Medium

The changes of the protonation constants of methyl and ethyl esters of L-tyrosine, L-tryptophane, L-cysteine, L-lysine, and methyl ester of L-histidine with increasing ethanol percentage are given in Table6. It can be seen that log
⁡
10
⁡
K
1 values of L-tyrosine, L-tryptophane, and L-cysteine esters increase with increasing percentage of ethanol. On the other hand, the log
⁡
10
⁡
K
2 values show a decrease with increasing ethanol concentration. The log
⁡
10
⁡
K
1 and log
⁡
10
⁡
K
2 values of L-lysine and L-histidine esters both decrease with increasing ethanol concentration in the solvent mixture.The protonation constants of methyl and ethyl esters of L-cysteine, L-tyrosine, and L-tryptophane are given in Table6. The log
⁡
10
⁡
K
1 values belong to equilibria related to the attachment of H
+ to the sulfur atom in the thiol group of cysteine esters, to the nitrogen atom in the indole ring of tryptophane, and to the phenolic oxygen group in tyrosine esters. The log
⁡
10
⁡
K
2 values are related to the protonation of –NH2 groups in the esters. The increase in the ethanol concentration decreases the dielectric constant of the medium. The amino acid ester (HL) is subjected to stronger solvation compared with the (L
-) and (H
2
L
+) forms in ethanol-rich media. In water-rich media, however, the reverse behavior will be the case. This shows that log
⁡
10
⁡
K
1 values are expected to increase and log
⁡
10
⁡
K
2 values to decrease with an increase in the ethanol ratio.Thelog
⁡
10
⁡
K
1 and log
⁡
10
⁡
K
2 values of L-lysine esters are related to the protonation of amino groups in side chain and α-amino groups, respectively. For L-histidine esters, these values correspond to the protonation of nitrogen in the imidazole and α-amino group in L-histidine. Contrary to the other esters investigated both log
⁡
10
⁡
K
1 and log
⁡
10
⁡
K
2 values of lysine methyl and ethyl esters and histidine methyl ester decrease with an increase in the ratio of ethanol in the solvent mixture. Since H
2
L
2
+ is better solvated than HL
+ and the latter is better solvated than L-molecular species in water-rich media, therefore, results obtained for these esters can be explained by the specific solvation effect.The stoichiometric protonation constants for t-butyl esters and benzyl esters ofα-amino acids determined in ethanol-water mixtures (20–80 volume % ethanol) are given in Table 7. These values are related to this equilibrium reaction L
+
H
+
⇌
L
H
+, where L and LH
⁡
+ show α-amino acid esters and their protonated species, respectively. The protonation constants given in Table 7 are considered in more detail in order to gain more information about the effect of solvent composition and specific effects of substituents on the basicities of the esters in solvent mixtures. The numerical log
⁡
10
⁡
K values for t-butyl esters of α-amino acids determined in ethanol-water mixtures decrease with increasing ethanol content in the solvent mixture. It is observed that a nearly linear relationship exists between the aforementioned protonation constants and the mole fraction of ethanol from 0.0331 to 0.4183 for all α-amino acid esters investigated. However, log
⁡
⁡
β values at a mole fraction of ethanol of 0.4183 are slightly higher than those expected from the linear trend. The linear variation for all α-amino acid esters that are given in Table 7 is very similar to that found for ammonia aliphatic alkyl amines, pyridine, and salicylideneanilines [68, 69]. The dissociation constants of charged acids in water-ethanol mixtures vary with solvent composition in a manner that is not completely understood. It is suggested that electrostatic charging effects resulting from the change in dielectric constant with solvent effects and the solute-solvent interactions have greater significance in the interpretation of solvent effects. Thus, results obtained for α-amino acid esters in water-ethanol medium can be explained by specific solvation effects. Since ethanol would solvate L better than LH
⁡
+, the log
⁡
⁡
β values, which are related to the formation of LH
⁡
+, would decrease upon addition of ethanol. The derivations of linearity in 80% ethanol may result from the preferential solvation of solute by one of the components of the solvent mixture that could change the effective dielectric constant value in the cybotactic region [70]. Furthermore, another factor which explains the increase in the log
⁡
⁡
β values of all α-amino acid esters in ethanol-rich regions is the differences in the solvent stabilization of the ionic species (H
+ and LH
⁡
+), brought about by changing the percentage of ethanol [5, 71]. Using the protonation constants obtained in this paper, the effects of the type of ester groups on the basicity of the amine groups of α-amino acid esters have been discussed. The most important factor that affects the basicity and therefore the protonation constant of a compound is the structural effect. Table 7 showed that the basicity of t-butyl esters of glycine, L-alanine, and L-valine is higher than that of the corresponding benzyl esters of the same amino acids in 20%–60% ethanol-water mixtures. This effect can be explained by taking the electronic effect of the t-butyl and benzyl groups investigated into account.


## 3.2.1. Protonation Equilibria of Amino Acids in Mixed Water-Ethanol Medium

(1) Protonation Equilibria of L-Cysteine and L-Tyrosine in Mixed Water-Ethanol Medium (Figure1). The protonation constants of L-cysteine and L-tyrosine given in Table 3 correspond to the following equilibria:
(1)
L
2
-
+
H
+
⇌
HL
-
K
1
=
[
H
L
-
]
[
L
2
-
]
[
H
+
]
H
L
-
+
H
+
⇌
H
2
L
K
2
=
[
H
2
L
]
[
H
L
-
]
[
H
+
]
H
2
L
+
H
+
⇌
H
3
L
+
K
3
=
[
H
3
L
+
]
[
H
2
L
]
[
H
+
]Figure 1Thelog
⁡
10
⁡
K
1 values are related to the attachment of H+ to the phenolic oxygen in L-tyrosine and attachment of H
+ to the –NH2 group in L-cysteine. The log
⁡
10
⁡
K
2 value corresponds to the attachment of a proton to the –NH2 and sulfur thiol groups in L-tyrosine and L-cysteine, respectively. The log
⁡
10
⁡
K
3 values are the smallest and are thought to correspond to the protonation of carboxyl groups. The decrease in the log
⁡
10
⁡
K
2 value with increasing ethanol concentration can be explained by the fact that the dipolar ionic form is more strongly solvated than the anionic form in aqueous media.(2) Protonation Equilibria of L-Tryptophane in Mixed Water-Ethanol Medium (Figure2). Two protonation equilibria were found for L-tryptophan. The first protonation constant was attributed to the protonation of –NH2 and the second one was related to protonation of carboxyl group. No protonation constants were found related to the protonation of the nitrogen atom of the indole ring of this amino acid. The decrease in log
⁡
10
⁡
K
1 values of L-tryptophane was observed with the increase of ethanol concentration. This can be explained by better solvation of the dipolar ionic form.Figure 2(3) Protonation Equilibria of L-Lysine and Histidine in Mixed Water-Ethanol Medium (Figure3). L-lysine contains –NH2 group in the side chain. Due to the fact that this amino group is far from the carboxyl groups, resulting in a lesser degree of inductive effect, it can be said that log
⁡
10
⁡
K
1 is related to the protonation of this amino group. The second and third protonation equilibria belong to the α-amino and carboxyl groups, respectively. For L-histidine, the log
⁡
10
⁡
K
1 and log
⁡
10
⁡
K
2 values are related to the protonation of –NH2 and the nitrogen in imidazole groups, respectively. The third constant (log
⁡
10
⁡
K
3) belongs to the protonation of the carboxylate group.Figure 3Thelog
⁡
10
⁡
K
1 and log
⁡
10
⁡
K
2 values of L-lysine decrease with increasing ethanol concentration in the solvent mixture. Although, in the first equilibrium involving K
1, HL seems to be an uncharged species, it belongs to the dipolar form of amino acid. If we consider the solvation of these species, the fact that both K
1 and K
2 decrease with increasing ethanol concentration seems highly reasonable. The fact that the log
⁡
10
⁡
K
1 value of L-histidine decreases with increasing ethanol concentration can be explained by the fact that dipolar HL is better solvated in water than anionic form (L
-). The reason for the decrease in log
⁡
10
⁡
K
2 is that solvation of the multicharged species formed by the attachment of a proton to the imidazole ring is very limited in ethanol.(4) Protonation Equilibria of Someα-Amino Acids in Mixed Water-Ethanol Medium. The numerical log
⁡
10
⁡
K
1 and log
⁡
10
⁡
K
2 values of eight α-amino acids (glycine, L-alanine, L-valine, L-leucine, L-isoleucine, L-phenylalanine, L-serine, and L-methionine) determined in water-ethanol mixtures are given in Table 4. When the change of log
⁡
10
⁡
K
1 given in Table 4 with the solvent composition is examined, it is observed that these values decrease with the increase in the mole fraction of ethanol. This can be explained by the structural changes in amino acids as the medium changed toward the ethanol-rich direction. The neutral form amino acid HL is subjected to higher solvation compared with L
- forms in ethanol rich media. In water-rich media, however, the reverse behavior will be the case. This shows that log
⁡
10
⁡
K
1 values are expected to decrease with an increase in ethanol ratio. Therefore, one can conclude that the dipolar form of amino acids, HL
±, dominates in these media. It is the only way to explain the decrease in log
⁡
10
⁡
K
1 as the amount of ethanol in the media is increased since L
- is expected to solvate more than HL
± in media rich in ethanol.The fact thatlog
⁡
10
⁡
K
2 values are observed to increase as the mole fraction of ethanol is increased in all media investigated supports that dipolar ion is predominant in water-ethanol mixtures as well as in water. This is due to the fact that ethanol solvated H
2
L better than HL
±. If the neutral form were the predominant species in water-ethanol mixture, log
⁡
10
⁡
K
2 values would decrease with increasing amount of ethanol since ethanol solvated HL better than H
2
L
+ [64, 65].


## 3.2.2. Protonation Equilibria of Dipeptides in Mixed Water-Ethanol Medium

Two protonation equilibria were found for all glycine-dipeptides (Figure4) mentioned in Table 5.Figure 4The first protonation constant was attributed to the protonation of amino group and the second one was related to the protonation of carboxylate group:(2)
L
-
+
H
+
⇌
HL
K
1
=
[
HL
]
[
L
-
]
[
H
+
]
HL
+
H
+
⇌
H
2
L
+
K
2
=
[
H
2
L
+
]
[
HL
]
[
H
+
]
where L is dipeptide monoanion; HL is total zwitterion and neutral forms of dipeptide.Chattopadhyay and Lahiri [66] examined the effect of a change of solvent composition on BH
+ ionization and the related Gibbs transfer energies in mixed solvent. They emphasized that the electrostatic charge effects due to changes in the dielectric constants with changes in solvent composition are of minor importance in explaining solvent effects, and that solute-solvent interactions have greater significance in interpretation of solvent effects.These protonation constants have been considered in some detail to gain more information about the effect of solvent composition on the corresponding equilibria. For this purpose, by examining the protonation equilibria of dipeptides with solvent composition, it is generally observed thatlog
⁡
10
⁡
K
1 values decrease with the increase in the mole fraction of ethanol, whereas log
⁡
10
⁡
K
2 values increase as the composition of ethanol increases in water-ethanol mixture. The decrease in log
⁡
10
⁡
K
1 values and increase in log
⁡
10
⁡
K
2 values of all dipeptides with the increasing ethanol concentration were the same as for the corresponding free amino acids [48].The fact that the variation oflog
⁡
10
⁡
K
1 and log
⁡
10
⁡
K
2 of dipeptides with solvent composition is similar to those of the corresponding free amino acids and can be concluded from the domination of the zwitterionic form of dipeptides in water-ethanol mixtures. In water-rich media, however, the reverse behavior will be the case. This shows that log
⁡
10
⁡
K
1 values are expected to increase with an increase in the ethanol ratio. This is the only way which explains the decrease in log
⁡
10
⁡
K
1 as the amount of ethanol in the media is increased because L
- is expected to solvate more than the zwitterionic form of dipeptides in the media rich in ethanol. The fact that log
⁡
10
⁡
K
2 values are observed to increase as the mole fraction of ethanol is increased in all media investigated supports that dipolar ion is predominant in water-ethanol mixtures as well as in water. This is due to the fact that ethanol solvated H
2
L better than zwitterionic form. If the neutral form were the predominant species in water-ethanol mixture, log
⁡
10
⁡
K
2 values would decrease with increasing amount of ethanol because ethanol solvated HL better than H
2
L
+. Thus, we can say that the log
⁡
10
⁡
K
2 values for these dipeptides increase as the mole fraction of ethanol increases.In order to compare thelog
⁡
10
⁡
K
1 and log
⁡
10
⁡
K
2 dipeptides with the log
⁡
10
⁡
K
1 value of glycine and log
⁡
10
⁡
K
2 value of the other amino acid forming the dipeptide the protonation constants of the corresponding amino acids are given in Table 5. According to the data given in Table 5, the log
⁡
10
⁡
K
2 values of amino acids increase and the log
⁡
10
⁡
K
1 values of glycine decrease due to peptide formation. For instance log
⁡
10
⁡
K
1 and log
⁡
10
⁡
K
2 values for glycine in 60% water-40% ethanol medium were 9.39 and 2.89 [55, 67], respectively, while same values were found as 6.92 and 3.54 for Gly-Gly. Similarly log
⁡
10
⁡
K
1 for glycine was 9.20 [55] and log
⁡
10
⁡
K
2 for phe was 3.05 [57] in 40% water-60% ethanol medium while these values were observed to be 6.98 and 3.91 for GlyPhe in the same medium. This may be attributed to the fact that the electron density of the amino group in dipeptide is lower than that of glycine, and the electron density of the carboxylate group is higher than that of the corresponding amino acid due to the formation of a peptide bond.Estimation of equilibrium concentrations of proton-ligand formation as a function of pH provides a useful picture of proton-ligand binding in media [16]. Figures 5, 6, 7, and 8 show the change of different species of dipeptides with pH. It is seen from these figures that the most predominant species is HL (total concentration of the zwitterion and the neutral form) between pH 5 and 7 in all ethanol-water media.Figure 5
Species distribution diagram (25°C,I
=
0.1 M NaCl) for Gly-Gly system as a function of pH in water (L
=
1.5
×
10
-
3 M).Figure 6
Species distribution diagram (25°C,I
=
0.1 M NaCl) for Gly-Gly system as a function of pH in 20% ethanol-80% water (L
=
1.5
×
10
-
3 M).Figure 7
Species distribution diagram (25°C,I
=
0.1 M NaCl) for Gly-Gly system as a function of pH in 40% ethanol-60% water (L
=
1.5
×
10
-
3 M).Figure 8
Species distribution diagram (25°C,I
=
0.1 M NaCl) for Gly-Gly system as a function of pH in 60% ethanol-40% water (L
=
1.5
×
10
-
3 M).


## 3.2.3. Protonation Equilibria of Amino Acid Esters in Ethanol-Water Medium

The changes of the protonation constants of methyl and ethyl esters of L-tyrosine, L-tryptophane, L-cysteine, L-lysine, and methyl ester of L-histidine with increasing ethanol percentage are given in Table6. It can be seen that log
⁡
10
⁡
K
1 values of L-tyrosine, L-tryptophane, and L-cysteine esters increase with increasing percentage of ethanol. On the other hand, the log
⁡
10
⁡
K
2 values show a decrease with increasing ethanol concentration. The log
⁡
10
⁡
K
1 and log
⁡
10
⁡
K
2 values of L-lysine and L-histidine esters both decrease with increasing ethanol concentration in the solvent mixture.The protonation constants of methyl and ethyl esters of L-cysteine, L-tyrosine, and L-tryptophane are given in Table6. The log
⁡
10
⁡
K
1 values belong to equilibria related to the attachment of H
+ to the sulfur atom in the thiol group of cysteine esters, to the nitrogen atom in the indole ring of tryptophane, and to the phenolic oxygen group in tyrosine esters. The log
⁡
10
⁡
K
2 values are related to the protonation of –NH2 groups in the esters. The increase in the ethanol concentration decreases the dielectric constant of the medium. The amino acid ester (HL) is subjected to stronger solvation compared with the (L
-) and (H
2
L
+) forms in ethanol-rich media. In water-rich media, however, the reverse behavior will be the case. This shows that log
⁡
10
⁡
K
1 values are expected to increase and log
⁡
10
⁡
K
2 values to decrease with an increase in the ethanol ratio.Thelog
⁡
10
⁡
K
1 and log
⁡
10
⁡
K
2 values of L-lysine esters are related to the protonation of amino groups in side chain and α-amino groups, respectively. For L-histidine esters, these values correspond to the protonation of nitrogen in the imidazole and α-amino group in L-histidine. Contrary to the other esters investigated both log
⁡
10
⁡
K
1 and log
⁡
10
⁡
K
2 values of lysine methyl and ethyl esters and histidine methyl ester decrease with an increase in the ratio of ethanol in the solvent mixture. Since H
2
L
2
+ is better solvated than HL
+ and the latter is better solvated than L-molecular species in water-rich media, therefore, results obtained for these esters can be explained by the specific solvation effect.The stoichiometric protonation constants for t-butyl esters and benzyl esters ofα-amino acids determined in ethanol-water mixtures (20–80 volume % ethanol) are given in Table 7. These values are related to this equilibrium reaction L
+
H
+
⇌
L
H
+, where L and LH
⁡
+ show α-amino acid esters and their protonated species, respectively. The protonation constants given in Table 7 are considered in more detail in order to gain more information about the effect of solvent composition and specific effects of substituents on the basicities of the esters in solvent mixtures. The numerical log
⁡
10
⁡
K values for t-butyl esters of α-amino acids determined in ethanol-water mixtures decrease with increasing ethanol content in the solvent mixture. It is observed that a nearly linear relationship exists between the aforementioned protonation constants and the mole fraction of ethanol from 0.0331 to 0.4183 for all α-amino acid esters investigated. However, log
⁡
⁡
β values at a mole fraction of ethanol of 0.4183 are slightly higher than those expected from the linear trend. The linear variation for all α-amino acid esters that are given in Table 7 is very similar to that found for ammonia aliphatic alkyl amines, pyridine, and salicylideneanilines [68, 69]. The dissociation constants of charged acids in water-ethanol mixtures vary with solvent composition in a manner that is not completely understood. It is suggested that electrostatic charging effects resulting from the change in dielectric constant with solvent effects and the solute-solvent interactions have greater significance in the interpretation of solvent effects. Thus, results obtained for α-amino acid esters in water-ethanol medium can be explained by specific solvation effects. Since ethanol would solvate L better than LH
⁡
+, the log
⁡
⁡
β values, which are related to the formation of LH
⁡
+, would decrease upon addition of ethanol. The derivations of linearity in 80% ethanol may result from the preferential solvation of solute by one of the components of the solvent mixture that could change the effective dielectric constant value in the cybotactic region [70]. Furthermore, another factor which explains the increase in the log
⁡
⁡
β values of all α-amino acid esters in ethanol-rich regions is the differences in the solvent stabilization of the ionic species (H
+ and LH
⁡
+), brought about by changing the percentage of ethanol [5, 71]. Using the protonation constants obtained in this paper, the effects of the type of ester groups on the basicity of the amine groups of α-amino acid esters have been discussed. The most important factor that affects the basicity and therefore the protonation constant of a compound is the structural effect. Table 7 showed that the basicity of t-butyl esters of glycine, L-alanine, and L-valine is higher than that of the corresponding benzyl esters of the same amino acids in 20%–60% ethanol-water mixtures. This effect can be explained by taking the electronic effect of the t-butyl and benzyl groups investigated into account.


## 3.3. Protonation Equilibria in Mixed Water-Dioxane Medium

Simple heterocyclic compounds, such as 1,4-dioxane, in pesticides are very important to living organisms and environment. Also, it is well established that the “effective” or “equivalent solution” dielectric constants in protein [72, 73] or active site cavities of enzymes [74] are small compared to that in bulk water. Estimates for the dielectric constants in such locations range from 30 to 70 [73, 75]. Hence by using aqueous solutions containing ~10–50% dioxane, one may expect to simulate to some degree the situation in active site cavities [76] and hence to extrapolate the data to physiological conditions. Mixed solvents such as water-dioxane mixtures provide a better model forin vivo reactions because the mixtures simultaneously show a lesser polar character and are partially aqueous contents, as do all biological systems [2]. On the other hand, water-1,4-dioxane mixtures are a very interesting binary solvent system. These mixtures are a favorite mixed solvent system in which to study association and mobilities of ions, because the dielectric constant can be varied over a large range. Changes in protonation constants upon addition of 1,4-dioxane to aqueous solutions are due to increasing ion-ion interactions resulting from the decreasing dielectric constant and from changes in solvent-ion and solvent-solvent interactions [77–79]. Also, the protonation constant of amino acid esters will be helpful in the determination of the microscopic equilibrium constants of the corresponding amino acids. Protonation constants of some α-amino acids at different water-dioxane mixtures are given in Tables 8 and 9 [49, 67].Table 8
Protonation constants of someα-amino acids in water and water-dioxane mixtures [Temp. = 25°C and I
=
0.1 mol·dm−3 NaClO4
]
a.
AA
Water
10% D—90% W
20% D—80% W
30% D—70% W
40% D—60% W
50% D—50% W
60% D—40% W
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
log
⁡
K
1
log
⁡
K
2
Gly
9.58
2.32
9.55
2.41
9.56
2.70
9.58
2.87
9.58
3.15
9.60
3.48
9.62
3.60
Ala
9.70
2.40
9.66
2.52
9.72
2.64
9.68
2.89
9.73
3.10
9.75
3.39
9.78
3.63
Val
9.61
2.38
9.56
2.53
9.60
2.68
9.58
2.86
9.60
3.08
9.62
3.43
9.66
3.65
Leu
9.75
2.50
9.59
2.64
9.64
2.80
9.66
2.99
9.70
3.31
9.72
3.55
9.75
3.75
Ile
9.76
2.47
9.72
2.58
9.63
2.74
9.67
2.97
9.65
3.26
9.70
3.55
9.71
3.74
Phe
9.20
2.43
9.15
2.57
9.10
2.64
9.12
2.81
9.15
3.10
9.16
3.36
9.15
3.50
Ser
9.15
2.42
9.07
2.57
9.10
2.62
9.13
2.82
9.16
3.04
9.20
3.17
9.17
3.35
Thr
9.04
2.45
9.02
2.53
9.06
2.60
9.05
2.75
9.07
2.96
9.10
3.18
9.17
3.40
Note. D = dioxane, W = water, Val = DL-valine, Ala = DL-alanine, Leu = L-leucine, Ile = L-isoleucine,  Phe = DL-phenylalanine, Ser = L-serine, and Thr = L-threonine.  aData taken from reference [61]. log
⁡
K
1 corresponds to 11 species (i.e., L−  +  H+
⇌   LH); log
⁡
K
2  corresponds to 12 species (i.e.,  LH  +  H+
⇌
L
H
2
+). The % is expressed as (volume/Volume).Table 9
Protonation constants ofα-amino acid esters in water-1,4-dioxane  mixtures [Temp. = 25°C and I
=
0.1 mol·dm−3 NaCl]
a.
Esters
p
q
a
Water
20% D—80% W(X
=
0.05)
40% D—60% W(X
=
0.12)
60% D—40% W(X
=
0.24)
log
⁡
K
log
⁡
K
log
⁡
K
log
⁡
K
GlyO-Me
1
1
7.67
7.47
7.18
6.85
GlyO-t-Bu
1
1
8.38
8.21
7.91
6.97
ValO-Me
1
1
7.53
7.38
7.09
6.08
ValO-Et
1
1
8.86
8.01
7.15
6.58
ValO-t-Bu
1
1
9.05
8.54
7.71
7.13
SerO-Me
1
1
7.10
6.99
6.69
6.61
SerO-Et
1
1
7.46
7.25
6.84
6.79
LeuO-Me
1
1
7.66
7.47
7.19
6.32
LeuO-Et
1
1
7.75
7.78
7.31
6.73
LeuO-t-Bu
1
1
7.89
8.04
7.60
6.95
PheO-Me
1
1
7.11
6.66
6.42
6.01
PheO-Et
1
1
8.12
7.24
6.81
6.24
PheO-t-Bu
1
1
8.26
7.39
6.96
6.64
Note. D = dioxane, W = water; adata taken from ref. [44].  log
⁡
K
1 corresponds to 11 species (i.e., L−  +  H+
⇌    LH).


### 3.3.1. Protonation Equilibria ofα-Amino Acids in Mixed Dioxane-Water Medium

The stoichiometric protonation constants of eightα-amino acids determined in various water-dioxane mixtures are listed in Table 8. The log
⁡
10
⁡
K
1 and log
⁡
10
⁡
K
2 values given in Table 8 refer to the protonation equilibria shown in Figure 9, respectively.Figure 9Many studies have shown that the equilibrium constant is linearly related to the fraction of organic solvent [80–83]. The results obtained in this study confirm this observation. Regarding the variation of log
⁡
10
⁡
K
1 values with the solvent composition, one can postulate that the zwitterionic to neutral form ratio decreases as the dioxane content increases. This can be inferred from considerations on specific solute-solvent interactions and structural changes in amino acids from water to dioxane-water media. In the related equilibrium, the HL form would be more solvated than the L
- and the HL
± forms in solvents containing a higher concentration of dioxane, whereas the opposite forms would hold for a solvent rich in water. If the zwitterionic to neutral form ratio (K
Z) was unaltered with increasing percentage of dioxane, the log
⁡
10
⁡
K
1 values would be expected to decrease, which is not observed in the present case. If the neutral form became the predominant species when dioxane content is increased, the log
⁡
10
⁡
K
1 values would increase as in the case of a neutral monobasic acid, whereas they are not altered [84]. Thus, even with the limitations, we can conclude that the K
Z values for all amino acids studied decrease as the concentration of dioxane increases. The fact that the log
⁡
10
⁡
K
2 values increase with increasing the dioxane content for all the amino acids suggests that the zwitterion is the predominant species in dioxane-water mixtures as well as in water. This is based on the assumption that dioxane solvates H
2
L
+ better than HL
±. If the neutral form was the predominant species in water-dioxane mixtures, the log
⁡
10
⁡
K
2 values would decrease on addition of dioxane since dioxane would solvate HL better than H
2
L
+.


### 3.3.2. Protonation Equilibria of Amino Acid Esters in Mixed Water-Dioxane Medium

The stoichiometric protonation constants of theα-amino acid esters and the effect of the solvents upon the values of these protonation constants of α-amino acid esters were discussed in water-1,4-dioxane mixtures and the results are tabulated in Table 9. These values are related to the protonation of –NH2 group of the esters, as follows:
(3)
where B and BH
+ denote the α-amino acid esters and their protonated species, respectively. The protonation constants given in Table 9 were considered in more detail in order to gain more information about the effect of solvent composition and specific effects of substituents on the basicities of the esters in solvent mixtures. When it comes to the variation of log
⁡
⁡
β with the solvent composition, it can be observed that these constants decrease as the concentration of 1,4-dioxane increases in going from water to 60% in dioxane-water mixtures. It is also observed that a nearly linear relationship exists between the protonation constants and the mole fraction of 1,4-dioxane between 0.05 and 0.24 for all α-amino acid esters investigated. Many studies have shown that the equilibrium constant is linearly related to the fraction of organic solvent [80–83]. Although the effect of solvent composition upon the protonation constants of various compounds has been the subject of extensive investigation, the variation of charged acid-base equilibrium constants with solvent composition is yet to be fully elucidated. The solvation of a solute in a mixed solvent is much more complicated than solvation in a single pure solvent, and in the literature there are several different explanations. In organic solvent-water mixtures, acid-base equilibria can be estimated from two effects: an electrostatic one that can be explained by the Barron equation and a nonelectrostatic one that includes specific solute-solvent interaction [85]. Bates [71] and Chattopadhyay and Lahiri [66] have examined the effect of changing solvent composition on ionization of BH+ and the related Gibbs energies of transfer in mixed solvents, and they found that the dielectric constant alone cannot serve for a quantitative explanation for the solvent effect [66, 71]. However, the solvent mixtures are characterized by dramatic changes in the physical constants of the solvent (i.e., their melting and boiling point, dielectric constants, etc.) and macroscopic solvent parameters (mole fraction of cosolvent, etc.), upon changes of their compositions that influence many solute properties of these mixtures. Also, according to Takamuku et al. [86], the 1,4-dioxane-water mixtures can be divided into three mole fraction ranges, that is, 0.3
≤
X
dioxane
≤
0.9, X
dioxane, 0.1
≤
X
dioxane
≤
0.2, and 0
≤
X
dioxane
≤
0.07. When 0.3
≤
X
dioxane
≤
0.9, the structure observed for pure dioxane obviously remains the same as it is mixed with 0.1
≤
X
dioxane
≤
0.2, small binary clusters consisting of one or two 1,4-dioxane molecules and several water molecules are formed by hydrogen bonding. When 0
≤
X
dioxane
≤
0.07, first-, second-, and third-neighbor O
⋯
O interactions are observed due to the hydrogen-bonded network of water. This suggests that the addition of a small amount of 1,4-dioxane breaks down the hydrogen-bonded network of water. In dioxane-rich mixtures the protonated base (BH
+) is mostly solvated by the less basic dioxane molecules, whereas the proton is more solvated in the water-dioxane mixtures than in water and dioxane, because the water-dioxane mixture is more basic than either pure water or dioxane. Also, the molecular composition of the solvation layer around a solute molecule is not the same as that of the bulk mixture ratio of water to organic solvent. The hydrophobic and hydrophilic properties of a solute may be reflected in preferential solvation in such mixed solvents. Therefore, the influence that the solvent exerts on the p
K values depends upon the extent and nature of the solute-solvent interactions, which involves species participating in the acid-base equilibrium [72]. From the theoretical explanation mentioned above, we can explain results for all α-amino acid esters in terms of specific solvation effects. A study of specific solvation effects on the α-amino acid esters showed that the log
⁡
10
⁡
K values show a decrease upon increase of the 1,4-dioxane ratio for all α-amino acid esters investigated. Because 1,4-dioxane should solvate B better than BH
+, the log
⁡
10
⁡
K values, which are related to the formation of BH
+, should decrease upon addition of 1,4-dioxane. We can say that the deviations from the trend with the log
⁡
⁡
β values for all amino acid esters appeared in 1,4-dioxane-rich regions (x
=
0.24), resulting from differences in solvent stabilization of the ionic species (H
+ and BH
+) being brought about by changing the percentage of 1,4-dioxane. The results show that each methyl ester is generally a slightly weaker base than its corresponding ethyl and t-butyl analogues. In addition, the alanine methyl ester is a weaker base than its benzyl ester. The acidity or basicity of a compound in a given medium is influenced by both the electronic effect of substituents and the solvent effect of the medium. Moreover, it is sometimes extremely difficult to assess how much each effect contributes to the acidity or basicity. In the literature it has also been shown that the methyl derivatives of aliphatic amines are weaker bases than their ethyl derivatives in water-methanol mixtures [49]. We can say that the basicity of a compound is a result of various factors such as (i)
the solvent effect (solvation power, the tendency of forming hydrogen bonds, selective solvation, dielectric constant, and the composition of the solution in the first solvation layer) in the case of mixed solvents,
(ii)
structural effects, electronic effects, and steric effects.


### 3.3.3. Protonation Equilibria of DNA in Mixed Water-Dioxane Medium

The solvent effect on the acid dissociation constants of a ligand [87] can be summarized as follows.(i)
As the solvent dielectric constant decreases, thelog
⁡
10
⁡
K of the ligand increases and vice versa.
(ii)
On decreasing the extent of hydrogen bonding in water by an organic solvent, the proton-accepting properties of the water increase, and consequently thelog
⁡
10
⁡
K of the ligand decreases.
(iii)
Increasing proton solvation by an organic solvent is accompanied by a decrease in thelog
⁡
10
⁡
K of ligand.The protonation constants of inosine and inosine 5′-monophosphate (Figure10) as representative examples of DNA constituents in different compositions of dioxane-water mixtures reveal the following features (Table 10) [88]: p
K
a (N1H) of inosine increases linearly with increasing percentage of organic solvent in the medium, Figure 11. This may be correlated with the ability of a solvent of relatively low dielectric constant to increase the electrostatic forces between the proton and the ligand anion and consequently the log
⁡
10
⁡
K value increases. The same trend was observed for the log
⁡
10
⁡
K’s of (N1H) and phosphoric acid (P-OH) groups of IMP, Figure 11. The log
⁡
10
⁡
K of the phosphate group in IMP is more affected by replacement of water molecules with dioxane. This is in agreement with the above finding, namely, the existence of hydrogen bonds between water molecules and phosphate ions. Thus, the phosphate ions are easily protonated in going from water to dioxane.Table 10
Protonation constants of inosine and IMP in dioxane-water media at 25°C andI
=
0.1 mol·dm−3 NaCl.
System
% dioxane
p
q
a
log
⁡
β
b
Inosine IMPc
0%
1
1
8.55
1
1
9.09
1
2
15.07
Inosine IMP
25%
1
1
8.85
1
1
9.12
1
2
15.83
Inosine IMP
37.5%
1
1
9.04
1
1
9.24
1
2
16.28
Inosine IMP
50%
1
1
9.14
1
1
9.35
1
2
16.63
Inosine IMP
62.5%
1
1
9.36
1
1
9.39
1
2
16.72
a
p and q are stoichiometric coefficients corresponding to ligand and H+, respectively; bdata taken from reference [83], cIMP =  inosine 5′-monophosphate. The % is expressed as (volume/Volume). Species 11 refers to L− + H+
⇌   LH; species 12 refers to L + 2H+
⇌
L
H
2
+.Figure 10Figure 11
Effect of dioxane on thep
K
a of inosine and inosine 5′-monophosphate (IMP).


## 3.3.1. Protonation Equilibria ofα-Amino Acids in Mixed Dioxane-Water Medium

The stoichiometric protonation constants of eightα-amino acids determined in various water-dioxane mixtures are listed in Table 8. The log
⁡
10
⁡
K
1 and log
⁡
10
⁡
K
2 values given in Table 8 refer to the protonation equilibria shown in Figure 9, respectively.Figure 9Many studies have shown that the equilibrium constant is linearly related to the fraction of organic solvent [80–83]. The results obtained in this study confirm this observation. Regarding the variation of log
⁡
10
⁡
K
1 values with the solvent composition, one can postulate that the zwitterionic to neutral form ratio decreases as the dioxane content increases. This can be inferred from considerations on specific solute-solvent interactions and structural changes in amino acids from water to dioxane-water media. In the related equilibrium, the HL form would be more solvated than the L
- and the HL
± forms in solvents containing a higher concentration of dioxane, whereas the opposite forms would hold for a solvent rich in water. If the zwitterionic to neutral form ratio (K
Z) was unaltered with increasing percentage of dioxane, the log
⁡
10
⁡
K
1 values would be expected to decrease, which is not observed in the present case. If the neutral form became the predominant species when dioxane content is increased, the log
⁡
10
⁡
K
1 values would increase as in the case of a neutral monobasic acid, whereas they are not altered [84]. Thus, even with the limitations, we can conclude that the K
Z values for all amino acids studied decrease as the concentration of dioxane increases. The fact that the log
⁡
10
⁡
K
2 values increase with increasing the dioxane content for all the amino acids suggests that the zwitterion is the predominant species in dioxane-water mixtures as well as in water. This is based on the assumption that dioxane solvates H
2
L
+ better than HL
±. If the neutral form was the predominant species in water-dioxane mixtures, the log
⁡
10
⁡
K
2 values would decrease on addition of dioxane since dioxane would solvate HL better than H
2
L
+.


## 3.3.2. Protonation Equilibria of Amino Acid Esters in Mixed Water-Dioxane Medium

The stoichiometric protonation constants of theα-amino acid esters and the effect of the solvents upon the values of these protonation constants of α-amino acid esters were discussed in water-1,4-dioxane mixtures and the results are tabulated in Table 9. These values are related to the protonation of –NH2 group of the esters, as follows:
(3)
where B and BH
+ denote the α-amino acid esters and their protonated species, respectively. The protonation constants given in Table 9 were considered in more detail in order to gain more information about the effect of solvent composition and specific effects of substituents on the basicities of the esters in solvent mixtures. When it comes to the variation of log
⁡
⁡
β with the solvent composition, it can be observed that these constants decrease as the concentration of 1,4-dioxane increases in going from water to 60% in dioxane-water mixtures. It is also observed that a nearly linear relationship exists between the protonation constants and the mole fraction of 1,4-dioxane between 0.05 and 0.24 for all α-amino acid esters investigated. Many studies have shown that the equilibrium constant is linearly related to the fraction of organic solvent [80–83]. Although the effect of solvent composition upon the protonation constants of various compounds has been the subject of extensive investigation, the variation of charged acid-base equilibrium constants with solvent composition is yet to be fully elucidated. The solvation of a solute in a mixed solvent is much more complicated than solvation in a single pure solvent, and in the literature there are several different explanations. In organic solvent-water mixtures, acid-base equilibria can be estimated from two effects: an electrostatic one that can be explained by the Barron equation and a nonelectrostatic one that includes specific solute-solvent interaction [85]. Bates [71] and Chattopadhyay and Lahiri [66] have examined the effect of changing solvent composition on ionization of BH+ and the related Gibbs energies of transfer in mixed solvents, and they found that the dielectric constant alone cannot serve for a quantitative explanation for the solvent effect [66, 71]. However, the solvent mixtures are characterized by dramatic changes in the physical constants of the solvent (i.e., their melting and boiling point, dielectric constants, etc.) and macroscopic solvent parameters (mole fraction of cosolvent, etc.), upon changes of their compositions that influence many solute properties of these mixtures. Also, according to Takamuku et al. [86], the 1,4-dioxane-water mixtures can be divided into three mole fraction ranges, that is, 0.3
≤
X
dioxane
≤
0.9, X
dioxane, 0.1
≤
X
dioxane
≤
0.2, and 0
≤
X
dioxane
≤
0.07. When 0.3
≤
X
dioxane
≤
0.9, the structure observed for pure dioxane obviously remains the same as it is mixed with 0.1
≤
X
dioxane
≤
0.2, small binary clusters consisting of one or two 1,4-dioxane molecules and several water molecules are formed by hydrogen bonding. When 0
≤
X
dioxane
≤
0.07, first-, second-, and third-neighbor O
⋯
O interactions are observed due to the hydrogen-bonded network of water. This suggests that the addition of a small amount of 1,4-dioxane breaks down the hydrogen-bonded network of water. In dioxane-rich mixtures the protonated base (BH
+) is mostly solvated by the less basic dioxane molecules, whereas the proton is more solvated in the water-dioxane mixtures than in water and dioxane, because the water-dioxane mixture is more basic than either pure water or dioxane. Also, the molecular composition of the solvation layer around a solute molecule is not the same as that of the bulk mixture ratio of water to organic solvent. The hydrophobic and hydrophilic properties of a solute may be reflected in preferential solvation in such mixed solvents. Therefore, the influence that the solvent exerts on the p
K values depends upon the extent and nature of the solute-solvent interactions, which involves species participating in the acid-base equilibrium [72]. From the theoretical explanation mentioned above, we can explain results for all α-amino acid esters in terms of specific solvation effects. A study of specific solvation effects on the α-amino acid esters showed that the log
⁡
10
⁡
K values show a decrease upon increase of the 1,4-dioxane ratio for all α-amino acid esters investigated. Because 1,4-dioxane should solvate B better than BH
+, the log
⁡
10
⁡
K values, which are related to the formation of BH
+, should decrease upon addition of 1,4-dioxane. We can say that the deviations from the trend with the log
⁡
⁡
β values for all amino acid esters appeared in 1,4-dioxane-rich regions (x
=
0.24), resulting from differences in solvent stabilization of the ionic species (H
+ and BH
+) being brought about by changing the percentage of 1,4-dioxane. The results show that each methyl ester is generally a slightly weaker base than its corresponding ethyl and t-butyl analogues. In addition, the alanine methyl ester is a weaker base than its benzyl ester. The acidity or basicity of a compound in a given medium is influenced by both the electronic effect of substituents and the solvent effect of the medium. Moreover, it is sometimes extremely difficult to assess how much each effect contributes to the acidity or basicity. In the literature it has also been shown that the methyl derivatives of aliphatic amines are weaker bases than their ethyl derivatives in water-methanol mixtures [49]. We can say that the basicity of a compound is a result of various factors such as (i)
the solvent effect (solvation power, the tendency of forming hydrogen bonds, selective solvation, dielectric constant, and the composition of the solution in the first solvation layer) in the case of mixed solvents,
(ii)
structural effects, electronic effects, and steric effects.


## 3.3.3. Protonation Equilibria of DNA in Mixed Water-Dioxane Medium

The solvent effect on the acid dissociation constants of a ligand [87] can be summarized as follows.(i)
As the solvent dielectric constant decreases, thelog
⁡
10
⁡
K of the ligand increases and vice versa.
(ii)
On decreasing the extent of hydrogen bonding in water by an organic solvent, the proton-accepting properties of the water increase, and consequently thelog
⁡
10
⁡
K of the ligand decreases.
(iii)
Increasing proton solvation by an organic solvent is accompanied by a decrease in thelog
⁡
10
⁡
K of ligand.The protonation constants of inosine and inosine 5′-monophosphate (Figure10) as representative examples of DNA constituents in different compositions of dioxane-water mixtures reveal the following features (Table 10) [88]: p
K
a (N1H) of inosine increases linearly with increasing percentage of organic solvent in the medium, Figure 11. This may be correlated with the ability of a solvent of relatively low dielectric constant to increase the electrostatic forces between the proton and the ligand anion and consequently the log
⁡
10
⁡
K value increases. The same trend was observed for the log
⁡
10
⁡
K’s of (N1H) and phosphoric acid (P-OH) groups of IMP, Figure 11. The log
⁡
10
⁡
K of the phosphate group in IMP is more affected by replacement of water molecules with dioxane. This is in agreement with the above finding, namely, the existence of hydrogen bonds between water molecules and phosphate ions. Thus, the phosphate ions are easily protonated in going from water to dioxane.Table 10
Protonation constants of inosine and IMP in dioxane-water media at 25°C andI
=
0.1 mol·dm−3 NaCl.
System
% dioxane
p
q
a
log
⁡
β
b
Inosine IMPc
0%
1
1
8.55
1
1
9.09
1
2
15.07
Inosine IMP
25%
1
1
8.85
1
1
9.12
1
2
15.83
Inosine IMP
37.5%
1
1
9.04
1
1
9.24
1
2
16.28
Inosine IMP
50%
1
1
9.14
1
1
9.35
1
2
16.63
Inosine IMP
62.5%
1
1
9.36
1
1
9.39
1
2
16.72
a
p and q are stoichiometric coefficients corresponding to ligand and H+, respectively; bdata taken from reference [83], cIMP =  inosine 5′-monophosphate. The % is expressed as (volume/Volume). Species 11 refers to L− + H+
⇌   LH; species 12 refers to L + 2H+
⇌
L
H
2
+.Figure 10Figure 11
Effect of dioxane on thep
K
a of inosine and inosine 5′-monophosphate (IMP).


## 3.4. Protonation Equilibria in Mixed Water-DMSO Medium

The use of this mixed solvent has some advantages over pure DMSO. Thus, pure DMSO is very hygroscopic and controlling its water content is difficult [89]. This fact would affect reproducibility of experimental results. However, water-DMSO 50% : 50% mixture has only small hygroscopic character. A further advantage is its compatibility with the standard glass electrode, so that the pH measurements may be carried out in a similar way to that employed in a purely aqueous solution. In contrast, the use of pure DMSO is not recommended for potentiometry. Another advantage of the water-DMSO 50% : 50% mixture is its large acidity range (p
K
w
=
15.50) which allows the investigation of deprotonation equilibria of weak acids which could be hardly studied in water [90, 91]. Protonation constants of some α-amino acids at different water-DMSO mixtures and different temperatures are given in Tables 11 and 12 [92, 93].Table 11
Protonation constants of someα-amino acids in water and DMSO-water media at 25°C and  I
=
0.1 mol·dm−3 Na
NO
3
a.
System
p
q
b
30% DMSO
50% DMSO
70% DMSO
Gly
1
1
9.51 ± 0.01
9.79 ± 0.01
10.29 ± 0.01
1
2
12.92 ± 0.01
13.85 ± 0.02
15.20 ± 0.01
L-Ala
1
1
9.65 ± 0.009
9.85 ± 0.02
10.48 ± 0.02
1
2
13.07 ± 0.01
13.96 ± 0.03
15.36 ± 0.03
L-Val
1
1
9.34 ± 0.01
9.57 ± 0.01
9.99 ± 0.01
1
2
12.61 ± 0.02
13.61 ± 0.02
14.88 ± 0.03
L-Ser
1
1
8.95 ± 0.01
9.12 ± 0.02
9.89 ± 0.01
1
2
12.16 ± 0.02
12.95 ± 0.03
14.71 ± 0.03
L-Isl
1
1
9.31 ± 0.01
9.51 ± 0.02
9.91 ± 0.02
1
2
12.75 ± 0.01
13.56 ± 0.03
15.02 ± 0.03
L-Leu
1
1
9.46 ± 0.01
9.77 ± 0.05
10.18 ± 0.01
1
2
12.80 ± 0.02
13.81 ± 0.05
14.98 ± 0.02
L-Phe
1
1
9.06d
9.35c
—
1
2
12.44
13.43
L-Met
1
1
9.07d
9.48c
—
1
2
12.69
13.48
L-Thr
1
1
8.95d
9.32c
—
1
2
12.30
13.38
aData taken from reference [87]; b
p and q are stoichiometric coefficients corresponding to ligand and H+, respectively; cdata taken from reference [88]. Species 11 refers to L− + H+
⇌  LH; species 12 refers to L + 2H+
⇌
L
H
2
+; the % is expressed as (volume/Volume).Table 12
Protonation constants (log
10
⁡
β
q
,
r) of some α-amino acids in solution containing 50% DMSO at different temperatures and I
=
0.1 mol·dm−3 NaNO3 and their thermodynamic parametersb.
System
p
q
a
20°C
25°C
30°C
35°C
Thermodynamicsd
Δ
H
°
Δ
S
°
Δ
G
°
Gly
1
1
10.04
9.79
9.53
9.30
−85.71
−100.14
−55.87
1
2
14.14
13.85
13.52
13.22
−21.04
6.87
−23.09
L-Ala
1
1
10.02
9.85
9.69
9.46
−63.51
−24.63
−56.17
1
2
14.19
13.96
13.75
13.46
−19.34
13.80
−23.46
L-Val
1
1
9.79
9.57
9.49
9.26
−57.69
−9.63
−54.82
1
2
13.89
13.61
13.47
13.17
−21.75
4.31
−23.03
L-Ser
1
1
9.42
9.12
8.82
8.52
−103.66
−173.88
−52.02
1
2
13.30
12.95
12.60
12.25
−17.28
15.34
−21.85
L-Isl
1
1
9.72
9.51
9.39
9.20
−58.07
−12.28
−55.74
1
2
13.81
13.56
13.42
13.19
−11.06
40.55
−23.01
L-Leu
1
1
10.09
9.77
9.45
9.13
−110.57
−184.0
−55.74
1
2
14.18
13.81
13.42
13.04
−21.07
6.51
−32.01
aDMSO% (vol/vol), bdata taken from reference [87], c
q and r are stoichiometric coefficients corresponding to ligand and H+ respectively; dWith respect to the thermodynamic parameters (Δ
H
°, Δ
S
° and Δ
G
°) these parameters referred to these reactions (L−  +  H+
⇌  LH±, corresponding to –NH2 group) and (LH± + H+
⇌
L
H
2
+, corresponding to carboxylate group); The % is expressed as (volume/Volume).


### 3.4.1. Protonation Equilibria of Amino Acids in Water-DMSO Medium

The stoichiometric protonation constants of glycine, L-alanine, L-phenylalanine, L-threonine, and L-methionine in water-DMSO mixtures are given in Table11. The protonation constants K
1 and K
2 are related to the protonation of the amino nitrogen and the carboxyl oxygen, respectively. When the change of log
⁡
10
⁡
K
1 with the solvent composition, given in Table 11, is examined for α-amino acids, it is observed that these values increase with an increasing percent of dimethyl sulfoxide. The linear relation is given in Figure 12, for glycine as a representative example of α-amino acids. This can be explained by structural changes in the amino acids as the medium becomes more like dimethyl sulfoxide. The neutral form of the amino acids (HL) is subjected to a larger amount of solvation, compared to the anionic form (L
-) in dimethyl sulfoxide rich media. In water-rich media, however, the reverse will be the case.Figure 12
The variation of protonation constants of glycine with the percentage of DMSO.Also, thelog
⁡
10
⁡
K
2 values increase for each α-amino acid as the percentage of dimethyl sulfoxide in the solvent mixtures increases. Many studies have shown that the equilibrium constant is linearly related to the fraction of the organic solvent [80, 81, 83]. In general, the protonation of compounds containing O–H increases with increasing organic content of the solvent, due to the decrease in the dielectric constant of the bulk solvents. As the dielectric constant decreases, the ion-ion interaction involving the proton and the anionic oxygen donor of the ligand increases to a greater extent than the ion-dipole interaction between the proton and the solvent molecule. Thus the log
⁡
10
⁡
K of carboxylic acid (COO−H+) groups of α-amino acids increases linearly with increasing percentage of organic solvent in the medium. The same trend was observed for the p
K
a’s of phosphoric acid (P-OH) groups of IMP [88].Serine isα-amino acid containing three functional groups (–OH, –NH2, and COOH), but only two protonation equilibria are determined corresponding to the protonation of amino and carboxylate groups, respectively. This can be attributed to the fact that, as the relative pH in most experiments did not extend to pH > 12, the protonation constant for the aliphatic OH group in serine could not be measured. Also, a qualitative explanation could be based on the formation of an intramolecular hydrogen bond as depicted in Figure 13. Assume that the proton on this –OH group is more difficult to remove in the amino acid form.Figure 13
Intramolecular H-bond of serine.Also, thelog
⁡
10
⁡
K
2 values increase for each α-amino acid as the percentage of dimethyl sulfoxide in the solvent mixtures increases. Many studies have shown that the equilibrium constant is linearly related to the fraction of the organic solvent [80, 81, 83].


## 3.4.1. Protonation Equilibria of Amino Acids in Water-DMSO Medium

The stoichiometric protonation constants of glycine, L-alanine, L-phenylalanine, L-threonine, and L-methionine in water-DMSO mixtures are given in Table11. The protonation constants K
1 and K
2 are related to the protonation of the amino nitrogen and the carboxyl oxygen, respectively. When the change of log
⁡
10
⁡
K
1 with the solvent composition, given in Table 11, is examined for α-amino acids, it is observed that these values increase with an increasing percent of dimethyl sulfoxide. The linear relation is given in Figure 12, for glycine as a representative example of α-amino acids. This can be explained by structural changes in the amino acids as the medium becomes more like dimethyl sulfoxide. The neutral form of the amino acids (HL) is subjected to a larger amount of solvation, compared to the anionic form (L
-) in dimethyl sulfoxide rich media. In water-rich media, however, the reverse will be the case.Figure 12
The variation of protonation constants of glycine with the percentage of DMSO.Also, thelog
⁡
10
⁡
K
2 values increase for each α-amino acid as the percentage of dimethyl sulfoxide in the solvent mixtures increases. Many studies have shown that the equilibrium constant is linearly related to the fraction of the organic solvent [80, 81, 83]. In general, the protonation of compounds containing O–H increases with increasing organic content of the solvent, due to the decrease in the dielectric constant of the bulk solvents. As the dielectric constant decreases, the ion-ion interaction involving the proton and the anionic oxygen donor of the ligand increases to a greater extent than the ion-dipole interaction between the proton and the solvent molecule. Thus the log
⁡
10
⁡
K of carboxylic acid (COO−H+) groups of α-amino acids increases linearly with increasing percentage of organic solvent in the medium. The same trend was observed for the p
K
a’s of phosphoric acid (P-OH) groups of IMP [88].Serine isα-amino acid containing three functional groups (–OH, –NH2, and COOH), but only two protonation equilibria are determined corresponding to the protonation of amino and carboxylate groups, respectively. This can be attributed to the fact that, as the relative pH in most experiments did not extend to pH > 12, the protonation constant for the aliphatic OH group in serine could not be measured. Also, a qualitative explanation could be based on the formation of an intramolecular hydrogen bond as depicted in Figure 13. Assume that the proton on this –OH group is more difficult to remove in the amino acid form.Figure 13
Intramolecular H-bond of serine.Also, thelog
⁡
10
⁡
K
2 values increase for each α-amino acid as the percentage of dimethyl sulfoxide in the solvent mixtures increases. Many studies have shown that the equilibrium constant is linearly related to the fraction of the organic solvent [80, 81, 83].


## 3.5. Thermodynamic Quantities Associated with the Protonation Equilibria

Changes in the temperature may have some significant influences on the equilibrium properties and cause spontaneous changes in aquatic systems. Solution equilibria may be shifted to the right or to the left by an increase in temperature. The direction and magnitude of the shift are given by the sign and magnitude of the heat of the reaction as a function of temperature.The thermodynamic parameters are useful tools for studying these interactions with bioligands and understanding the relative stability of the formed complexes.The enthalpy changes accompanying complex formation can be obtained directly through calorimetric measurements, or indirectly from the temperature dependence of the stability constants via the Van’t Hoff equation. The accuracy ofΔ
H values based on calorimetric measurements is generally higher than the accuracy of those based on the temperature dependence of log
⁡
10
⁡
K values. However, many parameters influence the accuracy of calorimetric data, for example, the enthalpy of dissociation of water, the heat of dilution, and so forth [93].Application of the temperature dependence oflog
⁡
10
⁡
K generally requires the determination of very accurate stability constants over as wide range in temperature as possible. However, the wider the range employed, the greater the uncertainty in the calculated values, because of the temperature dependence of the enthalpy values.A high electrolyte concentration was used in order to keep variations of the activity coefficients at a minimum. Precise thermodynamic data can only be obtained provided that an inert electrolyte of fairly high concentration (≥0.1 mol·dm−3) is used. The main differences between different procedures for the study of ionic equilibria in aqueous and nonaqueous solutions are due to the activity coefficients. As in most equilibria in nonaqueous medium, a background electrolyte is added to maintain constant ionic strength ranging from about 0.1 to 0.3 mol·dm−3. This is allowed in some water-ethanol or water-dioxane mixtures, but not in solvents of low dielectric constants where the solubility of electrolyte is very low. In rigorous thermodynamic calculations, the equilibrium constant should be expressed in terms of activities of the component ions at equilibrium. However, for convenience, concentrations are generally used, where, for a particular species i, a
i
=
(
C
i
)
F
i, where C
i is the concentration of the ion i, a
i is its activity, and F
i is its activity coefficient. The activity coefficients are maintained constant by working in a medium of constant ionic strength.The results are summarized in Tables12 and 13 and can be interpreted as follows.(1)
The protonation reactions (L
-
+
H
+
⇌
L
H
±, corresponding to –NH2 group) and (L
H
±
+
H
+
⇌
L
H
2
+, corresponding to carboxylate group) in Table 12 of the investigated amino acids are exothermic (Figure 14). Three factors affect these protonation reactions:
(i)
the neutralization reaction, which is an exothermic process,
(ii)
desolvation of ions, which is an endothermic process,
(iii)
the change in the configuration and the arrangement of the hydrogen bonds around the free and protonated ligands.
(2)
Thelog
⁡
10
⁡
K values decrease with increasing temperature revealing that the acidity increases with increasing temperature (Tables 12 and 13).Table 13
Effect of temperature and thermodynamic parameters of protonation equilibria of inosine and IMPa.
System
T (°C)
p
q
b
log
⁡
β
Equilibrium
Δ
H
°  (kJ mol−1)
Δ
S
°   (JK−1 mol−1)
Δ
G
°  (kJ mol−1)
Inosine
15
1
1
8.77
Inosine
−35.0
46.3
−48.8
IMP
15
1
1
9.33
(1) L−  +  H+
⇌  LH
1
2
15.28
Inosine
20
1
1
8.67
IMP
20
1
1
9.17
1
2
15.17
Inosine
25
1
1
8.55
IMP
−33.0
64.6
−52.2
IMP
25
1
1
9.09
(2) L3−+  H+
⇌   LH2−
0.32
114.6
−33.8
1
2
15.07
(3) LH2− + H+
⇌
L
H
2
-
Inosine
30
1
1
8.44
IMP
30
1
1
8.95
1
2
14.99
a
p and q are stoichiometric coefficients corresponding to ligand and H+, respectively; bdata taken from reference [83]; L denotes inosine and inosine 5′-monophosphate.Figure 14
Van’t Hoff plot oflog
⁡
⁡
K of glycine in 50% DMSO-50% H2O mixture (log
⁡
⁡
K
1 denotes protonation of amino group and log
⁡
⁡
K
2 denotes protonation of carboxylate group).The protonation reactions (1) and (2) in Table13 [88] of the N1 site of inosine and inosine 5′-monophosphate are exothermic.The phosphate ion in IMP is more solvated than its protonated form. Consequently it contributes more to the endothermic process upon protonation and the netΔ
H
° is only 0.32 kJ mol−1. Also, the phosphate ions form ordered hydrogen bonds with water molecules, which is confirmed by the large entropy change Δ
S
°
=
114.6 JK−1 mol−1. This contributes to a negative free energy change Δ
G
°
=
33.8 kJ mol−1. The positive entropies due to the release of ordered water molecules and the breaking of hydrogen bonds were found by Kramer-Schnabel and Linder [73] to give positive Δ
H
°, large Δ
S
°, and net negative Δ
G
° for the protonation and complexation reactions of organic monophosphates and copper ions. The log
⁡
10
⁡
K values decrease with increasing temperature revealing that its acidity increases with increasing temperature.


## 4. Conclusions

Great efforts should be made to collect data with as high quality as possible. This is often a time-consuming tough job but there are no short cuts in this respect. We would like to stress the importance of the following points.(a)
The amino acids have special importance among the other chemical molecules because they are the foundation stones of living organisms. The determination of the protonation constants is of great importance for the elucidation of numerous biological compounds.
(b)
The knowledge of the protonation constants for these bioligands is a prerequisite to an understanding of their mechanism of action in both chemical and biological processes. Thus, it is hoped that these data will be a significant contribution to workers carrying out mechanistic studies in biological media. However, a fundamental understanding of the effects that structural variations have on the properties of the amino acid is essential in order to interpret and predict the properties of proteins correctly and also in order to explain the behaviors of enzymes in living organisms.
(c)
In addition, the protonation constants of various amino acids investigated in this study not only are used for quantitative purposes but also can be used for the evaluation of the solvent effect and structural behavior of the amino acids in water containing media.
(d)
Also, the present study supports the idea that the solvent effect on the protonation constants for the amino acids can be utilized to predict whether the zwitterion or the neutral form is the predominant species and how the zwitterionic to neutral form ratio changes with the concentration of organic component in water-organic solvent mixtures.
(e)
The protonation constants of some dipeptides were determined in water and water-ethanol mixtures. Generally, it was observed that thelog
⁡
10
⁡
K
1 values of dipeptides decrease and the log
⁡
10
⁡
K
2 values increase with the increase in the mole fraction of ethanol. The chemical and biological activity of these substances would be expected to vary with the degree of ionization. Therefore, a good knowledge of the protonation constants of dipeptides in nonaqueous media is of considerable interest.
(f)
Additionally, the protonation constant of amino acid esters will be helpful in the determination of the microscopic equilibrium constants of the corresponding amino acids due to knowledge of the microscopic equilibrium constants of some compounds being necessary for the calculation of the concentration of each ionized species at any pH, which is important for the complete understanding of the physicochemical behavior of such molecules.
(g)
The basicity of a compound is a result of various factors such as (i) the solvent effect (solvation power, the tendency of forming hydrogen bonds, selective solvation, dielectric constant, and the composition of the solution in the first solvation layer) in the case of mixed solvents and (ii) structural effects, electronic effects, and steric effects.
(h)
Equilibrium constants are directly related to the solvent composition.
(i)
The search for the equilibrium model must be unambiguous. The search for the speciation scheme as a service to the reader should be based on computerized methods and visualized in the form of diagrams.
(j)
The discussed data in this study will make an important contribution to the literature.



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*Source: 626719-2014-08-14.xml*