1886
P. L.
The Journal of Physical Chemistty, Vol. 82, No. 17, 1978
Huyskens and Y. 0. Lambeau
Specific Interactions of Citric Acid with Anions in Acetonitrile P. L. Huyskens" and Y. 0. Lambeau Depatfment of Chemistry, University of Leuven, 200-F, Celestljnenlaan, B 3030 Heverlee, Belgium (Received February 14, 1978) Publication casts assisted by N. V. Citrique Belge
From conductance measurements the dissociation constants Kd of several triethylammonium salts were determined at 288.16,298.16, and 308.15 K in acetonitrile in presence of concentrations of citric acid varying from 0 to 0.08 mol dm-3. From the variation of Kd with respect to L , it can be inferred that C1- and HSO, can bind two molecules of citric acid and that the undissociated ion pairs are still able to become complexed by one molecule ligand. The corresponding equilibrium constants kc, k f , and Kl were determined and are weaker for HSO,. For Br-, NO,, and I- the first addition constant is still smaller and k , becomes too small to be detected. Although it is possible for the anions forming two hydrogen bonds to chelate with the same ligand molecule, it was shown that the majority of 1:l complexes of these anions with citric acid remain unchelated. The enthalpy of the first bond although negative is rather small and the standard entropy of formation of this bond is positive. This paradox can be explained by the formation of a specific bond between citric acid and acetonitrile. Owing to the lower symmetry of the solvent molecule compared with the anions, and to the higher number of active sites on the latter, the liberation of the solvent molecule when the anion takes its place is accompanied by an increase in disorder and entropy which is greater than the reverse effect which accompanies the fixation of the ion. The interaction of acetonitrile with citric acid is confirmed by the low value of the vOH stretching vibration frequency in this solvent. The formation of the first bond triggers a very important reduction, described by the ratio kZ-/ki, of the electron donor power of the second site on the anion. For various anions and acids In .Izz/kl-varies approximately linearly with In kl- in acetonitrile. However the function is different in CC4. The difference can also be explained by the existence of a hydrogen bond between the acids and acetonitrile. Most anions, such as halogenide ions, NO3-, or SO4-, possess external electron pairs which enable them to form hydrogen bonds with acid molecules present in solution. It is possible to determine the stability constants kl-, k2-,... of these bonds from conductivity measurements. This was done for a great variety of anions and acids by Kolthoff, Chantooni, and their co-w0rkers.l In a previous work2 we determined in this way the stability constants of the complexes between C1-, Br-, and NO; and a series of substituted benzoic acids in nitrobenzene and acetonitrile. In some cases the existence of twice complexed ions, for instance C6H6COOH. -Br-. .HOOCC6H5 was shown. A hydrogen bond can also be formed between the ligand and an ion which is still bound to the counterion. This was shown for instance by Rulinda and ZeegersH ~ y s k e n s using ,~ infrared spectrometry, for iodide or bromide ions, bound to tetraalkylammonium ions, which can be complexed by one or two phenol molecules: CGH5O-H. *I--R4N+ In this work we investigate the complexation of halogenide ions, HSO, and NO,, by citric acid. This acid was chosen because, in addition to the important role it plays in the metabolism of carbon hydrates, lipids, and proteins in living organisms, it is particularly rich in proton donor 0-H groups. In the case of citric acid the question also arises if a given anion can be bound by several 0-H groups of the same ligand molecule, giving chelates such as
-
*;XH
oi COOH
I
'C
TH2
$0
0022-3654/78/2082-1886$01.00/0
As a matter of fact in citric acid geometrical conformation like these occur which bring the 0-H groups in positions suitable for such interactions. In this work we try, among other things, to determine whether the formation of such chelates by citric acid is significant or not. The solvent is acetonitrile and the temperature 298.16 K. In the case of HS04- and Brmeasurements were also performed a t 308.16 and 288.16 K in order to obtain information on the enthalpies of the hydrogen bonds. Principle of the Method The overall dissociation constant Kd of the ionophore, the formal concentration of which is F, is defined as (1) Kd = [iI2y2/(F- [i]) where [i] is the concentration of the positive ions which is equal to the sum of the concentrations of the anions in the various states of solvation. y is the activity coefficient which in this range of [i] can be computed by means of the Debye-Huckel equation. Kd is influenced by the presence of the ligand L. If the permittivity of the medium remains constant, the ratio Kd/Kdorespectively in presence and in the absence of the ligand is in this case given by the relation
L is the concentration of the free ligand, and kl-, kz-, and K1 are the addition constants of the ligand respectively on the first and second sites of the anion and on the ion pair. This equation is a generalization of the sort used by Gilker~on.~ In this case however the presence of the ligand a t concentrations as high as 10-1 mol dm-3 can modify the dielectric constant. It is therefore necessary when using eq 2 to correct Kd for this change in the permittivity. In order to make this correction, the variation of log Kdo of the ionophores in mixed solvents acetonitrile-benzene of 0 1978 American Chemical Society
Interactions of Citric Acid with Anions in Acetonitrile
The Journal of Physical Chemistry, Vol. 82, No. 17, 1978 1007
TABLE I: Dissociation Constant Kd of Et,NH+Br in Acetonitrile at 298 K Computed from the Conductivity Data Using Eq 5 F/10'3 mol dm-3
[i]/10-3 mol dm-3
Y
Kd / 1 v4 mol dm"
0.614 0.923 1.151 1.671 2.09 2 2.523 2.98 7 3.507
0.311 0.409 0.470 0.595 0.686 0.770 0.851 0.936
0.9 36 0.927 0.922 0.912 0.906 0.901 0.896 0.891
2.80 2.79 2.75 2.73 2.75 2.74 2.72 2.71
Mean value K d = (2.75 f 0.03) X
remain fairly constant. The light systematic decrease may be due to inaccuracy in the values of the activity coefficients or to the fact that the higher terms of FUOSS' conductance equations were ignored.
mol dm-3
decreasing permittivity was determined. On the other hand the presence of citric acid also modifies the viscosity 7 of the solutions and this has an influence on the mobility and the ionic conductances of the ions. This effect was taken into account assuming Walden's rule holds for such changes in viscosity. It was further assumed that the effects of the other ions' presence on the mobility, in this range of [i], are sufficiently taken into account by the use of Onsager limiting equations5 Under these circumstances the overall ion concentration [i] is related to the conductivity K of the solution by the equation (K - K o ) ~ x 1000 [i] = (3) (w+ + w-) - [a'(w+ + w-) + 2p'][i]'/2
w+ and w- are the Walden products of the cation and of the anions. a' and p' are known coefficients which depend on the dielectric constant and on the temperature, and which appear in the Onsager equation. K~ is the residual conductivity in the absence of the ionophore. The Walden product w+ of the triethylammonium ion is constant but that of the anions w- depends on the extent of solvation of these ions and thus on L. It was experimentally determined for given values of L from the conductivity of solutions of tetraalkylammonium salts which are practically completely dissociateda6 Following a procedure proposed by Fuoss7 it is possible to solve eq 3, introducing the variable z defined by 1ooo1/2[ol'(W++ W-) + 2p](K - Ko)'/~~'/* z= (4) (w+ + w-)3/2 and calculating the function f(z) described and tabulated by this author. Under these circumstances eq 1becomes ((1ooo(K - K O ) ~ )? Y2 w+ w-)f(2)
+
Kd=
1ooo(K - K o ) ~ w+ w-)f(z)
(5)
+
Kd can thus be computed for several concentrations of the ionophore, leaving the concentration of the ligand L constant. An example is given in Table I which refers to Et3NHBr in acetonitrile in the absence of ligand. The values of Kd
Experimental Section Conductivities were measured by Wayne Kerr universal bridge B221 and B224 a.c. conductometers operating a t a frequency of 1592 Hz. Philips PW 95/2/01 cells were used. The cell constants were determined according to the procedure proposed by Lind, Zwonelik, and Fuossag Viscosities were measured using the automatic MGM Lauda viscosimeter. The densities of the solutions were determined with a hydrostatic balance. Their dielectric constants were measured by means of a WTW dekameter operating at 2 MHz. The infrared spectra were determined by means of a Perkin-Elmer 180 spectrophotometer. Products. Et3NH+C1-, Fluka purum, was purified by recrystallization from a benzene-methanol solution. Et3NH+Br-, Eastman Kodak, was purified by recrystallization from a butanol-methanol solution. Et3NH+I- was prepared from H I and Et3N. It was crystallized from a benzene-acetone solution. Et3NH+HS04- and Et3NH+N03-were prepared from Et3NH+Br--methanol solutions by ion exchange on a Baker resin, Et3NH+HS04was further purified by recrystallization from methanol solutions. Et3NH+N03-was recrystallized from methanol-benzene solutions. These products are highly hydroscopic. Citric acid (Citrique Belge) was used without further purification. The solvent acetonitrile (Fluka puriss.) was purified according to the method used by Coetzee and his co-workers.1° The residual conductivity does not exceed ohm-' cm-'. Results For each ionophore and for each concentration L of the ligand, the conductivities of a series of ten solutions with a formal ionophore concentration ranging from 0.2 X to 4 X mol dm-3 were measured. The conductivities lie between 0.1 X and 5 X ohm-' cm-l. For each ionophore the conductivities of eight solutions containing respectively 10,20, and 30 wt ?& benzene, with F ranging from 0.1 X to 5 X mol dm-3, were measured in order to determine the influence of the dielectric constant on the dissociation constant. The conductivities here lie between 0.4 X and 4 X ohm-l cm-l. All these direct experimental data are available upon request. For each ionophore and for each concentration L of the ligand, the dielectric constant D and the viscosity 7 were measured (these values are practically not affected by the variations of F in our experiments). It is then possible using the appropriate values of w+ and w-to compute ten separate values of Kd from eq 5. The mean value of Kd and the values of D , 7, w+, and w- for a given ionophore at a given ligand concentration and at a given temperature are available as supplementary material (see paragraph at end of text regarding supplementary material). From these data we summarize in Tables I1 and I11 those which are needed for the discussion.
TABLE 11: Dissociation Constant (Kd0/10-5mol d m - 3 )of the Ionophores in Media of Various Constants, D , in the Absence of Ligands ( T =298.16 K) ionophore Et,NH+ClEt,NH+BrEt,NH+NO,Et,NH+HSO,Et,NH+I-
D = 36.10 (pure acetonitrile) 3.05 27.0 79.3 81.1 316
D = 32.10 (10%benzene)
D = 28.45 (20%benzene)
D = 24.75 (30%benzene)
1.95 14.7 45.6 47.4 206
0.97 7.5 24.1 26.3 108
0.43 3.2 11.1 13.7 51
1000
P. L. Huyskens and Y. 0. Lambeau
The Journal of Physical Chemistry, Vol. 82, No. 17, 1978
TABLE 111: Dissociation Constants mol dm-l) of Triethylammonium Salt in Acetonitrile at Various Concentrations (L/mol dm-l) of Citric Acida
T,K
D
298.16 288.16 298.16 308.16 298.16 288.16 298.16 308.16 298.16
36.10 37.71 36.10 34.61 36.10 37.71 36.10 34.61 36.10
anion
c1-
L = 0.00
0.01
3.1 25.0 27.0 27.6 79.3 78 81 83 316
26.8
0.02
0.03
49.4 67.9 60.0 77.1 61.7 43.2 78.5 63.5 79.6 132 179 233 307 167 237 163 230 300 HS0,158 226 284 I334 352 370 a The dissociation constants are corrected for the change in permittivity resulting from BrBrBrNO,: HS0,HS0,-
0.04
0.06
0.08
86.3 91.9 94.2 94.6 28 5 373 354 343 388
118.6 120 125.2 126.8 378 487 46 3 440 433
150.3
570 541 468 the addition of the ligand.
TABLE IV: Standard Free Energy, Enthalpy, and Entropy of Hydrogen Bond Formation between Et,"' the Anions in Acetonitrile at 298.16 K (D= 36.1)a AHh',
anion
AGh", kJ mol-'
kJ
mol-I
Ash", J
K-'
A H " * , kJ
mol-'
A G h o * , kJ mol-'
and
mol-'
- 2 5 . 8 * 0.1 -14.9 * 1.6 -4i. 2 55 f 4 - 8 . 8 c 1.0 -7 f 2 - 2 0 . 4 ?: 0.1 -17.7 * 0.1 - 7 . 3 f 1.0 -17.7 * 0.1 -2* 1 53 f 4 - 8 . 2 f 1.2 -5* 2 -14.3 * 0.1 -4.2 * 0.8 Extrapolated values AGh" *, AHh'*, and AS"* in a medium of infinite dielectric constant.
5f 9
BrNO,HS0,-
Discussion A. Influence of the Dielectric Constant on the Dissociation of the Ionophores. When log Kdo is plotted against 1 / D (Figure 1)the functions are linear. The slopes d log Kdo/d(l/o) are respectively -68 for C1-, -73 for Br-, -67 for NO3-, -61 for HSO;, and -62 for I-. These values were used to correct the experimental values of Kd in the presence of the ligand for the change in the dielectric constant of the medium. According to Denison and Ramseyll these slopes can be related to the distance of closest approach (a) between the charges by the relation
3.6 A).
B. Hydrogen Bonds between Et3"+ and the Anions. In Table I1 important differences appear between the dissociation constants of the various triethylammonium salts. These differences cannot be explained only by the change in the size of the ions. Of course, a modification of the parameter a in the equation of Denison and Ramsey
10 i 10
t 0 02
0 01
The a values which are computed from the slopes (3.5 A for C1-, 3.3 for Br-, 3.6 for NO3-, 4 for HS04-, and 3.9 for I-) are effectively on the order of magnitude of the distances between the centers of the ions in the ammonium salts in the solid state (respectively, 3.1,3.2, 3.4, 3.8, and
K-'
Asho*, J
mol - I
c1-
a
163.0 159.4 481
0 03
004
1/D
-
Figure 1. Logarithm of the dissociationconstant K: of triethylammonium salts at 298.16 K in nitrobenzene-benzene mixtures as a function of the reverse of the dielectric constant.
bonds can be computed from log Kdoand its variation with 1/T. This can also be done for the thermodynamic characteristics AGho*, mho*, and Asho* which would be found in a medium of infinite dielectric constant where the electrostatic attraction between the ions of opposite charge is eliminated. From eq 7, one finds
(7)
will bring about changes in the constant Kdo. However, as can be seen from Figure 1,strong differences also appear in the extrapolated values log KO* a t infinite dielectric constant for the various ions. The differences in Kd0 are thus not provoked by differences in the electrostatic attraction between the ions resulting from changes in their sizes, but must be mainly ascribed to the changes in the characteristics of the hydrogen bonds between the ions. Evidence of hydrogen bonding can also be found in the lowering of the VN-H stretching frequencies of the triethylammonium salts in the sequences of the anions C104- > I- > Br- > Cl-.12J3 The standard free energy AGho, the standard enthalpy mho, and the standard entropy Asho of formation of these
For acetonitrile d log D l d log T = -1.297. In Table IV these thermodynamic characteristics of the reaction EtSNH+ X- Et3NH'. * *X-
+
+
are tabulated. In acetonitrile the stability of the hydrogen bonds between the anions and Et3"+ decreases following the sequence C1- > Br- > HS04- > NO3- > I-. These bonds are characterized by very weak enthalpies. This is partly due to the fact that the exothermic polarization of the solvent molecules around the free ions is
The Journal of Physical Chemistry, Vol. 82, No. 17, 1978 1889
Interactions of Citric Acid with Anions in Acetonitrile
TABLE V: Equilibrium Constants k,-,k2-,and K, for the Addition of a First and a Second Citric Acid Molecule on the Anions and of Citric Acid Molecule on the Triethylammonium Ion Pairs'"
ki-,dm3
T,K
anion
mol
k2-,dm3 mol - *
K, , dm3
AH,,kJ
AS,, J K-'
mol-'
mol-'
mol-'
298.16 c18 8 5 + 25 lo+5 20 c 6 288.16 Br78 c 6 3 + 1.5 298.16 Br72i 6 -6i: 6 + 1 5 i 20 308.16 Br66 i 6 298.16 NO,64c 2 119 i: 8 4 i 2 9 + 3 288.16 HS0,298.16 HS0,108 i 8 3i 2 8i 3 - 8 k 6 + 1 4 i 20 308.16 97i 8 2 i 2 7c3 HS0,298.16 I6c 1 a Standard molar enthalpy A Hlo and standard molar entropy A s l o of the first bond. The solvent was acetonitrile.
more important than around the ion pair. As a consequence m h o , which is the sum of all the changes in enthalpy brought about by the formation of the bond, is rather small. The orientation of the polarized solvent molecules around the free ions is drastically reduced around the ion pair. This is the main reason for the high positive value of the entropy A s h o . However, even when the share due to the polarization of the solvent molecules is substracted from AGh", M h o , ASh" (what is done in effect when considering AGho*, mho*, and Asho*), the enthalpies mho* still remain very small and the entropies Asho* are positive. Such behavior can be explained if it is assumed that the cation Et3"+ forms a specific bond with the solvent molecules. In this case, in effect, the reaction must be written Et3N+H**.N=CCHS
+ X-
-+
+
Et3N+H*.*X- CH~CEN
In such a reaction the hydrogen bond between the ions replaces the hydrogen bond between the cation and the solvent. Furthermore, the anion X- is much more symmetric than the solvent molecule CH3CH. The limitations in the orientations of X- when it is bound by Et3"+ may thus be less important than those of the solvent molecule when it forms a hydrogen bond with the cation. This can explain why the value of Asho* remains positive. The existence of a specific bond between Et3"+ and acetonitrile is confirmed by the fact that in nitrobenzene the dissociation constants of the triethylammonium salts are systematically one order of magnitude smaller than in acetonitrile. The values of AGho* of Et3NH+Cl- and Et3NH+Br- in nitrobenzene, computed from our previous results,2 are respectively -34.4 and -30.2 k J mol-I which are 8.6 and 9.8 k J mol-I lower than in acetonitrile. In the case of triethylammonium picrate, the later difference is 9.4 k J mol-l. From these data it can be concluded that the formation of the specific bond between triethylammonium and acetonitrile lowers the standard free energy by some 9.3 kJ mol-l. C. Complexation of the Anions and of the Ion Pairs by Citric Acid. The data of Table I11 clearly demonstrate the positive influence of the presence of citric acid on the dissociation constant of all the ionophores. When the ratio R 3 Kd/Kdo is plotted against the concentration L of the ligand, downward curved curves are obtained for C1-, Br-, and HS04-. An example is given in Figure 2. For NO3and I- the functions remain linear within the accuracy of the determination. This demonstrates that for the three first ions complexation of the ion pairs by citric acid occurs to a marked extent. Furthermore the constant Kl must be greater than the addition constant k2 of the second ligand molecule on the once solvated anion. When the function ( R - 1)/L is
F
Kd/ K:
7 / 298 16 K 308 16 K
m o l e dr6'
0
2
4
6
8
Figure 2. Ratio K d I K 2 for the dissociation of Et3N+H-HSO4- in acetonitrile vs. the concentration L of citric acid.
plotted against R, significant deviations from linearity occur for Cl- and HS04-. In the case of Br- the deviations remain within the limits of experimental errors on R. This shows that C1- and HSOc can become solvated twice by the citric acid molecules. In Table V we give the values of hl-, hz-, and K1 which best fit relation 2 for the experimental values of R vs. L. The maximum absorption for the AVOHstretching band of citric acid in acetonitrile is located at 3240 cm-l. This frequency lies already fairly lower than the one characterizing a carboxylic group of similar acid strength in more inert solvents, where the band appears in the vicinity of 3550 cm-l. When C1- is added to the solution, a second band appears at 2770 cm-l, which can be ascribed to the 0-H- stretching vibration in the complex. This represents thus a shift AVOHof the order of 780 cm-l, relative to the "free" 0-H group. It was not possible to determine the position of the band of the complex for the other ions because of overlapping of the bands of the solvent. As for the interaction with Et3N-H+ the -AHl values are rather small in acetonitrile. This is also the case for the interaction between C1- and monochloroacetic acid for which Lam et al.14 determined a AH1value of -9 k J mol-l in the same solvent. The interactions between the anions and citric acid are also characterized by positive (and a t any rate low) values of the standard entropies of reaction ASo. Thus, in acetonitrile, the reaction Br- + Cit-H Cit-H. * .Bris accompanied by an increase in disorder! This behavior can be explained by the existence of a specific interaction between citric acid and acetonitrile. The reaction must then be written Br- Cit-H*..N=CCH3 Cit-Ha ..Br- N=CCH3
-
+
-
+
As a consequence of the larger number of specific sites of
1890
P. L. Huyskens and Y. 0. Lambeau
The Journal of Physical Chemistry, Vol. 82, No. 17, 1978
11 a
10
-
14
-
12
-
10
-
8 -
" 2
0
0
ri
0
1
2
3
L
5
6
d
I
Flgure 3. Standard free energy -4Glo/kJ mol-' for the addition of a first molecule of benzoic or carboxylic acids on CI- (0)and (A) against the pK, of the acid in water; (W) citric acid (references from Table VI).
the Br- ion and of its higher symmetry compared with CH3C=N, the formation of a specific bond by the latter involves a more drastic reduction of its orientation possibilities. Thus the increase in order as a result of the bromide ion's fixation and of the stronger bonding of citric acid is overcompensated by a decrease in order accompanying the liberation of the solvent molecule previously bonded to citric acid. This explanation is corroborated by the frequency shift of 310 cm-' which the vOH bond of citric acid in acetonitrile exhibits compared to the situation in inert solvents. Furthermore, it was shown in a previous work2 that passing from acetonitrile to nitrobenzene strongly enhances the kl-values of the anions complexed by benzoic acids. In Table VI we compare the standard free energy AGlo of formation of the 1:l complex between the anions and citric acid with values concerning similar complexes in the literature. For C1-, as shown in Figure 3, the following relation holds approximately AGlo = 15.6 kJ mol-l (single)
. i
0
+ 2.35(3 - pKa) kJ mol-l
Most of the points obey this relation within 0.6 kJ mol-'. (A strong deviation is observed only for acetic acid.) The experimental value of -4G10 for citric acid lies only 0.9 kJ mol-' higher than the value predicted by this relation. This difference lies still within the margin of the possible deviations. It may be noted that, with the exception of citric acid, all the acids in Table VI bear only one carboxylic group and are thus unable to form more than one hydrogen bond with the anions. In the case of citric acid, the overall kl constant is related to the formation constant klsingle of a single bond and to the constant kchelation governing the formation of a second bond by the same ligand molecule on the same anion by the expression
kl = klsingle(1 + kchelation) From the data above it thus appears that RT In (1 + kchelation) does not surpass an order of magnitude of 0.9 kJ mol-l. This means that kchelationis smaller than 0.5. Similar conclusions can be drawn for the interaction of citric acid with Br- where 4G10 lies only 0.9 kJ mol-l above the line which describes the influence of the pKa on 4G10 for the
I
8G
E
2
I
9
d
Lo rl
d U
The Journal of Physical Chemistty, Vol. 82,No. 17, 1978
Interactions of Citric Acid with Anions in Acetonitrile
l
-12
o
1
/
r
/
Figure 4. Difference between the standard free energy for the addition of a ligand molecule on the first and on the second site of anions vs. the standard free energy of the first bond: 0 , phenols on CI- in acetonitrile (ref 1); 0, benzoic and acetic acids on phenols in acetonitrile (ref 15a and 15b); H,citric acid on HS0,- in acetonitrile (this work H,citric acid on CI- in acetonitrile (this work); X, phenols on I-(R4N ) in CCI, (ref 3); A, phenols on Br- (R,N+) in CC14 (ref 3).
1;
complexes with a single bond. In the case of NO3-, the point lies even below the predicted value. There are insufficient data available in the case of HS04- to draw definitive conclusions, but as the increase of -AGIO in going from benzoic acid to citric acid is of the same order of magnitude as for the three previous ions, one can infer that here again, the chelation is not important. No data exist for comparing with monobasic acids for I-. Owing to the low value of kl-,an important chelation seems improbable a priori. It may thus be concluded that the majority of the 1:l complexes of citric acid with these anions are not chelated (kchelation
