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The Journal of Physical Chemistty, Vol. 82, No. 11, 1978
TABLE IV: Change in Polarizability of Organic L i q u i d s with Change in Density a t 296.15 K a n d 0.1 M P a
101'(dn/ lO"(Bh/ - 1 0 3 ( d a / dP)/Pa-*
Compound
Benzene n-Hexane
n-Octane n-Decane Carbon
PwWl
(Coumou6) Pa-'" 52,3 66.5 52.8 46.9 52.8
tetrachloride 2-Butanone 44.2 Cyclohexane 50.8 a B = (Lam,m, cot e)-1.
dp)/
cm'lg
55.5 70.3 55.9 49.8 56.9
5.24 6.84 6.41 6.53 1.87
49.1 54.8
8.54 7.43
compound to 296.15 K. From eq 10 it was then possible to calculate da/dp for these compounds if n was known at 296.15 and 546.1 nm. It is not yet possible to measure dn/dP or da/dp directly in the ultracentrifuge with conventional interference optics for comparison with other interferometrically determined values. The reason is that if one tries to centrifuge a double sector cell with liquid in one sector and vacuum in the second sector, the fringes are lost because coherence is lost due to the sizeable change in refractive index between the two sectors. The advent of laserB interferometry which does not suffer from loss of coherence may permit high enough precision in fringe measurement, in the ultracentrifuge to observe the term in da/dp of eq 10. It would then be possible to generate true, low pressure, dn/dP data relatively simply and quickly for large numbers of transparent liquids, and to demonstrate the validity of the Lorentz-Lorenz formula in light scattering and other applications.
H. P. Hopkins, C. J. Alexander, and S. Z. All
References and Notes J. Dayantis, C . R. Acad. Sci. Paris, Ser. C , 267, 223 (1968). A. J. Richard, J. Giick, and R. Burkat, Anal. Biochem., 37, 378 (1970). K. S.Rogers, R. Burkat, and A. J. Richard, Can. J . Chem., 51, 1183 (1973). R. K. Burkat and A. J. Richard, J. Chem. Thermodyn.,7, 271 (1975). L. I.Epstein, P. Nixon, and A. J. Richard, Can. J. Chem., 51, 3309 (1973). 8. P. Sahli, H. Gager, and A. J. Richard, J. Chem. Thermodyn., 8, 179 (1976). V. Raman and K. S. Venkataraman, Proc. R. SOC.London, Ser. A , 171, 137 (1939). E. Reisier and H. J. Eisenberg, Chem. Phys., 43, 3875 (1965). D. J. Coumou, E. L. Mackor, and J. Hijmans, Trans. Faraday Soc., 60, 1539 (1984). R. M. Waxier, C. E. Welr, and H. W. Schamp, J . Res. Nafl. Bur. Stand. ( U . S . ) , Sect. A , 68, 489 (1964). L. S. Shraiber, N. V. Kartsev, G. N. Beiousova, and P. K. Ivanov, Zh. Fir. Khim., 48, 1044 (1974). R. Josephs and A. P. Minton, J. Phys. Chem., 75, 716 (1971). A. Michels, C. Micheis-Veraart, and A. Biji, Nature (London), 138, 509 (1936). A. Michels, J. DeBoer, and A. Biji, Physics, 4, 981 (1937). A. Michels and J. Hamers, Physica, 4, 995 (1937). J. L. Jasper, J . Phys. Chem. Ref. Data, 1, 841 (1972). D. M. Cowan, G. H. Jeffery, and A. I.Vogei, J . Chem. SOC.,171 (1940). M. Hennaut-Roland and M. Lek, Bull. SOC. Chim. Belg., 40, 177 (193 1). K. Owen, 0. R. Quayle, and W. J. Ciegg, J. Am. Chem. SOC.,64, 1294 (1942). A. I. Vogei, J. Chem. SOC.,810 (1948). B. J. Zwolinski, "Selected Values of Properties of Chemical Compounds", Thermodynamics Research Center, Texas A&M University, College Station, Tex. R. Ceuterick, Bull. SOC. Chim. Belg., 45, 545 (1936). R. C. Wiihoit and B. J. Zwolinski, J. Phys. Chem. Ref. Data, 2, suppl. 1 (1973). A. E. Lutskii and V. N. Soion'ko, Russ. J . Phys. Chem., 38, 778 (1964). R. C. Williams, Anal. Biochem., 48, 164 (1972). J. S. Rosen, J . Opt. SOC.Am., 37 932 (1947).
A Quantitative Comparison of the Relative Hydrogen Bonding Basicities and the Gas-Phase Proton Affinities of Substituted Pyridines Harry P. Hopklns, Jr,,* C. J. Alexander, and Syed Zakir Ali Department of Chemistfy, Georgia State Universlty, Atlanta, Georgia 30303 (Received November 3, 1977; Revised Manuscrlpt Received March 13, 1978)
The equilibrium constants for the association of phenol with 19 substituted pyridines have been measured at 25 "C via an infrared technique. In a comparison with the gas-phase proton affinities the log K H values for all the 3- and 4-substituted pyridines correlated with the PA values (correlation coefficient = 0.986), but several of the sterically hindered pyridines deviated markedly from this correlation line.
For some time it has been recognized1i2that the relative hydrogen bonding strengths of a closely related series of acids or bases are related at least qualitatively to the pK, values of the aqueous ionization equilibria. Since it is now possible to quantitatively determine the concentration of ions in the gas phase at equilibrium condition^,^^^ the gas-phase acidities and basicities of a wide variety of compounds are becoming available. The gas-phase enthalpy or Gibbs free-energy values for the ionization equilibria are a direct measure of the intrinsic acidity or basicity of a molecule, unincumbered by solvation effects. Thus, the gas-phase thermodynamic parameters might be 0022-365417812082-1268$01 .OOlO
expected to correlate much better with the hydrogen bonding data than the aqueous values. Aue, Webb, and Bowers have5>6established the quantitative gas-phase basicities for a variety of 2-, 3-, and 4-substituted pyridines and it is the purpose of this study to compare these data with a consistent set of hydrogen bonding data for the same pyridines. In a recent comparison between aqueous and gas-phase basicities for substituted pyridines: the gas-phase proton affinities (PA) were found to linearly correlate with the aqueous enthalpies of ionization and pK, values. For the 3- and 4-substituted pyridines a single proportionality 0 1978 American Chemical Society
The Journal of Physical Chemistry, Vol. 82, No. 11, 1978
Phenol-Pyridine Equilibrium Constants
constant was found to adequately correlate the data, whereas, a slightly different linear correlation was observed for the data for the 2-substituted pyridines. Since several w ~ r k e r s have ~ ~ - reported ~ approximately linear relationships between hydrogen bonding basicities and the corresponding aqueous pK, values for the substituted pyridines, the gas-phase P A values are also expected to correlate with the hydrogen bonding strengths. However, the pK,-hydrogen-bonding strength correlation was only moderate and covered only a limited number of substituted pyridines. According to Joris and Schleyer7,the correlation is nearly the same for 2-, 3-, and 4-substituted pyridines and "there appears to be no significant deviation in the pK, - log KH plot attributable to steric effects". With these observations in mind it seemed appropriate to directly compare the gas-phase PA values of the substituted pyridines with the hydrogen bonding basicities. This comparison will establish the extent to which PA values in a closely related series are related to hydrogen bonding strengths. It will also be possible to assess the influence of steric requirements on hydrogen bonding equilibria since steric requirements do not appreciably influence the PA values in the gas phase for substituted pyridines.6 The equilibrium constants for the hydrogen bonding of substituted pyridines reported in the literature vary widely because the experiments were performed with different acids, solvents, and temperatures. Several of the hydrogen bonding equilibrium constants were evaluated at acid concentrations where self-association of the acid may be occurring and are subject to large errors even when corrections are applied because of the ambiguity regarding the association models. Furthermore, the pyridines previously studied do not cover the wide range of PA values found for substituted pyridines in the gas phase. In order to have a self-consistent and extensive set of hydrogen bonding basicities, the equilibrium constants on a molar basis, KH,were determined for 19 substituted pyridines by an infrared procedure in CCll for the reaction base
+ phenol ?r? complex
(1)
Special precautions were employed to minimize the effects of water and the self-association of phenol on the equilibria.
Experimental Section All spectral data were recorded with a Beckman IR-12 spectrometer operated in the double beam mode using a 5-mm variable pathlength cell equiped with Teflon stoppers. The temperature of the cell was regulated to 25 f 1 "C by circulating thermostated water through a jacket attached to the cell. The temperature of the cell during the measurements was monitored with a copper-constantan thermocouple in direct contact with the window of the cell. Final absorbance data were recorded when the thermocouple reading was constant. Spectroscopic grade carbon tetrachloride was dried by storing it over freshly activated (600 "C) Linde 4A molecular sieves. Analytical reagent grade phenol was vacuum sublimed to remove traces of water. All the substituted pyridines were purchased from the Aldrich Chemical Co., except 2,6-di-tert-butylpyridine and 2-tert-butylpyridine which were purchased from the Chemical Samples Co. Each of the pyridines was purified by vacuum distillation or vacuum sublimation if it was a solid at 25 "C. After the vacuum line purification of the pyridine, the samples were transported directly to the nitrogen flushed drybox without exposure to the atmosphere. Stock solutions of phenol and the substituted pyridines were prepared directly in the drybox with the
1269
dried carbon tetrachloride. The final solutions were prepared by a volumetric dilution technique and placed in the IR cell in the drybox. The equilibrium concentration of phenol in the carbon tetrachloride solutions containing the bases was obtained by monitoring the absorbance of the band centered at 3613 cm-l assigned to the OH stretching frequency of moM the nomeric phenol. At concentrations below 6.5 X apparent extinction coefficient at 3613 cm-l was independent of concentration indicating that the monomer is the predominate species in these solutions. Since the pathlength of the cell could not be directly determined and the actual extinction coefficient is not needed in the equilibrium calculations, the pathlength was routinely held fixed during a particular set of experiments. Small variations (less than 9%) were observed in the apparent extinction coefficient from one run to another. The equilibrium constants, KH, were calculated for a 1:l model using the observed equilibrium concentration of free phenol calculated from the peak absorbance at 3613 cm-l along with the initial concentrations. All the bases studied have small absorbances in this region, therefore, a correction to the absorbance at 3613 cm-' was made due to the base. In each experimental determination, the absorbances of three solutions were used to obtain the concentration of the monomeric phenol at equilibrium. These included the absorbance (APh) for phenol at concentration coph, the absorbance of the mixture (Amix) containing the base (COB), and the absorbance of the base (AB) at the base concentration COB. The concentration of the phenol at equilibrium in the presence of the base was taken to be c p h = (Amix- AB)(Coph/Aph).
Results and Discussion The concentration and absorbance data employed in the calculation of the hydrogen bonding equilibrium constants for the substituted pyridines are listed in Table I. For 4-dimethylaminopyridine (CDMAP) the KH values were found to be dependent on the concentration of the base above 0.0035 M. Only KH values determined below this concentration are listed in Table I for this compound. The average KH values are uncertain by at least 5% due to the uncertainty in the absorbance measurements and the influence of trace quantities of water in these solution on the derived KH. The KH values derived from this study are within 15% of the available literature values, except for pyridine and 3-cyanopyridine. The literature data for pyridine vary from 48.5 to 29.3; the value derived from this study, 38.5, with a range of f1.0, is in good agreement with the value of 41.0 reported by Rao and Singhag Attempts were made to determine the maximum of the absorption due to the hydrogen bonded OH stretching mode of the complex. For the majority of the substituted pyridines this peak overlapped the aromatic CH absorption region of the phenol and the pyridine. Because of this overlap and the broad nature of the hydrogen bonded peak it was not possible to accurately locate the maxima for the hydrogen bonded complexes. In Figure 1the PA values6for the substituted pyridines are plotted vs. the hydrogen bonding Gibbs free energy change (-AG,). A linear correlation exists for all the 3and 4-substituted pyridines studied: slope = 9.84 and correlation coefficient = 0.986. The 2-methyl-, 2-fluoroand 2-cyanopyridinesfall within experimental error on this correlation line: slope = 9.72, correlation coefficient = 0.987. The variations in the PA values are attenuated by a factor of 10 in the hydrogen bonding equilibria, whereas the corresponding attenuation factor for the aqueous ionization comparision6 is -3. Thus, while the absolute
1270
The Journal of Physical Chemistry, Vol. 82, No. 11, 1978
H. P. Hopkins, C. J. Alexander, and S. 2.All
TABLE I: Equilibrium Data for the Complexation of Phenol with Substituted Pyridines in Carbon Tetrachloride at 25 C C'Phr
M
lo3
C'B,M
Av K H
Lit valuea
PAs16
4.67 4.67
3.06 6.12
0.575 0.575
0,283 0.184
AB 0.004 0.008
Kn
H
37.4 39.5
38.5
217.4
2-CH3
5.04 5.04 5.04 5.04 5.04 5.04 5.04 4.35 4.35 3.79 3.79 3.86 3.86 3.86 3.78 1.83 2.75 2.17 4.39 3.51 4.33 4.33 4.33 5.51 5.51 5.51 5.51 5.51 5.51 4.43 4.43 4.43 4.43 4.43 4.43 4.43 4.43 4.43 5.04 5.04 5.04 5.04 5.04 5.04 6.01 6.01 6.01 6.01 6.01 6.01 6.01 5.10 5.10 5.10 6.01 6.01 6.01
3.93 1.96 1.18 3.94 1.97 1.57 7.88 1.29 1.54 1.03 1.54 1.76 1.18 1.47 1.415 1.415 0.849 1.77 1.74 2.08 0.61 0.91 1.53 4.99 15.00 24.93 0.417 4.17 8.35 4.37 6.56 1.66 3.32 6.64 1.13 3.39 5.65 7.91 2.20 3.09 1.12 3.36 5.61 7.85 0.92 2.76 4.60 0.127 0.191 0.318 0.302 2.04 3.07 4.09 2.92 5.84 7.77
0.702 0.702 0.702 0.696 0.696 0.696 0.696 0.575 0.575 0.501 0.501 0.510 0.510 0.510 0.464 0.188 0.320 0.225 0.455 0.364 0.449 0.449 0.449 0.741 0.741 0.741 0.756 0.756 0.756 0.591 0.591 0.584 0.584 0.584 0.589 0.589 0.589 0.589 0.717 0.717 0.684 0.684 0.684 0.684 0.849 0.849 0.849 0.843 0.843 0.843 0.837 0.674 0.674 0.674 0.845 0.845 0.845
0.243 0.354 0.447 0.220 0.332 0.378 0.130 0.348 0.323 0.332 0.280 0.299 0.339 0.308 0.387 0.151 0.282 0.112 0.235 0.170 0.403 0.383 0.344 0.729 0.710 0.721 0.623 0.222 0.143 0.455 0.403 0.472 0.392 0.295 0.566 0.513 0.468 0.416 0.596 0.569 0.632 0.537 0.469 0.413 0.701 0.525 0.412 0.799 0.730 0.627 0.644 0.598 0.571 0.546 0.701 0.602 0.562
0.012 0.007 0.005 0.008 0.006 0.006 0.009 0.010 0.012 0.010 0.012 0.005 0.004 0.005 0.000 0.000 0.000 0.001 0.001 0.001 0.000 0.000 0.000 0.008 0.030 0.049 0.005 0.013 0.024 0.023 0.036 0,001 0.001 0.001 0,010 0.010 0.010 0.010 0.002 0.002 0.009 0.010 0.012 0.014 0.000 0.001 0.007 0.001 0.002 0.003 0.006 0.000 0.000 0.000 0.006 0.009 0.013
56.8 59.9 58.9 63.6 65.9 65.2 63.8 60.6 59.7 57.4 57.8 50.0 49.2 51.0 15.0 17.8 16.6 61.9 61.9 60.9 19.9 20.3 21.3 0.56 0.60 0.41 70.5 69.3 67.9 7.8 8.9 15.1 15.2 15.2 5.7 4.7 5.1 5.3 6.6 6.3 9.1 9.3 9.1 9.3 25.9 24.5 25.6 513 632 551 507 6.5 6.1 5.9 7.7 7.5 7.2
58.5
48.5 29.3 41.0 57.5
220.7
64.6
57.5
219.8
Substituent
3-CH,
4-CH3
2-CH2CH 3 2-t-BU 4-t-BU 2,4-Di-t-Bu 2,6-Di-t-Bu 2,6-DiMe 2-c1
341 2-CN
3-CN 4-CN
2-N(CH s 11 4-N(CH3)2
2-F 2-Me0
x
X
lo2
~~
~
a
220.6
50.1
221.9
16.5
(223)b
61.6
68.0
20.5
228.0
0.52 69.2
222.7 (228)b
80.5
224.1
8.4
211.8
15.2
212.7
5.2
205.9
6.5
12.5
206.5
9.2
10.6
207.6
226.3
25.3
232.7
551
6.2
208.8
7.5
218.3 12.5 57.5
3-Br
2-NH2 All values taken from ref 2, corrected to 25 "C.
58.9
213.3 220.8
Estimated.
magnitude of the substituent effect in the gas-phase protonation of 3- and 4-substituted pyridines is substantially greater than in the formation of the hydrogen bonded complexes, it appears that the same electronic factors are operative in both cases. This result is consistent with the conclusions of Arnett, Mitchell, and MurtylO derived from their comparison of hydrogen bonding enthalpies with enthalpies of protonation in fluorosulfuric acid. These workers found that a general correlation did not exist for all functional groups, but a well-defined
correlation line did exist for each functional group, including a good correlation for a limited number of substituted pyridines.ll For most of the 2-substituted pyridines the PA values do not appear to correlate with the hydrogen bonding Gibbs'free energies in the same way as the 3- and 4substituted pyridines; nor does a second linear correlation appear to exist just for 2-substituted pyridines as was found in the comparison with the aqueous ionization data. 2-Aminopyridine has a -AGH value within 100 cal of the
The Journal of Physical Chemistv, Vol. 82, No. 11, 1978 1271
Phenol-Pyridine Equilibrium Constants
a
P
*oo]
,
1.o
,
,
2.0 3.0 ~ G (kcal/mol) H
,
4.0
Figure 1. A plot of PA vs. the relative hydrogen bonding basicity of substituted pyridines. The open circles are the compounds included in the least-squares treatment to obtain the straight line.
line but 2-ethyl-, 2-tert-butyl-, and 2,6-dimethylpyridine are off the line by 250, 1000, and 600 cal, respectively. 2-Fluoropyridine, 2-chloropyridine, and 2-bromopyridine are to the left of the line by approximately 150,400, and 500 cal, respectively. Much larger deviations are found for 2-dimethylamino- (1100 cal), 2,4-di-tert-butyl- (1300 cal), 2,6-di-tert-butyl- (2800 cal, not shown in Figure l),and 2-methoxypyridine (1000 cal). These deviations from the line by the 2-substituted pyridines appear to be qualititively related to the steric requirements of the substituted pyridines in the immediate vicinity of the lone pair of electrons on nitrogen. When a substituent is located at the 2 position, the size of the substituent can conceivably adversely influence the approach of the proton to the basic site because of the space requirements of the phenol molecule and the substituent. From space filling models for the 2-substituted pyridine-phenol complexes, it is also apparent that the internal rotational motions of the complex are hindered relative to the 3- and 4-substituted pyridine-phenol complexes. For example, in the phenol-2,6-di-tert-butylpyridine complex the rotational motion of the phenyl ring is severely hindered when the OH proton of phenol is a t an appropriate hydrogen bonding distance from the nitrogen. This hinderance to internal rotational motion increases with the size of the substituent and appears to be substantial when methyl or tert-butyl groups occupy both the 2 and 6 positions. Both of these phenomena are steric effects that lead to smaller KH values than predicted from the gas-phase protonation equilibria. In the gas-phase studiess it has been shown that even in 2,6-di-tert-butylpyridinethere is no evidence to support the concept of steric hinderance to protonation. However, the deviation of this compound from the linear correlation of Figure 1 is too large to be shown (-3 kcal mol-I). Thus the deviations from the linear correlation shown in Figure 1 must be primarily related to the steric requirements in the hydrogen bonded complex. In this respect it is interesting to note that 2,4-di-tert-butylpyridineis nearly 5 kcal mol-l more basic in the gas phase than 2-tert-butylpyridine, but the KH values for these two compounds are virtually identical. Apparently, the tert-butyl group restricts the approach of
the proton in the phenol complex so that the increase in basicities as measured by the PA values is not evident in the hydrogen bonding basicity. It should also be noted that for pyridines containing basic sites other than the ring nitrogen, it is possible that two different hydrogen bonded complexes are formed in solution. When this occurs, the measured KH is the sum of the two equilibrium constants and two hydrogen bonded hydroxyl stretching frequencies should appear in the infrared spectrum. For the cyanopyridines two broad, but well-defined, bands were observed in the 3100-3600-~m-~ region. The maxima for the first band appeared 110-140 cm-l below the 3600-cm-l peak, whereas the second band was 300-360 cm-l below the free phenol OH peak. The peaks found at the lower wavenumber values are in the same region as the hydrogen bonded OH stretching frequencies for other pyridines and can be assigned to a complex where the hydrogen bonding is at the pyridine nitrogen. In hydrogen bonding studies on CC1, solutions of benzonitrile and phenol the KH and OH stretching frequency have been found to be 3.4 (L/mol) and 156 cm-l below the free phenol peak, respectively.2 Therefore, the broad peaks appearing 110-140 cm-l below the free phenol peak in the solutions containing the cyanopyridines are readily assigned to a complex where the hydrogen bonding occurs at the nitrile nitrogen. From the available hydrogen bonding data on nitriles? the basicity of the nitrile nitrogen in the cyanopyridines is expected to be similar to that of benzonitrile, i.e., KH 3.0 (L/mol). With this estimate the values in Table I can be corrected to provide approximate KHvalues for the pyridine nitrogen complexes. This correction moves the 3- and 4-cyanopyridines to the left by -0.3 and 0.2 kcal, respectively, but does not appreciably affect the correlation shown in Figure,l. For 2-cyanopyridine this correction places the corresponding point nearly 0.6 kcal to the left of the correlation line, which is consistent with deviations of the other 2-substituted pyridines from the correlation line. An alternative interpretation of the results for the cyanopyridines is that both nitrogens can be protonated in the gas phase. This is unlikely since the PA for benzonitrile can be estimated from the data of Beauchamp, Staley, and KleckeP as -190 kcal/mol which when compared to the cyanopyridines on the same PA scale,3 gives a difference in PA values of -17 kcal/mol. This means that the protonation at the pyridine nitrogen is favored by a factor of -10I2 and would occur almost exclusively in the gas phase.
-
References and Notes G. C. Pimental and A. L. McClellan, “The Hydrogen Bond”, W. H. Freeman, San Francisco, Cailf., 1960. M. D. Joesten and L. J. Schaad, “Hydrogen Bonding”, Marcel Dekker, New York, N.Y., 1974. M. T. Bowers, D. H. Aue, H. M. Webb, and R. T. McIver, Jr., J . Am. Chem. Soc., 93, 4314 (1971). R. W. Taft, “Gas Phase Proton Transfer Equillbrii” in “Proton Transfer Reactions”, E. F. Caklin and V. Gold, Ed., Chapman and Hall, London, 1976. D. H. Aue, H. M. Webb, and M. T. Bowers, private communication. D. H. Aue, H. M. Webb, M. T. Bowers, C. L. Liotta, C. J. Alexander, and H. P. Hopkins, Jr., J . Am. Chem. Soc., 98, 854 (1976). (a) L. Joris and P. v. R. Schleyer, Tetrahedron,24, 5991 (1968); (b) T. 0. Gramstad, Acta Chem. Scand., 16, 807 (1962); (c) A. M. Dlerchx, P. Huyshens, and T. Zeegen-Huyshens, J . Chim. Phys., 82, 336 (1965); (d) D. Neerinck and L. Lamberts, Boil. Soc. Chim. Be@., 75, 473, 484 (1966); (e) A. Halleux, ibid., 88, 381 (1959). D. Bostwick, H. F. Henneike, and H. P. Hopkins, Jr., J . Am. Chem. Soc., 97, 1505 (1975). S. Singh and C. N. R. Rao, Can. J . Chem., 44, 2611 (1966). E. M. Arnett, E. J. Mitchell, and T. S. S. R. Mum, J . Am. Chem. Soc., 98, 3875 (1974). A referee has pointed out that a llmlted and less precise quantitative comparison of relative hydrogen bonding basicities and PA (AH,(g))
1272
The Journal of Physical Chemistry, Vol. 82,No. 11, 1978 values for 3- and 4-substituted pyridines has very recently appeared in the literature. (E. M. Arnett, B. Chawla, L. Bell, M. Taagepera, W. J. Hehre, and R. W. Taft, J. Am. Chem. Soc., 99, 5729 (1977). This comparison is based on the hydrogen bonding of p-fluorophenol to nlne substituted pyridines, with 6AGm/6AHdg) ratios ranging from 0.05 to 0.14 and an average ratlo of 0.10 0.03. These results
*
Surender K. Jain are in remarkable agreement with the slope found in Figure 1, i.e., a 6AHB/6AH,(g)ratio of 1/9.84. (12) R. H. Staley, J. E. Kleckner, and J. L. Beauchamp, J. Am. Chem. Soc., 98, 2081 (1976). (13) D. H. Aue, H. M. Webb, and M. T. Bowers, J . Am. Chem. SOC., 98, 318 (1976).
Volumetric, Surface, and Flow Properties of Mixed Molten Hydrates. 1. Cr(N03)3*9H204- Ca(N03)2*4H20and Cr(N03)3*9H20 Cd(N03)2*4H20Mixtures
+
Surender K. Jaln Hindu College, University of Delhi, Delhi-110007, India (Received August 16, 1977; Revised Manuscript Received January IO, 1976)
Densities, surface tensions, and viscosities of Cr(N03)3,9H20+ Ca(N03)2.4Hz0and Cr(N03)3-9H20+ Cd(N03)2*4H20systems have been measured as a function of temperature and composition. Both systems show departure from ideal behavior in their volumetric, surface, and flow properties. The results suggest the existence of hydration-dehydration phenomenon in these mixed hydrated melts. The fluidity data were fitted to the equations, 4 = A,T1/2 exp[-B,/(T- To,,)]and 4 = A; expi-B,'/(V- Vo,,)l based on free volume model. A,, B,, To,,,A+', B,', and Voi, are the empirical parameters. The composition dependences of these parameters have been discussed. The intrinsic volumes, Vo)s,for both systems show positive deviations to an extent of is consistent 1-2.5% from additive values. A lower VOvalue for Cd(NO3)2-4Hz0relative to that for Ca(N03)2.4Hz0 with the smaller radius of Cd2+(0.97 A) relative to that of Ca2+(0.99 A).
Introduction Fused salt hydrates have been discussed as analogues of fused salts with large weak field cations of the type M(H20),"; this has been evidenced by spectral,lp2 cond ~ c t a n c eand , ~ ~volumetric6-8 ~ studies. However, the assumption that in hydrated melts the coordinated water exists as an integral part of the hydrated cations appears to be untenable due to the fact that there is no stoichiometry at which the transference number of water relative to the cation is zero and in view of the competition between hydration and ion association during the formation of associated species. Volumetric studiesgon melts containing the hexahydrate of nickel nitrate and the tetrahydrate of cadmium nitrate show pronounced negative deviations from additive values. Discontinuities in volume-mole fraction plots have been interpreted in terms of hydration-dehydration phenomenon in such systems. Similar but relatively smaller negative deviations (0.2-0.370) from the additive volumes in the system Ca(N03)2.4.09H20+ Cd(N03)z.4.07H20have also been reported by Moynihan et al.1° Recent investigationsl1J2 on a number of Lewis acid salt hydrates of metals belonging to the 3d series of transition elements reveal that the volume per equivalent of nitrate ion depends upon the number of moles of water per equivalent of cationic charge, irrespective of the charge and radius of the cation. The surface properties of these salts in the moltdn state were found12to be governed by the electrostricted dipoles of the coordinated water. This communication reports some of the observations made on systems containing hydrates of trivalent chromium nitrate and divalent calcium and cadmium nitrates. Experimental Section The source and quality of the salts used in this study are as follows: Ca(N03)2.4H20,BDH (India), AnalaR; Cd(N03)z.4Hz0,E. Merck, Extrapure; Cr(N03)3-9H20, 0022-3654/78/2082-1272$01 .OO/O
Ortanal (Italy), Analytical Reagent. The melting temperatures of these salts as determined by the cooling curve method were found to be 42.5, 59.5, and 36.8 "C; the corresponding literat~re'~ liquidus temperatures were 42.7, 59.4, and 37 "C, respectively. Agreement between the experimentally determined melting temperatures and those reported in the literature leads one to believe that the salts were of stoichiometric composition as reported by the manufactures. This, however, was verified for the tetrahydrates of calcium and cadmium nitrates gravimetrically, by vacuum dehydration of samples (ca. 5 g) at temperatures starting from near melting to about 130-140 "C. Repeated cross checks established the water content of the calcium and cadmium salts within fO.O1 of the stoichiometric value of 4. Each mixture was prepared separately by melting the requisite amount of the salts in hard glass flasks (ca. 100 cm3) fitted with an air-tight ground glass joint at the top and filtered in situ through glass filters (porosity G-3) under a slight positive pressure of dry, C02-free air. Filtered melts were then maintained at about 60-70 "C for about 1 h for maturing. Densities were measured by measuring the volumes of a certain amount of melt in a precalibrated densitometer capable of reading to 0.01 cm3. Details regarding the design, calibration of the densitometer, and the measuring technique have been described earlier.6 The estimated accuracy of experimental densities is f0.1%. Viscosities were measured with Cannon-Fenske type viscometers, calibrated with aqueous solutions of sucrose and glycerol; the viscometer constants were 0.2009 and 0.2647 CPs-l. The inherent accuracy of these measurements is estimated as &0.5%. Surface tensions were measured by the differential capillary rise method as described earlier.12 Under conditions of thermal equilibrium, the temperature did not vary by more than f O . l "C, over the region of interest. The 0 1978 American Chemical Society
