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Ultraviolet spectra and complex formation in mixtures of polar organic solvents and aromatic hydrocarbons
Spectrochimiea Acta,1966, Vol.22.pp.77to80. Pergamon PressLtd. Printed inNorthern Imland
Ultraviolet spectra and complex formation in mix&es of polar organic solvents and aromatic hydrocarbons R. F. WEIMER and J. M. PRAUSNITZ Chemical Engineering Department, University of California, Berkeley (Received 29 May 1966) Abstract_Ultraviolet spectra at 25°C have been measuredfor mixtures of p-xylene and thirteen polar organic solvents dissolved in n-hexane. The data support the hypothesis that weak complexes are formed between aromatic hydrocarbons and such common polar solvents as aliphatic ketonee, nitroparaff?nsand nitriles. When reduced with Dn~ao’s equation the spectral data yield equilibrium constants for complex formation in the range 0*03-O-3 litem/gmol. SPECIFIC chemical interactions between common polar organic solvents and unsaturated hydrocarbons, can contribute to the thermodynamic selectivity of these solvents, for unsaturated, relative to saturated, hydrocarbons [l]. Since many sromstic molecules containing strongly polar functional groups, such as the nitro, cyano, and carbonyl groups, are known to act as electron acceptors relative to benzene and its methylated derivations [2, 31, it is expected that aliphatic materials containing these groups can also be capable of accepting charge. A standard method for the determination of donor-acceptor complexing in solution is through ultraviolet spectrophotometry. This method, initially described by BENESI and HILDEBRAND [4] and reviewed by BRIEGLEB [3], has been applied, in the present study, to the spectra of a number of mixtures of polar organic solvents, withp-xylene as the electron donor species and n-hexane as the inert diluent. EXPERIMENTAL
PROCEDURE
Optical measurements were made with a Beckman DU ultreviolet spectrophotometer in spectral regions, where enhanced absorption of the polar solvent+xylene solutions indicated the formation of complexes. The absorbance data were reproducible to within f l%, and were reduced by the general equation for 1: 1 complexing derived by DRAGO and ROSE [5]: lCAOCDO A -
lCa”cs -
CA0i- CD0
lCDo~D = (cc -
E__,-
EJ + KC(cC -
1
Ed -
eD) ’
(1)
where I is the path length, E is the molar extinction coefficient, C” is the initial concentration and K the equilibrium constant; the subscripts A, D and C stand for acceptor, donor and complex. The concentration ranges of polar acceptor and [l] R. ANDERSON, R. CAMBIO and J. M. PRAUSNITZ, A.I.Ch.E.J. [2] [3] [4] [6]
L. G. H. R.
8, 66 (1962). J. ANDREWS, Chmm Rev. 54, 713 (1954). BRIEOLEB, Elektronen-Donator-Acceptor-Komplexe, Berlin Springer-Verlag (1961). BENESI md J. H. HILLDEB-, J. Am. Chem. Sot. 71, 2703 (1949). S. DRAGO md N. J. ROSE, J. Am. Chem. Sot. 81,6138 (1969). 77
R. F. WEIMER
78
and J. M. PRAUSNITZ
Table 1. Concentrationranges of polar solvent and p-xylene in n-hexane
Polar solvent Acetone Cyclohexanone Triethyl phosphate Methoxyacetone Cyclopentanone y-Butyrolactone N-methyl pyrrolidone Propiontrile Nitromethane Nitroethane 2-Nitropropane Citraconicanhydride 2-Nitro-S-methyl propsLne
Concentrationrange of polar solvent, gmol/liter
Concentrationrange of p-xylene, gmol/liter
0.40-0.56 0*54-0.95 0.94-2.80 0.23-0.85 0.32-0.45 1.25-6.86 1.65-9.95 2.07-11.25 o-52-0.53 0.27-0.29 O-35-0.36 0.09-0.24 0.30-0.33
1.29-6.45 1.28-6.45 0*16-0~17 1.28-6.50 1.28-6.46 O-25-4.83 0.33-4.85 0~16-0~17 1.29-6.45 1.29-6.46 1.29-5.20 1.31-6.48 1.30-6.49
p-xylene used in the investigation of each system ctre listed in Table 1; the large values of the total concentration of acceptor zlnd donor were needed in order to obtain a significant amount of complex. All of the experimental solvents used were of the highest purity commercially available, and those which were not spectra grade were freshly distilled before use. In most cases the path length was 1 mm but in some cases it was 1 cm. A more detailed account of the experimental procedure may be found elsewhere [6]. RESULTS The equilibrium constants and extinction coefficients for the complexes studied are listed in Table 2. A comparison of the results for systems containing acetone or nitroethane with equilibrium constants previously obtained by ANDERSON [l, 71, shown in Table 3, indicates that the strength of the complex increases as the donor strength of the aromatic increases; therefore, the polar solvents are in fact acting as electron acceptors. Furthermore, measurement of the temperature variation of the equilibrium constants for the acetone-p-xylene and nitroethane-pxylene complexes between -5 and +4O”C indicated that the heat of formation of these complexes is about 2-3 kcal/mole, which is consistent with the formation of a weak chemical bond. The equilibrium constants reported in Table 2 are considerably smaller than those quoted by BRIEGLEB [3] for complexes between sromatic donors and strong acceptors such as trinitrobenzene, iodine, tetracyanoethylene, chlorenil and derivatives of quinone. Most previous studies on complexes have been made with a constant acceptor and a variety of donors; the data reported in Table 2 are different since they are concerned with a variety of acceptors and a constant donor. [S] R. F. WEIMER, Ph.D. Dissertation, Department of Chemical Engineering, University of California, Berkeley ( 1965). [7] R. ANDERSON, Ph.D. Dissertation, Department of Chemical Engineering, University of California, Berkeley (1961).
Polar organic solvents and aromatic hydrocarbons
79
Table 2. Spectroscopicequilibrium constants and extinction coefficientsfor polar solvent+-xylene complexes at 25°C KO liter/gmol
Polar Solvent
I,g%c~
80 - l - aI z/g&n-m
Acetone
288 289 290
0.26 f 0.06 0.26 & 0.06 0.24 f 0.06
12.96 12.66 12.36
4.66 4.36 4.40
Cyclohexanone
286 288 290
0.14 f 0.16 f 0.16 f
0.06 0.06 0.06
14.14 14.24 14.36
6.76 6.68 6.80
Triethyl phosphate
229 230
0.16 f 0.12 f
0.06 0.06
0.08 0.08
Methoxyacetone
286 288 290
0.13 f 0.13 f 0.10 f
0.07 0.06 0.06
14.96 14.70 14.46
6.24 6.30 6.12
Cyclopentmone
296 306
0.10 f 0.14 f
0.03 0.06
16.16 12.60
7.79 4.29
y-Butyrolmtone
285 286
0.09 f 0.09 f
0.06 0.06
0.02 0.02
2.6 1.7
N-methyl pymolidone
285 286
0.09 * 0.06 0.09 + 0.06
0.02 0.02
Propionitrile
229 229.6 230
0.071 f O-063 f 0.061 f
0.002 0.002 0.002
Nitromethane
286 288 290
0.059 f 0.060 f 0.066 f
0.010 0.010 0.011
14.4 13.3 12.4
48 47 39
286 288 290
0.057 f 0.012 0.064 $= 0.012 0.062 & 0.012
17.8 16.7 16.8
42.9 39.1 34.6
2-Nitropropane
286 288 290
0.062 f 0.042 f 0.040 f
19.6 18.6 17.6
32 41 39
Citraconic anbydride
343 346 347
0.052 k 0.003 0.041 f 0.003 0.027 f 0.003
0.38 0.23 0.13
173 166 179
2.Nitro-2-methyl-propane
286 288 290
0.031 & 0.016 0.029 f 0.017 0.022 + 0.017
21.6 20.86 20.0
62 63 60
0.024 0.021 0.022
0.044 0.041 0.039
111 71
44 28 137 119 90
Bubscript A refers to acceptor, D to donor and C to complex.
Table 3. Equilibrium constants for complexes of polar acceptors with polymethylbenzenes at 25°C Acceptor
Benzene
Acetone Nitroethane Nitromethane
0.07 f 0.04 0.020 f 0.010 0.016 f O*OlO
K,, lit&m01 Toluene p-Xylene 0.022 f 0.010 -
0.25 f 0.06 0.054 f 0.012 0.065 f 0.010
Mesitylene 0.31 f 0.20 0.093 * 0.020 0.048 f 0.020
80
R. F. WEIMER and J. M. PRAUSNITZ
Furthermore, the acceptors studied are all common organic solvents which are used in many typical chemical operations. The reduction of spectral data with Equation (1) involves at least three important assumptions: that only 1: 1 complexes are formed; that BEER’S law is valid for the absorption of the complex as well as for that of the donor and acceptor; and that the ratio of the activity coefficient of the complex to those of the uncomplexed donor and acceptor is independent of concentration. Because of the uncertainties known to be associated with these assumptions [6, 81 the equilibrium constants listed in Table 1 should be considered to be no more than a rough approximation of the strength of electron donor-acceptor interaction in polar solventaromatic hydrocarbon solutions, and they probably should not be used for making quantitative comparisons between polar species containing different functional groups. The data reported here do, however, support the hypothesis that common polar organic solvents form weak complexes with unsaturated hydrocarbons. AckmmdedgemRntThe authors are grateful to the National Science Foundation and to the National Institutes of Health for Cmncial support. [8] N. B. JURINSKIand P. A. D. DEMAINE,J. Am. Chem. Sot. 86, 3217 (1964).
Ultraviolet spectra and complex formation in mix&es of polar organic solvents and aromatic hydrocarbons R. F. WEIMER and J. M. PRAUSNITZ Chemical Engineering Department, University of California, Berkeley (Received 29 May 1966) Abstract_Ultraviolet spectra at 25°C have been measuredfor mixtures of p-xylene and thirteen polar organic solvents dissolved in n-hexane. The data support the hypothesis that weak complexes are formed between aromatic hydrocarbons and such common polar solvents as aliphatic ketonee, nitroparaff?nsand nitriles. When reduced with Dn~ao’s equation the spectral data yield equilibrium constants for complex formation in the range 0*03-O-3 litem/gmol. SPECIFIC chemical interactions between common polar organic solvents and unsaturated hydrocarbons, can contribute to the thermodynamic selectivity of these solvents, for unsaturated, relative to saturated, hydrocarbons [l]. Since many sromstic molecules containing strongly polar functional groups, such as the nitro, cyano, and carbonyl groups, are known to act as electron acceptors relative to benzene and its methylated derivations [2, 31, it is expected that aliphatic materials containing these groups can also be capable of accepting charge. A standard method for the determination of donor-acceptor complexing in solution is through ultraviolet spectrophotometry. This method, initially described by BENESI and HILDEBRAND [4] and reviewed by BRIEGLEB [3], has been applied, in the present study, to the spectra of a number of mixtures of polar organic solvents, withp-xylene as the electron donor species and n-hexane as the inert diluent. EXPERIMENTAL
PROCEDURE
Optical measurements were made with a Beckman DU ultreviolet spectrophotometer in spectral regions, where enhanced absorption of the polar solvent+xylene solutions indicated the formation of complexes. The absorbance data were reproducible to within f l%, and were reduced by the general equation for 1: 1 complexing derived by DRAGO and ROSE [5]: lCAOCDO A -
lCa”cs -
CA0i- CD0
lCDo~D = (cc -
E__,-
EJ + KC(cC -
1
Ed -
eD) ’
(1)
where I is the path length, E is the molar extinction coefficient, C” is the initial concentration and K the equilibrium constant; the subscripts A, D and C stand for acceptor, donor and complex. The concentration ranges of polar acceptor and [l] R. ANDERSON, R. CAMBIO and J. M. PRAUSNITZ, A.I.Ch.E.J. [2] [3] [4] [6]
L. G. H. R.
8, 66 (1962). J. ANDREWS, Chmm Rev. 54, 713 (1954). BRIEOLEB, Elektronen-Donator-Acceptor-Komplexe, Berlin Springer-Verlag (1961). BENESI md J. H. HILLDEB-, J. Am. Chem. Sot. 71, 2703 (1949). S. DRAGO md N. J. ROSE, J. Am. Chem. Sot. 81,6138 (1969). 77
R. F. WEIMER
78
and J. M. PRAUSNITZ
Table 1. Concentrationranges of polar solvent and p-xylene in n-hexane
Polar solvent Acetone Cyclohexanone Triethyl phosphate Methoxyacetone Cyclopentanone y-Butyrolactone N-methyl pyrrolidone Propiontrile Nitromethane Nitroethane 2-Nitropropane Citraconicanhydride 2-Nitro-S-methyl propsLne
Concentrationrange of polar solvent, gmol/liter
Concentrationrange of p-xylene, gmol/liter
0.40-0.56 0*54-0.95 0.94-2.80 0.23-0.85 0.32-0.45 1.25-6.86 1.65-9.95 2.07-11.25 o-52-0.53 0.27-0.29 O-35-0.36 0.09-0.24 0.30-0.33
1.29-6.45 1.28-6.45 0*16-0~17 1.28-6.50 1.28-6.46 O-25-4.83 0.33-4.85 0~16-0~17 1.29-6.45 1.29-6.46 1.29-5.20 1.31-6.48 1.30-6.49
p-xylene used in the investigation of each system ctre listed in Table 1; the large values of the total concentration of acceptor zlnd donor were needed in order to obtain a significant amount of complex. All of the experimental solvents used were of the highest purity commercially available, and those which were not spectra grade were freshly distilled before use. In most cases the path length was 1 mm but in some cases it was 1 cm. A more detailed account of the experimental procedure may be found elsewhere [6]. RESULTS The equilibrium constants and extinction coefficients for the complexes studied are listed in Table 2. A comparison of the results for systems containing acetone or nitroethane with equilibrium constants previously obtained by ANDERSON [l, 71, shown in Table 3, indicates that the strength of the complex increases as the donor strength of the aromatic increases; therefore, the polar solvents are in fact acting as electron acceptors. Furthermore, measurement of the temperature variation of the equilibrium constants for the acetone-p-xylene and nitroethane-pxylene complexes between -5 and +4O”C indicated that the heat of formation of these complexes is about 2-3 kcal/mole, which is consistent with the formation of a weak chemical bond. The equilibrium constants reported in Table 2 are considerably smaller than those quoted by BRIEGLEB [3] for complexes between sromatic donors and strong acceptors such as trinitrobenzene, iodine, tetracyanoethylene, chlorenil and derivatives of quinone. Most previous studies on complexes have been made with a constant acceptor and a variety of donors; the data reported in Table 2 are different since they are concerned with a variety of acceptors and a constant donor. [S] R. F. WEIMER, Ph.D. Dissertation, Department of Chemical Engineering, University of California, Berkeley ( 1965). [7] R. ANDERSON, Ph.D. Dissertation, Department of Chemical Engineering, University of California, Berkeley (1961).
Polar organic solvents and aromatic hydrocarbons
79
Table 2. Spectroscopicequilibrium constants and extinction coefficientsfor polar solvent+-xylene complexes at 25°C KO liter/gmol
Polar Solvent
I,g%c~
80 - l - aI z/g&n-m
Acetone
288 289 290
0.26 f 0.06 0.26 & 0.06 0.24 f 0.06
12.96 12.66 12.36
4.66 4.36 4.40
Cyclohexanone
286 288 290
0.14 f 0.16 f 0.16 f
0.06 0.06 0.06
14.14 14.24 14.36
6.76 6.68 6.80
Triethyl phosphate
229 230
0.16 f 0.12 f
0.06 0.06
0.08 0.08
Methoxyacetone
286 288 290
0.13 f 0.13 f 0.10 f
0.07 0.06 0.06
14.96 14.70 14.46
6.24 6.30 6.12
Cyclopentmone
296 306
0.10 f 0.14 f
0.03 0.06
16.16 12.60
7.79 4.29
y-Butyrolmtone
285 286
0.09 f 0.09 f
0.06 0.06
0.02 0.02
2.6 1.7
N-methyl pymolidone
285 286
0.09 * 0.06 0.09 + 0.06
0.02 0.02
Propionitrile
229 229.6 230
0.071 f O-063 f 0.061 f
0.002 0.002 0.002
Nitromethane
286 288 290
0.059 f 0.060 f 0.066 f
0.010 0.010 0.011
14.4 13.3 12.4
48 47 39
286 288 290
0.057 f 0.012 0.064 $= 0.012 0.062 & 0.012
17.8 16.7 16.8
42.9 39.1 34.6
2-Nitropropane
286 288 290
0.062 f 0.042 f 0.040 f
19.6 18.6 17.6
32 41 39
Citraconic anbydride
343 346 347
0.052 k 0.003 0.041 f 0.003 0.027 f 0.003
0.38 0.23 0.13
173 166 179
2.Nitro-2-methyl-propane
286 288 290
0.031 & 0.016 0.029 f 0.017 0.022 + 0.017
21.6 20.86 20.0
62 63 60
0.024 0.021 0.022
0.044 0.041 0.039
111 71
44 28 137 119 90
Bubscript A refers to acceptor, D to donor and C to complex.
Table 3. Equilibrium constants for complexes of polar acceptors with polymethylbenzenes at 25°C Acceptor
Benzene
Acetone Nitroethane Nitromethane
0.07 f 0.04 0.020 f 0.010 0.016 f O*OlO
K,, lit&m01 Toluene p-Xylene 0.022 f 0.010 -
0.25 f 0.06 0.054 f 0.012 0.065 f 0.010
Mesitylene 0.31 f 0.20 0.093 * 0.020 0.048 f 0.020
80
R. F. WEIMER and J. M. PRAUSNITZ
Furthermore, the acceptors studied are all common organic solvents which are used in many typical chemical operations. The reduction of spectral data with Equation (1) involves at least three important assumptions: that only 1: 1 complexes are formed; that BEER’S law is valid for the absorption of the complex as well as for that of the donor and acceptor; and that the ratio of the activity coefficient of the complex to those of the uncomplexed donor and acceptor is independent of concentration. Because of the uncertainties known to be associated with these assumptions [6, 81 the equilibrium constants listed in Table 1 should be considered to be no more than a rough approximation of the strength of electron donor-acceptor interaction in polar solventaromatic hydrocarbon solutions, and they probably should not be used for making quantitative comparisons between polar species containing different functional groups. The data reported here do, however, support the hypothesis that common polar organic solvents form weak complexes with unsaturated hydrocarbons. AckmmdedgemRntThe authors are grateful to the National Science Foundation and to the National Institutes of Health for Cmncial support. [8] N. B. JURINSKIand P. A. D. DEMAINE,J. Am. Chem. Sot. 86, 3217 (1964).