Interaction of pNitrosalicylic Acid and Ethylenediamine References and Notes J. L. Kurz and J. M. Farrar, J. Am. Chem. SOC.,91, 6057 (1969). J. Jordan, Rec. Chem. Prog., 19, 193 (1958). K. S. Pitzer, J. Am. Chem. SOC., 59, 2365 (1937). J. L. Kurz and G. J. Ehrhardt, J. Am. Chem. SOC., 97, 2259 (1975). (5) M. Braid, H. Iserson, and F. E. Lawler. J. Am. Chem. SOC., 76, 4027 (1954). (6) C. W. Davies, "Ion Association", Butterworths, London, 1962, p 41. (1) (2) (3) (4)
157 (7) D. H. Everett and W. F. K. Wynne-Jones, Trans. Faraday SOC., 35, 1380 (1939). (8) The total variation of the apparent pK.'s through the range 0 < I< 0.04 was ca. 0.02 at each temperature and, as expected, was linear within the experimental scatter. (9) A. A. Clifford, "Multivariate Error Analysis", Wiley, New York, N.Y., 1973. (10) W. J. Middleton and R. V. Lindsey, Jr., J. Am. Chem. SOC., 86, 4948 (1964).
Interaction of p-Nitrosalicylic Acid and Ethylenediamine in Mixed Solvents. A Proton Donor-Acceptor Equilibrium S. P. Moulik,* S. Ray, Department of Chemistry, Jadavpur University, Calcutta 700032, India
and A. R. D a s Depatfment of Physical Chemistry, Indian Associatlon for the Cultivation of Science, Calcutta 700032, India (Received June 25, 1974; Revised Manuscript Received June 23, 1975)
The proton transfer complex of p-nitrosalicylic acid with ethylenediamine has been studied in mixed solvents with a view toward understanding the electrostatic and specific effects of protic and aprotic solvents. The equilibrium constants have been determined and the thermodynamic parameters, AGO, AHo, and ASo, evaluated in several solvent systems. It has been observed that the solvent polarity fails to be a proper guide of complex formation. The formation can be appreciably arrested by the addition of a high polar solvent such as formamide.
Introduction A phenolic acid usually releases proton from the -OH group in a polar medium depending on the pK of the compound. In doing so, the anion exhibits a characteristic absorption band whose intensity can be a convenient measure of its dissociation. In a medium of low polarity electrostatic factors reduce the above ionization; in such a medium a withdrawing agent (an amine) is necessary for favorable release of the proton. The released proton is preferentially transferred to the base and a proton transfer complex or hydrogen bonded ion pair is formed. When the acid is not that weak, some of the released hydrogen ions may be accepted by the polar solvent molecules (of course in a less preferential manner than the added base), thus making the whole transfer process complex.' Besides, specific solvation, solute-solvent interaction, and solvent structural effects may also influence the ionization process.1,2 Since in a solvent medium the electrostatic or solvent polarity effect on the proton transfer to a base is complicated by the additional acceptance of the protons by the solvent molecules, a condition is, therefore, needed to reduce the second process to a minimum for the understanding of the first. Under normal conditions, p-nitrosalicylic acid has been observed to satisfy this by not releasing its phenolic proton even into pure water when a base is absent. An orderly measure of the electrostatic effects on the proton donor-acceptor equilibrium of it with an amine (ethylenediamine) is expected in mixed solvent media. Results of such investigations using 12 nonaqueous solvents are re-
ported in what follows. Such an investigation can be important for the basic understanding of the hydrogen bonded interactions between the acid and the basic groups in proteins responsible for various biological activities. Experimental Section Materials and Method. p-Nitrosalicylic acid of Eastman Organic Chemicals, mp 227-228OC) was recrysta!lized twice from conductivity water, and its melting point was checked for purity. All the solvents used were either E. Merck pro analysi or BDH AnalaR grades. They were further purified prior to use by standard chemical and physical treatments. The sample of urea (BDH, AnalaR) was used without further purification. Aqueous solutions were prepared using double distilled conductivity water (specific conductance = 1.5 X mho cm-' at 3OOC). Spectral measurements were taken in a Beckman DU spectrophotometer using 1-cm matched silica cells. The prepared samples were placed in a temperature controlled thermostatic bath (working in the deviation range of f0.2OC) for a period of 1 hr, and the absorbances were measured in the spectrophotometer having arrangements for circulating the water of the bath around the cell compartment. Two types of experiments were performed. In the first, p-nitrosalicylic acid samples of a fixed concentration were mixed with increasing amounts of ethylenediamine in a fixed solvent composition by placing equal amounts of the proton donor in different containers, and adjusting the final composition by the required addition of The Journal of Physical Chemistry, Vol. 80,No. 2, 1976
158
S.P. Moulik, S. Ray, and A. R. Das
the amine and the solvent. The absorbances of each of these solutions were then measured over a wide range of wavelengths. In the second set of measurements, the concentrations of the proton donor and the acceptor were kept constant, and the solvent compositions were varied over a wide range. The absorbances of these solutions were measured a t the characteristic wavelength of the complex, which is 414 nm in the present case.
Results Spectra of p-Nitrosalicylic Acid. As long as the phenolic proton is unreleased, the spectra of the acid in water and in nonaqueous solvents, viz., alcohols, acetone, dioxane, etc., are all almost alike with a broad band at 320 nm, falling sharply in the visible direction (practically zero absorbance for a 4.03 X M acid at 400 nm). In the presence of increasing concentrations of the diamine a characteristic band appears at 414 nm, and sharpens with a concomitant decrease of the 320-nm band of the undissociated acid. This is a general tendency of a proton transfer process in which one form changes into another with a noteworthy occurrence of an isosbestic point' (the isosbestic point in the present system is at 355 nm). Equilibrium Constant of the Complex. To understand the thermodynamics of the process, determination of the equilibrium constant of the complex is obvious. This was done in the following way. For the reaction
AH donor
+
K
B + A - . . .H+B acceptor complex
- Cc][ c b - Cc] (2) [Ca - c,], [cb - e,], and [c,]are the equilibrium
where concentrations of the donor acid, the acceptor base, and the 1:l complex, respectively. The concentration of the free base (equilibrium concentration) can seldom be determined. This was avoided by the use of large excess of the base, and eq 2 then becomes (3)
Now the measured absorbance is
A = ec[CcI
+ ca[Ca - Cc]
(4)
where c, and ca are the extinction coefficients of the complex and the proton donor, respectively (the concentration of the amine is assumed zero, which is a reality). Combining eq 3 and 4 we have ICa]/{A - €a[Ca]}= I/{€, - ea}
1/K{ec - %}[Cb]
(5)
Equation 5 can be further reduced to [Ca]/A
l/tc
+ I/Kc[Cb]
(6)
selecting a wavelength at which only the complex absorbs but the proton donor does not (in the present case the complex only absorbs a t 414 nrn and above). In all the measurements the concentrations of the base were kept nearly 100 times more than the p-nitrosalicylic acid and eq 6 was used to evaluate cc and K from the intercepts and the slopes, respectively. Representative least-squares plots are illustrated in Figure 1, and the comprehensive results are given in Table I with appropriate standard deviations. The very close fitting of the data to straight lines supports the presence of 1:l complex formation. Chances of higher order complexes were, however, not truly zero.' The formation of The Journal of Physical Chemistry, Vol. 80, No. 2, 1976
0 "
U
i
_
-
_
_
4
-
8
I
-
-
12
l/[cbl xio-'lit mole-'
Graphical determination of equilibrium constants of the complex in various mixed solvent media: (0, 0 ) water, open and half-closed points represent duplicate experiments; (m) urea, 37.25 wt %; (6)ethylene glycol, 14.24 wt %; ( 0 )methanol, 8.01 wt % (line I), 16.46 wt % (line 2); (A)ethanol, 9.66 wt %; ( 0 )1-propanol, 9.83 wt %; (@) 2-propanol, 12.15 wt %. Figure 1.
(1)
K = [CcI/[Ca
K = [CcI/ICa - Cc][cb]
.
8 n
the complex was maximum in pure water, decreased with the addition of nonaqueous solvents, and was appreciably low in the presence of small concentrations of acetone and dioxane. The protic solvents behaved differently than the aprotic ones. Dependence of K on Solvent Polarity. The dielectric constants reported in Table I were taken from the work of Aker10f.~Whenever direct data were not available, a method of interpolation was adopted. Those on dimethylformamide and formamide were taken from the work of Rohdewald and Moldner? and that of aqueous urea solutions were taken from the data of Wyman5 computed by Kundu and Majumdar.6 Dielectric constants of aqueous mixtures of acetonitrile are not avaiIable in the literature, therefore, the effects of this solvent could not be brought into the general line. To test the requirements of Coulomb's law, a plot of log K against D-l was made for the hydroxylic solvents; the dependence was not linear (Figure 2). An apparent conclusion is that solvent effects other than electrostatic are present. In this figure, the 30° set is seen to be above the 5 5 O set. It was observed that use of increased concentrations of p-nitrosalicylic acid yielded higher K values. This may then explain the increased K values observed for the 30° set which contained appreciably higher concentration of the proton donor. The altered activities of all the species at higher concentration of the donor probably made the equilibrium constant more for the 30° set than the 55' set. For a comfortable comparison a detailed knowledge of the activities of all the species is required. A sharp distinction among the mixed solvents appear in that the dipolar aprotic solvents (acetone and DMF) reduced the equilibrium constant more than the protic ones. Effect of Temperature. The effects of temperature on complex formation in various solvent media were studied
Interaction of pNitrosalicylic Acid and Ethylenediamine
159
TABLE I: Physicochemical Properties of the Complex in Various Environmentsa ec x 10-3,
M-'cm-'
Solvent Water Glycol-water (14.24) Methanol-water (8.01) Methanol-water (16.46) Methanol-water (25.21) Ethanol-water (9.66) Acetone-water (3.96) 1-Propanol-water (9.83) 2-Propanol-water (12.1 5) 2-Methyl-2-propanol-water (12.06) Dioxane-water (2.05) Urea-water (37.25)
15.06 i 12.79 f 12.17 f 8.67 f 5.59 f 11.61 f 11.93 f 9.77 f 9.91 f 11.61 i
0.16 0.38 0.90 0.16 0.17 0.10
0.43 0.10 0.15 0.55
D 76.7 72.5 73.4 69.3 65.2 71.4 74.4 70.0 68.2 66.7
K x lo-*, M-'
9.60 f 6.80 i 7.03 f 5.52 f 3.44 f 6.30 f 5.24 i 6.10 f 4.98 f 3.91 f
-AG", kcal mol-'
0.04 0.09 0.07 0.001 0.02 0.02 0.02 0.02 0.14 0.03
4.14 t 0.003 3.93 f 0.008 3.95 f 0.006 3.80 r 0.003 3.52 i 0.004 3.88 i: 0.002 3.77 It 0.002 3.86 r 0.002 3.74 f 0.02 3.60 i 0.005
74.2 23.31 f 2.36 1.70 f 0.09 3.09 i: 0.03 92.9 15.85 f 0.38 5.01 f 0.001 3.75 i 0.003 a Weight percent of the first named compound in parentheses. Temperature 30 i 0.2"C. Transition energy at 414 nm = 70.2 kcal mol-'. Donor concentration = 5.16 x M. TABLE 11: Solvent Controlled Thermodynamic Parameters of the Complex at 25"Ca -ASo, -AGO, -M, cal mol-' Solvent kcal mol-' kcal mol-' deg-' 3.0
Water Methanol-water (16.46) Ethanol-water (12.1 5) Propanol-water (8.20)
2.8
4.07 3.36
f
i
0.02 6.11 i: 0.085 6.84 i: 0.30 0.006 3.72 i 0.200 1.22 i 0.64
3.51 f 0.010 4.99
i
0.153 4.97
f
0.48
0.007 4.88
f
0.085 4.06
f
0.31
3.67
f
a Weight percent of the first named compound in parentheses. Donor concentration = 5.16 x M.
x 0,
0
2.6
2.4
2.2
I/o x 1 0 2
Flgure 2. Effect of solvent polarity on complex formation. Dependence of K on dielectric constants of solvent mixtures at five different temperatures: for 25, 35, 45, and 55' sets:,(o)water,.(P)l-propanol (8.2 wt %), (A)ethanol (12.15 wt %), ( 0 )methanol (16.46 wt %); for 30' set: (0)water, ( 0 )methanol (8.01, 16.46, and 25.21 wt %), (0)ethylene glycol (14.24 wt %), (A)ethanol (9.66 wt %), (0) 1-propanol (9.83wt %), (@) 2-propanol (12.15 wt %), (A)2-methyl2-propanol (12.06 wt %), (0)acetone (3.96 wt %), (0)dioxane (2.05wt %).
to evaluate the standard enthalpy and entropy changes. The reaction was spectrophotometrically studied at four different temperatures at 10' intervals in the range 2555'C. The equilibrium constants were determined applying eq 6, and log K vs. T-l plots were made to obtain AHo from the slopes. The thermodynamic parameters obtained at 25'C are given in Table 11. The standard states of the reactants were considered to be the hypothetical states of ideal solutions of unit molarity. As stated above the concentration effects of the proton donor on the derived K values were checked. Extrapolation of these values to zero
donor concentration in water and in 16.46% methanol showed a 5% decrease of the reported K values at 5.16 X M donor concentration. The effect of temperature may also be considered to be equivalent to the electrostatic effect, since the dielectric constant of the solvent changes (usually decreases) with increasing temperature. Plots of log K vs. D-' for the working temperatures for several mixed solvent systems are presented in Figure 3. For each solvent system a good linear variation was observed. The slopes of these lints were, on the other hand, different. The scope of specific solvent effects was further supported. Detailed Solvent Effects. For an understanding of the solvent effects in detail, identical experimental conditions were maintained, where to the aqueous samples having a fixed p-nitrosalicylic acid (3.26 X M ) and ethylenediamine (0.424 M )concentration, varied proportions of nonaqueous solvents were added, and the absorbances were measured at 414 nm. The fixed concentrations of the donor and the acceptor could thus give a direct estimate of the extent of the complex on a comparative basis of the solvent polarities and compositions. The results are elaborated in Figure 4. It is seen that at constant mole fraction, the efficiencies of the solvents are in the following order: dioxane < acetone < 2-methyl-2-propanol = dimethylformamide < 2-propanol < 1-propanol < ethanol < acetonitrile < methanol < glycol < urea. In this comparison formamide behaves rather peculiarly. The general trend of the solvent effects is almost in the order of increasing dielectric constant having notable variations with acetone, DMF, and FA. These findings encouraged us to measure the results at isoThe Journalof PhyskalCherni.stry,Vol. 80, No. 2, 1976
160
S. P. Moulik, S. Ray, and A. R. Das
O"
7 1
3.c
2.e
r cn
-
2.E
D 2 L
Figure 5. Absorbances of Figure 4 plotted against the dielectric constants of the media with their correspondlng signs.
2.:
I35
I
1.55
I45 I/D
1.65
x 102
Flgure 3. Effect of solvent polarity on complex formation. Dependence of K on temperature induced dielectric constant for the solvent media recorded in Figure 2 for the sets at 25, 35, 45, and 55O.
0.7
1
0.8
0.5
0.4
6 0.3
0.2
0.1
0
0.1
0.2
0.3
0.4
0.5
0.6
0.7
Mole f r a c t i o n
Figure 4. Influence of solvent compositions on the absorbance of the complex at fixed donor and acceptor concentrations of 3.26 x and 0.424 M, respectively: (H) urea, (0)ethylene glycol, ( 0 ) methanol, ( 0 )formamide, ( 0 )acetonitrile, (A)ethanol, ( 8 )2-propanol, (P) 1-propanol, (0) acetone, ( 0 )dimethylformamide, (A)2methyl-2-propanol, (0)dioxane. dielectric conditions. In Figure 5, absorbances are plotted against the dielectric constants for several of the solvents. It is seen that the order of efficiency is the reverse. Specific solvent effects other than electrostatic is further supported. The Journal of Physical Chemistry, Vol. 80, No. 2, 1976
Discussion The high stability of p-nitrosalicylic acid in releasing its phenolic hydrogen (pK = 10.37)8 in the absence of a strong proton-withdrawing agent can be reconciled through the formation of intramolecular hydrogen bonding. In the presence of the diamine this bond is weakened owing to a greater pull, and the released proton is lodged on the amine via a hydrogen bonded ion pair f ~ r m a t i o n .This ~ , ~ ion pair formation can be expected to depend on the polarity of the medium. That this electrostatic effect is not independent of the type of the solvents is evidenced from the nonlinear variation of log K with D-I when various mixed solvents are compared (Figure 2). On the other hand, linear variation of log K with D-I for a fixed solvent composition a t different temperatures warrants electrostatic effects to be a proper guide (Figure 3). In this context importance of a threshold polarity of the medium for the proton transfer complex should be recognized.1,2J0 A comparison of the added nonaqueous solvents on the basis of mole fraction shows that the order of complexation follows the polarity order of these solvents. This is not manifested when the comparison is made at isodielectric conditions (Figures 4 and 5). It is anticipated that the complex becomes stabilized through solvation by the water molecules." This is supported by the fact of more complex formation in water-dioxane medium than in water-methano1 medium at isodielectric conditions, since more water molecules are needed in the former than in the latter to yield equidielectric states. The same principle holds for other solvent systems. With this end in view, log A values a t 0.1, 0.2, 0.25, and 0.4 mole fractions have been plotted against D-' for many solvent systems in Figure 6 (log A values have been used to give a comparative basis with log K , A being directly proportional to K). Recognizing the probable uncertainties in the measurements the correlation is noteworthy. The necessity of water molecules for the etabilization of the complex by solvation may further be indicated by the maximum -ASo and -AH"values12J3 in this medium (Table 11).A sound comparison of thermodynamic parameters given in this table is restricted for want of a single reference state for different solvent systems. The results given in Figure 6 further show that the protic
161
Limitations of Mn(ll) Selective Broadening in NMR
i e
7
16
-
2o
ties. In this regard the pronounced effect of formamide at low concentrations is a special feature. At high concentrations it is less effective than either methanol or glycol (see Figure 4). All these results adjudicate a complex analysis when aprotic solvents are used. Further experiments are needed for a meaningful dissection.
-
I I
14-
12
-
Acknowledgment. The authors extend their thanks to Professor M. N. Das, Head of the Physical Chemistry Section, for laboratory facilities.
4
g -
10-
I
08
-
06
-
References and Notes
04-
02-
1 0
I 0.01
002
003
004
005
(1)S.P. Moulik, A. K. Chatterjee, and K. K. Sen Gupta, Spectrochim. Acta, Sect. A, 29,365 (1973). (2)S . P. Moulik, A. K. Chatterjee, and K. K. Sen Gupta, Ind. J. Chem., 12, 92 (1974). (3)G. Akerlof, J. Am. Chem. SOC.,54, 4125 (1932). (4)P. Rohdewald and M. Moldner. J. Phys. Chem., 77, 373 (1973). (5)J. Wyman, Jr., J. Am. Chem. SOC.,55, 41 16 (1933). (6)K. K. Kundu and K. Majumdar, J. Chem. SOC., Faraday Trans. 1. 69, 806 (1973). (7)E. H. Lane, S. D. Christian, and J. D. Childs, J. Am. Chem. SOC.,96,38 (1974). (8)J. L. Ernst and J. Manashi. Trans. faraday SOC..59, 230 (1963). (9)J. W. Smith, J. Chim.,Phys.,58, 182 (1964). (lp)R. M. Scott and S. N. Vinogradov, J. Phys. Chem., 73, 1890 (1969). (11) S. N. Vinogradov, R. A. Hudson, and R. M. Scott, Biochim. Biophys. Acta, 214, 6 (1970). (12)G. C. Pimentel and A. L. McClellan, "The Hydrogen Bond", W. H. Freeman, San Francisco. Calif., 1960. (13)E. M. Arnett, E. J. Mitchell, and T. S. S.R. Murty, J. Am. Chem. SOC., 96, 3875 (1974).
Limitations Concerning Use of Manganese(II) Selective Broadening in Nuclear Magnetic Resonance Spectroscopy for Determination of Ligand Binding Sites William G. Espersen and I?.Bruce Martin' ChemistryDepartment, University of Virginia, Charlotfesville, Virginia 2290 1 (Received July 3 1, 1975) Publication costs assisted by the National Science Foundation
From comparison of proton and carbon-13 spin-lattice and transverse relaxation times in the presence of Mn(II), it is concluded that in many cases the dipolar term is not the dominant contribut'or to line broadening. Hence selective broadening experiments to determine the site of Mn(I1) binding to small ligands based on an r6dependence between Mn(I1) and the affected nucleus are not generally applicable. Binding sites and distances may be estimated from selective TI measurements, but in this case it must be established that the predominant dipolar interaction contributing to relaxation is that between the paramagnetic ion and the affected nucleus and that other, closer interactions from unpaired spin density on the ligand do not contribute importantly.
Sites of Mn(I1) binding to molecules such as nucleic acid bases, nucleosides and their phosphates, histidine and derivatives, and other amino acids and peptides have often been characterized by selective broadening of resonance lines in nuclear magnetic resonance spectroscopy due to
hydrogen and carbon-13 atoms that are near to the presumed metal ion binding site. I t is the purpose of this paper to assess the validity of selective broadening results with Mn(1I) on a variety of small ligands. Previously we have shown that there are severe limitations to use of selecThe Journal of Physical Chemistry, VOI. 80,NO 2, 1976
