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Solvent effect on acid salt of dichloroacetic acid containing short and asymmetric OH…O bond
Journal of Molecular Structure, 143 (1986) 333-336 Elsevier Science Publishers B.V., Amsterdam -Printed
333 in The Netherlands
SOLVENT EFFECT ON ACID SALT OF DICHLOROACETICACID CONTAINING SHORT AND ASXWETRIC OH**=0 BOND Z. DEGA-SZAFRAN and M. SZAl?RAN Chemistry Dept., A. Mickiewicz University, 60780 Poznarl(Poland) ABSTRACT The IR and 'H NMR spectra are bis-dichloroacetatein five dr city rity (& = 2.27 - 64.4) and bas;Y tegrated intensity (A) and the centre of gravity (3 absorption show weak dependence upon solvent used. The chemical shift (6) depends on the proton acceptor properties of solvents and is consistentwith an equilibrium (Bu) NA -tAH =+ (Bu)qN(AHA).The weak dependence of A and 3 upon the so4vent polarity is consistent with the single minimum energy surface postulated in the literature for the strong OH0 bonds. INTRODUCTION In this paper we extend our investigationsof the solvent effect (ref.l-3) on IR and 'H NMR spectra of (nC4H9)4NH(OOCCHC12), in order to get new evidence of this effect on the strong hydrogen bond. The'hydrogen bond in KH(OGCCHC1,)2is short (2.498 8) and the hydrogen atom is placed asymmetricallywith respect to the oxygen atoms (ref.4).This diffractionwork is supported by studies of 35Cl NQR (ref.5) and IR (ref.6) spectra. In solution the interactionsof (Bu)~N(OOCCHC~~)with C12CHCOOH are described by the following equilibria: (Bu) NA + HA A Kl 4 'di (Bu)~N(AHA) %A assoc
(Bu)~N(AHA)
(1)
(Bu)~N+ + (AHA)-
(2)
The values of log K, = 4.5 in nitromethsne (ref.7) and enthalpy -AH= 25.52 KJ/mol in propylene carbonate (ref.8) have been determined. EXPERIMENTAL (Bu)~N(OOCCHC~~)was prepared by standard method and recrystallized from ethylacetate,m.p.<30°C. (Bu)~NH(OOCCHC~~)~was prepared by dissolving (Bu)~N(OOCCHC~~)in C12CHCOOH and removing the excess acid by evaporation at 1OO'C in a vacuum and then by'cooling slowly, whereupon white crystal was obtained, m.p. 31'~. 0022-2860/86/$03.50
6 1986 Elsevier Science Publishers B.V.
334 The solvent were purified and dried by standard methods and stored over molecular sieves (3A or 4A). IR spectra were measured on a Perkin-Elmer 580 spectrophotometer in cells with I(Brwindows. Two methods are used to calculate integrated intensities (A = 1/C.lJlog(Tbl/T)d3) and centres of gravity(g=jVlog(Tbl/T)dV/ log(Tbl/T)dJ), where Tbl and T are transmittance of the base line J and the investigated compounds, respectively, C is the concentration (0.2 mol/liter), 1 the path length in centimeters (1 = 0.0131), and L?the frequency in cm-'. The first, a continuous proton absorption was separated graphically from sharp bands ascribed to other internal vibrations (see e.g. ref.2). The second, the integrated intensity is given by A = A(nuj4N(AHAj - A(Bu)4NA - AHA. The spectra of_, both butylammonium salts were integrated in the region 2800-400 cm and the spectrum of acid in the region 2300-400 cm-1 . The compensation value of the solvent was used as a base line. 1 H NkiR spectra were measured on a Tesla ES 46'7 spectrometer. RESULTS AND IjISCUSSION As illustrated in Figure formation of the acid salt caused a continuum absorption over the frequency region from 2800 to 400 cm-1 . The shape of absorption of the acid salt is practically independent of the solvent used, but its integrated intensity slightly varies with the solvent (Table 1). The most sensitive to solvent is integrated intensity of the acid. The larger value observed in acetonitrile is caused by interaction of the acid with solvent molecules.
3500
3000
2500
2000
1800
1600
l&O0
1200
1000
800
600
cm-l
Fig. IR spectra of C12CHCOOR (......I, (C4H9)4N(OBCCHC12) (- - - -1, and (C4H9)4N(C12CHCOOH~OOCCHC12) ( -1
in benzene.
336
TABI;G1 Solvent effect on the integrated intensity (AX10s4 1 m01"cm'~) Compounds (Bu)~NH(OOCCHCL~)~ (Bu)~N(OOCCHC~~) C12CHCOOH
'6%
CH2C12
CH3N02
CH3CN
27.41 6.82 4.81
27.16 6.72 4.79
27.39
28.12 6.91 8.22
5.63
TABIZ 2 Data and values for (Bu)4Xl(OOCCHC12)2 Solvent
&
6
DN a
C6H6
2.27
CH2C12
8.93
0.1
b
Method I A~l0'~ ij
Lethod II A~l0'~ 3
19.05
19.13
19.23
1176
15.78
1165.7
18.20
18.39
20.46
1191
15.65
1156.5
35.87
2.7
18.23
18.33
20.09
1168
15.88'
l143.8d
37.5
14.1
17.90
18.37
21.28
1182
16.42'
1120.2'
64.4
15.1
17.83
18.18
a 0.2 M solutions; b 0.2 kisolutions with 0.5 u of (8~)~N(00CChCl~); ' calculatedwith AAh = 4.79; d calculatedwith A (Bu)~NA = 6*72 and AAh = 4.79 ; e PC - propylene carbonate. Table 2 lists values of A and 3 of the continuous absorption. The difference between two methods is ca. 4 units in intensity and ca. 30 cm-' in frequency. The larger intensitiesobtained in method I are probably caused by superpositionof broad absorption of the acid (see Fig., region 1450-600 cm-') and uncertainty in the carbonyl region. Iogansen (ref.9) noted that, for ca. 250 complexeswith hydrogen bond, there appear to be relationship between A H and integrated intensity. The similar integrated intensitiesand centres of gravity (Table 2) suggest that the enthalpy of hydrogen bond in the acid salt does not change with the solvent. Another interestingfeature are two C=O stretching bands at 1750 and 1712 cm-1 , which indicate that the hydrogen bond is asymmetric. The chemical shift of H-bonded protons in the investigatedsalt varies with the solvent (Table 2). These variations can be explained by eq. (1). On addition of (Bu)~N(OOCCHC~~)the signal shifts to lower field. Liagnitudeof these shifts is proportionalto DN.
336
Another explanationof the observed variation of the chemical shift with the solvent can be based on the equilibrium (2). Jones and Dyer (ref.10)have studied the System of triflUOrOaCetates in trifluoroaceticacid and found that the value of the chemical shift of the H-bonded protons varied according to the cation (6 = 15.73 (Li), 16.94 (Na), 19.17 (K), 19.44 (Rb), 19.80 (Cs) and 19.66 (Re4N). They concluded that the chemical shift depends largely upbn the nearness of approach of the two ions; a maximum downfield chemical shift is expected from the "free" (AHA)- anion. Since solvents of high dielectric constant promote the dissociation ion pair (ref.ll),one can expect a downfield shift with an increase in the solvent dielectric constant, if the equilibrium (2) is operating.Tha data in Table 2 shows that this is not the case. In complexes of pyridinium with carboxylicacida both 3 and 6 strongly depend upon solvent polarity (ref.l,2).This is a consequence of a tautomerismbetween two forms with different &values: + A-H***B 'c- A-***H-B . The weak sensitivityof A, 3 and 6 on solvent polarity in the investigatedsalt and complexes of some pyridine N-oxides with trifluoroacetic acid (ref.3) is consistentwith a "quasi-symmetrical" hydrogen bond postulated in the literaturefor complexes of oxygen bases. REFERENCES Z. Dega-Szafranand E. Dulewicz, Org. Kagn. Reson., 16 (1981) 214. Z. Dega-Szafran,E. Dulewicz and MI.Szafran, J. Chem. Sot., Perkin Trans. 2, 1984, 1997. B. Brycki and 13.Szafran, J. mol. Liquids, 29 (lY84) 135. D. Had%, L. Leban B. Ore1 m. Iwata and J.lir. Williams, J. Cryst. kiol.Structure, 9 flY7r) 113. H. Ratajczak, W.J. Orville-Thomasand I. ChoLoniewska,Chem. Phys. Letters, 45 (1977) 208. Z. Pawlak and L. Sobcsyk, Adv. Mol. Relax. Processes, 5 (1973) 99. B.A. Korolev and E.I. Kashkovskaya,Zh. Obshch. Khim., 49 (1982) 909. Z. Pawlak and R.G. Bates, J. Chem. Thermodynamics,14 (1982) 1035. A.V. Iogansen, in N.D. Sokolov (Ed) Hydrogen Bond, kioacow,1981, pp. 112-155. R.G. Jones and J.R. Dyer, J. Am. Chem. Sot., 95 (1973) 2465. C. Reichardt, Solvent Effects in Organic Chemistry, Verlag Chemie, Weinheim, New York, 1979. (This work was supported by Polish Academy of Sciences; kR.I.9.4.4.4)
333 in The Netherlands
SOLVENT EFFECT ON ACID SALT OF DICHLOROACETICACID CONTAINING SHORT AND ASXWETRIC OH**=0 BOND Z. DEGA-SZAFRAN and M. SZAl?RAN Chemistry Dept., A. Mickiewicz University, 60780 Poznarl(Poland) ABSTRACT The IR and 'H NMR spectra are bis-dichloroacetatein five dr city rity (& = 2.27 - 64.4) and bas;Y tegrated intensity (A) and the centre of gravity (3 absorption show weak dependence upon solvent used. The chemical shift (6) depends on the proton acceptor properties of solvents and is consistentwith an equilibrium (Bu) NA -tAH =+ (Bu)qN(AHA).The weak dependence of A and 3 upon the so4vent polarity is consistent with the single minimum energy surface postulated in the literature for the strong OH0 bonds. INTRODUCTION In this paper we extend our investigationsof the solvent effect (ref.l-3) on IR and 'H NMR spectra of (nC4H9)4NH(OOCCHC12), in order to get new evidence of this effect on the strong hydrogen bond. The'hydrogen bond in KH(OGCCHC1,)2is short (2.498 8) and the hydrogen atom is placed asymmetricallywith respect to the oxygen atoms (ref.4).This diffractionwork is supported by studies of 35Cl NQR (ref.5) and IR (ref.6) spectra. In solution the interactionsof (Bu)~N(OOCCHC~~)with C12CHCOOH are described by the following equilibria: (Bu) NA + HA A Kl 4 'di (Bu)~N(AHA) %A assoc
(Bu)~N(AHA)
(1)
(Bu)~N+ + (AHA)-
(2)
The values of log K, = 4.5 in nitromethsne (ref.7) and enthalpy -AH= 25.52 KJ/mol in propylene carbonate (ref.8) have been determined. EXPERIMENTAL (Bu)~N(OOCCHC~~)was prepared by standard method and recrystallized from ethylacetate,m.p.<30°C. (Bu)~NH(OOCCHC~~)~was prepared by dissolving (Bu)~N(OOCCHC~~)in C12CHCOOH and removing the excess acid by evaporation at 1OO'C in a vacuum and then by'cooling slowly, whereupon white crystal was obtained, m.p. 31'~. 0022-2860/86/$03.50
6 1986 Elsevier Science Publishers B.V.
334 The solvent were purified and dried by standard methods and stored over molecular sieves (3A or 4A). IR spectra were measured on a Perkin-Elmer 580 spectrophotometer in cells with I(Brwindows. Two methods are used to calculate integrated intensities (A = 1/C.lJlog(Tbl/T)d3) and centres of gravity(g=jVlog(Tbl/T)dV/ log(Tbl/T)dJ), where Tbl and T are transmittance of the base line J and the investigated compounds, respectively, C is the concentration (0.2 mol/liter), 1 the path length in centimeters (1 = 0.0131), and L?the frequency in cm-'. The first, a continuous proton absorption was separated graphically from sharp bands ascribed to other internal vibrations (see e.g. ref.2). The second, the integrated intensity is given by A = A(nuj4N(AHAj - A(Bu)4NA - AHA. The spectra of_, both butylammonium salts were integrated in the region 2800-400 cm and the spectrum of acid in the region 2300-400 cm-1 . The compensation value of the solvent was used as a base line. 1 H NkiR spectra were measured on a Tesla ES 46'7 spectrometer. RESULTS AND IjISCUSSION As illustrated in Figure formation of the acid salt caused a continuum absorption over the frequency region from 2800 to 400 cm-1 . The shape of absorption of the acid salt is practically independent of the solvent used, but its integrated intensity slightly varies with the solvent (Table 1). The most sensitive to solvent is integrated intensity of the acid. The larger value observed in acetonitrile is caused by interaction of the acid with solvent molecules.
3500
3000
2500
2000
1800
1600
l&O0
1200
1000
800
600
cm-l
Fig. IR spectra of C12CHCOOR (......I, (C4H9)4N(OBCCHC12) (- - - -1, and (C4H9)4N(C12CHCOOH~OOCCHC12) ( -1
in benzene.
336
TABI;G1 Solvent effect on the integrated intensity (AX10s4 1 m01"cm'~) Compounds (Bu)~NH(OOCCHCL~)~ (Bu)~N(OOCCHC~~) C12CHCOOH
'6%
CH2C12
CH3N02
CH3CN
27.41 6.82 4.81
27.16 6.72 4.79
27.39
28.12 6.91 8.22
5.63
TABIZ 2 Data and values for (Bu)4Xl(OOCCHC12)2 Solvent
&
6
DN a
C6H6
2.27
CH2C12
8.93
0.1
b
Method I A~l0'~ ij
Lethod II A~l0'~ 3
19.05
19.13
19.23
1176
15.78
1165.7
18.20
18.39
20.46
1191
15.65
1156.5
35.87
2.7
18.23
18.33
20.09
1168
15.88'
l143.8d
37.5
14.1
17.90
18.37
21.28
1182
16.42'
1120.2'
64.4
15.1
17.83
18.18
a 0.2 M solutions; b 0.2 kisolutions with 0.5 u of (8~)~N(00CChCl~); ' calculatedwith AAh = 4.79; d calculatedwith A (Bu)~NA = 6*72 and AAh = 4.79 ; e PC - propylene carbonate. Table 2 lists values of A and 3 of the continuous absorption. The difference between two methods is ca. 4 units in intensity and ca. 30 cm-' in frequency. The larger intensitiesobtained in method I are probably caused by superpositionof broad absorption of the acid (see Fig., region 1450-600 cm-') and uncertainty in the carbonyl region. Iogansen (ref.9) noted that, for ca. 250 complexeswith hydrogen bond, there appear to be relationship between A H and integrated intensity. The similar integrated intensitiesand centres of gravity (Table 2) suggest that the enthalpy of hydrogen bond in the acid salt does not change with the solvent. Another interestingfeature are two C=O stretching bands at 1750 and 1712 cm-1 , which indicate that the hydrogen bond is asymmetric. The chemical shift of H-bonded protons in the investigatedsalt varies with the solvent (Table 2). These variations can be explained by eq. (1). On addition of (Bu)~N(OOCCHC~~)the signal shifts to lower field. Liagnitudeof these shifts is proportionalto DN.
336
Another explanationof the observed variation of the chemical shift with the solvent can be based on the equilibrium (2). Jones and Dyer (ref.10)have studied the System of triflUOrOaCetates in trifluoroaceticacid and found that the value of the chemical shift of the H-bonded protons varied according to the cation (6 = 15.73 (Li), 16.94 (Na), 19.17 (K), 19.44 (Rb), 19.80 (Cs) and 19.66 (Re4N). They concluded that the chemical shift depends largely upbn the nearness of approach of the two ions; a maximum downfield chemical shift is expected from the "free" (AHA)- anion. Since solvents of high dielectric constant promote the dissociation ion pair (ref.ll),one can expect a downfield shift with an increase in the solvent dielectric constant, if the equilibrium (2) is operating.Tha data in Table 2 shows that this is not the case. In complexes of pyridinium with carboxylicacida both 3 and 6 strongly depend upon solvent polarity (ref.l,2).This is a consequence of a tautomerismbetween two forms with different &values: + A-H***B 'c- A-***H-B . The weak sensitivityof A, 3 and 6 on solvent polarity in the investigatedsalt and complexes of some pyridine N-oxides with trifluoroacetic acid (ref.3) is consistentwith a "quasi-symmetrical" hydrogen bond postulated in the literaturefor complexes of oxygen bases. REFERENCES Z. Dega-Szafranand E. Dulewicz, Org. Kagn. Reson., 16 (1981) 214. Z. Dega-Szafran,E. Dulewicz and MI.Szafran, J. Chem. Sot., Perkin Trans. 2, 1984, 1997. B. Brycki and 13.Szafran, J. mol. Liquids, 29 (lY84) 135. D. Had%, L. Leban B. Ore1 m. Iwata and J.lir. Williams, J. Cryst. kiol.Structure, 9 flY7r) 113. H. Ratajczak, W.J. Orville-Thomasand I. ChoLoniewska,Chem. Phys. Letters, 45 (1977) 208. Z. Pawlak and L. Sobcsyk, Adv. Mol. Relax. Processes, 5 (1973) 99. B.A. Korolev and E.I. Kashkovskaya,Zh. Obshch. Khim., 49 (1982) 909. Z. Pawlak and R.G. Bates, J. Chem. Thermodynamics,14 (1982) 1035. A.V. Iogansen, in N.D. Sokolov (Ed) Hydrogen Bond, kioacow,1981, pp. 112-155. R.G. Jones and J.R. Dyer, J. Am. Chem. Sot., 95 (1973) 2465. C. Reichardt, Solvent Effects in Organic Chemistry, Verlag Chemie, Weinheim, New York, 1979. (This work was supported by Polish Academy of Sciences; kR.I.9.4.4.4)