THEI\'-NETHYLPROPIONAMIDE-WATER SYSTEM
5'9
and OH radicals could explain the additional iodine. In other words, step 10 should be replaced by reaction l G , as suggested by Bates and Spence.la Such a reCH302 +HCHO
+ OH
(W
action has also been suggested by Norrish in his extensive work on the autooxidation of methane a t high temperatures20 and in the flash photolysis of methyl iodide a t 140°.21 However, other workers22have concluded that the dissociation of CH302 radicals according t o reaction 16 is insignificant up to 50". The absence of additional iodine a t [ C H d ] = 0.01 M (Table I) also indicates that reaction 16 may not be important in solution a t room temperature. The significant point, ho.wever, is the fact that while CH302radicals preferentially dissociate or disproportionate in the gas phase, their main reaction in solution in the presence of reactive anions is the formation of hydroperoxide. It would be interesting to study the chemistry of CK302 radicals in solution when no anions likely to react with them are formed. Photolysis of methyl iodide solutions in organic solvents has given ample evidence for the formation of methyl radicals and iodine atoms.23
However, this work has not been done in the presence of oxygen. It is quite possible that some scavenging of spurelectrons a t high methyl iodide concentration is responsible for the additional iodine observed. For degassed solutions of 0.1 ill CHJ we found that G(H2) = 0.3. I n conclusion, we have considered the possible reactions of CH302 and 1 2 - radicals and have proposed a mechanism consistent with our results. A kinetic evaluation of the reactions of the intermediates is quite complex. Direct observation of the transient-species spectra by the pulse radiolysis technique would be more convenient for determining the kinetics. (19) J. R. Bates and R. Spence, J . Amer. Chem. Soc., 53, 1689 (1931). (20) R. G. W.Norrish, Discussions Faraday Soc., 10, 269 (1951). (21) J. F. Mcltellar and R. G. W. Norrish, Proc. Roy. Soc., A263, 51 (1961). (22) (a) J. A. Gray, J . Chem. Soc., 3150 (1962); (b) W.J. Blaedel, R. A. Ogg, and A. A. Leighton, J . Amer. Chem. Soc., 64,2499 (1942). (23) R. F. Pottie, W. H. Hamill, and R. R. Williams, Jr., ibid., 80, 4224 (1958).
The N-Methylpropionamidewater System. Densities and Dielectric Constants at 20-40°1 by Thomas B. Hoover National Bureau of Standards, Washington, D . C. aO&%$
(Received May 24, 1968)
Densities and dielectric constants of the N-methylpropionamide (?jMP)-mater system were measured atj 20, 30, and 40'. The curve of excess molar volume of the system us. mole fraction of water was essentially independent of temperature and had a minimum of -1.32 cm3at 0.6 mol fraction water. The molar volume of transfer of NMP from the pure liquid t o a dilute solution in water had a minimum of -6.0 em3 at 0.9 mol fraction of water. The dielectric data for the system were analyzed in terms of correlation factors calculated on the basis that the square of the gas-phase moment and the high-frequency dielectric constant for the mixture were linear functions of the mole fractions of the components. The structural properties of the system are discussed with respect to the corresponding data for the mixtures of water with N-methylacetamide, ethanol, or acetic acid. NMP and water each exert a strong effect on the structure of the other.
Introduction The S-monosubstituted amides are of interest for their remarkably high dielectric constants, a property which has been attributed to extensive chainwise association through hydrogen bonding. Furthermore, the miscibility of these compounds with water and several nonpolar organic liquids permits the preparation of mixed solvents covering an extremely wide
range of dielectric constants. S-Methylpropionamide (NIP)-water mixtures have been employed i n this (1) Presented in part at the 154th National Meeting of the American Chemical Society, Chicago, Ill., Sept 1967. (2) (a) S. J. Bass, W. I. Nathan, R. M. Meighan, and R. H. Cole, J . Phvs. Chem., 68, 509 (1964); (b) R.-Y. Lin and W.Dannhauser, ibid., 67, 1805 (1963); (c) S. Mizushima, T. Simanouti, S. Nagakura, K. Kuratani, M.Tsuboi, H. Baba, and 0. Fujioka, J . Amer. Chem. Soc., 72, 3490 (1950). Volume 78, Number 1
January 1969
58 laboraiory in an investigation of solvent effects on the kinetics of the hydrolysis of acetaL3 The dielectric constant measurements obtained for that study have recently been extended and supplemented by density determinations. The structural properties of the mixed system XXPwater are of interest since each component is highly structured through hydrogen bonding, but in different ways. X a n y of the unique properties of water are determined by its three-dimensional network of bondi n g 4 while X\IP exhibits remarkably long chains of h e a r associations.2a Density and dielectric constant measurements were expected to shed some light on the configuration of the mixtures. Hovermale, Sears, and Plucltnett5 have used these properties to study the structure of binary mixtures of S-methylacetamide (XAIA) with water and aliphatic alcohols arid have interpreted the associations in terms of the apparent Debye polarizations of the mixtures.
Experimental Section The preparation and purification of XlIP have been describedS6 The material was reclaimed by repeated fractional distillatioris after having been used for conductivity nieaburements and had a specific conductance of 1.16 X lo-' ohm-l cm-l (30") ( Y Z ~ ~=D 1.4350 and f p = -30.9"). Gas chromatography on a Carbowax 20lI liquid phase showed 0.11% water and no propionic acid, while Karl Fischer titration of a 100-cm3 sample showed less than 0.01% water. Density measurements were made in a bicapillary pycnometer of approximately 20 em3 capacity. The pyconometer was calibrated with distilled water a t each temperature. The standard error of the calibration was 0.01%. Densities of the three most dilute solutions of S J I P in water were also measured in each of three 5O-cm3 single-capillary pycnometers at 30". The staiidard deviations of the densities of each of these solutions mere less than O.0lo&. Dielectric constants mere measured a t audiofrequenciea by the substitutiori method with a Jones-Dike conductivity bridge (Leeds and Xorthrup KO. 4666). The bridge had, in series with the X arm, a 14-ohm coil used to balance the resistance of the slide-wire fineadjustment circuit of the S arm. This coil was shorted so that only the lead resistance (less than 0.02 ohm) was in series with the reference capacitor. The inequality of the distributed capacitances of the ratio arms was adjusted to less than 2 pF. A calibrated General Radio Type 722-H capacitor and a fixed 1000ohm shunt were connected in parallel across the unlinown (X) arm. The shunt ensured that no changes in the bridge setting above the 100-ohm decade would have to be made as the cell was inserted or removed from the circuit. The difference between the bridge coil Capacitances for the two settings under these circumstances was assumed to change the balance condiThe Journal of Physical Chemistry
THOM'IAS B.
HOOVER
tion by less than 5 pF. All of the solutions having dielectric constants greater than 90 actually required adjustment only of the 10-ohm and lower decades. The maximum specific conductance of the solutions was 8.8 X low6ohm-' cm-l (0.884 mol fraction of water a t 40"). A two-terminal borosilicate glass cell with concentric, cylindrical platinum electrodes was used. The dielectric constant, E, was calculated from the observed capacitance, C, in picofarads, by the equation E
=
(C - a ) / b
(1)
The constants u and b and their standard errors were obtained by fitting the equivalent form
C=u+be
(2)
to the five data points for nitrobenzene ( 3 5 O ) , mater (ZOO), and NAIP (20, 30, and 40"). The reference values of these standards mere obtained with an absolute (three-terminal) cell and a General Radio Type 1615-A capacitance bridge. The constants of eq 2 and their standard errors were a = 2.7 rt 0.9 and b = 8.149 f 0.007. The precision of the calibration supports the assumption that variations in the bridge coil intercapacitances were not a serious source of error. AIeasuremenls were made at 1, 2 , 4, 8, and 16 1&z. The variation of the measured capacitance over the experimental frequency range was greatest for the water-rich mixtures but did not exceed 1%. The extrapolated capacitance did not differ from the value measured at 8 lrHz by more than 0.2% and the uncertainty in the correction for polarization is probably not greater than 0.27,. The indicated precision of the cell calibration was O . l % , and errors in the calibration of the reference capacitor or in temperature control were less than this. The over-all uncertainty is f1% based on a standard error of 0.2% and an allowance of =to.;% for bias. The density and dielectric constant values for pure W I P mere obtained as a part of a more extensive characterization program,' based on several additional samples of KMP. The densities obtained with the pycnometer of the present study were within +0.02% of the pooled results obtained with three quartz singlecapillary pycnometers and two dilatometers. Since S I I P was used as a calibrating fluid for the dielectric cell, values of E obtained with the three-terminal cell are reported in Table I. The dielectric constants reported (3) R. K. Xolford and R. G. Bates, J . Phys. Chem., 6 6 , 1496 (1962).
(4) (a) J. D. Bernal and R. H . Fowler, J . Chem. Phys., 1, 515 (1933); (b) L. Pauling in "Hydrogen Bonding," D. Hadzi, Ed., Pergainon Press Ltd., London, 1959, p 1. (5) R. A. Hovermale, P. G. Sears, and M ,K. Plucknett, J. Chem. E ~ QData, . 8 , 490 (1963). (6) T. B. Hoover, J . Phys. Chem., 6 8 , 876 (1964). (7) C. G. Malmberg and T. B. Hoover, unpublished results.
THEN-~IETHYLPROPIONAMIDE-WATER SYSTEM
59
Table I : Dielectric Constants and Densities of N-Rlethylpropionamide-Water Mixtures mol fraction of Hz0
0 0.0174 0.0433 0.0880 0.2779 0.3541 0.4909 0.7571 0.8842 0,9283 0.9513 0.9846 1.0000
P,
PI
t
185.3 178 171
g/cma
97 83 83
0.93449 0.9355 0.9364 0.9388 0.9501 0.9563 0,9665 0,9912 0,9992
80.10
0.99820
111
e
166.7 161 155
103
here for pure iYAIP are 9% higher than those of ref 2a8 and 1.5% higher a t 30” than those of Leader and Gormley.9 The temperature variation of E is in good agreement with ref 2a but is significantly greater than that of ref 9. Solutions were made by weight, with buoyancy corrections, and both the NMP and the solutions were stored and transferred under argon. The concentrations (given in mole fraction of water), densities, and dielectric constants are listed in Table I.
91
Pl
dcma
e
g/cma
0.92647 0.9275 0.9286 0.9309 0.9421 0.9473 0 9586 0.9834 0.9929 0.99325 0.99331 0.99360 0,99565
150.6 146 141
96 86 77 76
0.91847 0.9194 0.9205 0.9229 0.9341 0.9393 0.9503 0 9752 0.9860
73.15
0.99222
I
81 79
76.55
AVz =
-
?z
I
= V’/X~
(3
where V2is the molar volume of the pure organic component and a2is its apparent molar volume in the s o h tion a t mole fraction z2. Both amides, like ethanol, show a pronounced minimum in AV2 near 0.1 mol fraction. The interpretation of this behavior in the case of 0
Volumetric Behavior The volumetric behavior of the system NMP-water is shown as the excess molar volume, VE, in Figure 1. This quantity, values of which at 30” are plotted, is nearly independent of temperature over the range of observations. For comparison, the corresponding data for the systems water-NAlA, water-acetic acid, and water-ethanol are included. These examples were intended to represent a wide variation of possible organic solute-water interactions through hydrogen bonding, but within a small range of molar volumes. They include the next homolog of NIIP, the strongly dimerized acetic acid, and the (presumably) more waterlike ethanol. The phenomenon of contraction (negative VE) when a hydrogen-bonding liquid is added to water is quite general, but apparently only the poly(ethy1ene glycol) ethers’O show a greater negative excess volume than SAIP. Such a contraction is frequently inter?r.etedl1in terms of the two-structure model for water.12 i hi the basis of the density functions used by n’6methy h i d Scheraga,13 a shift in the equilibrium between lowdensity structured “icebergs” and more dense “monomer” water by the addition of XMP cannot account for the minimum VE for this system. I n the predominantly aqueous range of compositions it is helpful to consider the apparent volume change on transferring 1 mol of the organic component from the pure liquid to a dilute solution in water. I n Figure 2 this change, AV2, is shown for four polar solutes
-0.5
E
”>
-1.0
0
04 06 08 MOLE FRACTION ORGANIC COMPONENT
02
IO
Figure 1. Excess molar volume a t 3 0 ” : 0 , NMP, this work; -, NNIA;6 - - -, ethanol; - - acetic acid.
--
.,
(8) The low result seems to have been almost completely accounted for by adventitious water in the sample (R. H. Cole, private commiinication). (9) G. R. Leader and J. F. Gormley, J . Amer. Chem. Soc., 73, 5731 (1951). (10) W. J. Wallace, C. S,Shephard, and C. Underwood, J . Chem. Eng. Data, 13, 11 (1968), (11) F. Franks and D. J. G. Ives, Quarf. Rev., 20, 1 (1966). (12) (a) H. S. Frank and W-Y. Wen, Discussions Faraday SOC.,24, 133 (1957); (b) A thorough and recent review of theories of water structure is given by W. Drost-Hansen, “Equilibrium Concepts in Natural Water Systems,” Advances in Chemistry Series, No. 67, American Chemical Society, Washington, D. C., 1967. (13) G. NBmethy and H. A. Scheraga, J . Chem. Phys., 36, 3401 (1962). If the addition of a small amount of K M P converted all the water to the dense form, the maximum contraction would be 1.325 cma/mol at 30°. At 60 mol % water the calculated contraction is 0.6 X 1.325 = 0.795 cms/mol, while the experimental value is 1.32 cm3/mol.
Volume 73, Number 1
January 1960
60
THOMAS B. HOOVER I
I
I
I
I
I
-2
-
-3
2
%
-4
9 Q -5
-6
-7
stabilization occurs a t the concentrations corresponding to the minimum in the AV, curves. These concentrations, representing 6-9 mol of water/mol of organic solute, are consistent with the stoichiometry of known clathrate hydrates.15 I n the concentration range where the curves of AV2 for NMA and NAIP are nearly parallel, the vertical displacement is a measure of the chain length, or -CH2-, effect.13 The size of this effect is considerably smaller than the average value of -1.2 cm3/-CHzfound by Friedman and Scheraga‘o from measurements on a series of normal alcohols but is closer to the values found when the -CHz- was introduced between the hydroxyl groups of glycols.17
Dielectric Behavior I
0.I
I
I
0.2
0.3
I
0.4
I
0.5
MOLE FRACTION ORGANIC COMPONENT
Figure 2. Molar volume of transfer from organic liquid to aqueous solution at 30’; 0, NMP, this work; A, NMA;6 c-- , ethanol; - acetic acid. e
The dielectric constants of NJIP-water mixtures are portrayed in Figure 3 in terms of a deviation function, Ae, representing the deviation from an ideal volume mixture relation. As =
€12
-
$%el
-
(4)
pzez
a ,
the alcohols has been discussed by Franks and Ives.” The minimum in AVz results from competing effects on water due to the polar and nonpolar ends of the organic molecule. The hydroxyl end of the alcohol is structure breaking while the hydrocarbon end reinforces a lowdensity structure in the water. A similar interpretation probably applies as well to the amides. The apparent molar volume of transfer for acetic acid is characteristic of solutes, such as dioxane, having strong specific interactions with water, and reaches its extreme value in the infinitely dilute solution, That the minimum in AV2 for ethanol and both amides occurs at a finite concentration indicates a cooperative effect among the organic molecules. The contraction on adding the solute is greater when some is already present in the solution, within rather narrow limits. Although it is tempting to consider that selfassociation of the solute through hydrogen bonding may be reducing the structure-breaking effects of the polar groups upon water, the association constant of NMA in aqueous solution14 rules out such an explanation. At mole fractions less than 0.5, WMA is almost completely unassociated. A more satisfying interpretation lies in the cooperative, specific participation of the solute molecules with water to form a clathrate type of structure.’l As Franks and Ives” are careful to point out, it is not clear what structure is stabilized, since crystalline clathrate hydrates of these solutes are not known but evidently the organic molecules help to form or to stabilize “holes” in the water structure. These holes are able to accommodate, a t least partially, the solute molecules with a consequent economy of volume. On this basis, the maximum structural The Journal of Phgsical Chemistry
where cpl and 9 2 are the volume fractions of the components based on the densities of the pure liquids. It has been found that the dielectric constants of many polar mixtures can be represented as linear functions of the volume fraction,l* and the alcohol-water mixtures show only moderate positive deviations from this relati0n.l‘ Comparison on a volume fraction basis largely compensates for the “dipole-dilution” effect discussed by Pranks and Ives.ll The very large negative deviations shown in Figure 3 for the N I P system, then, are clear evidence of changes in the degree of alignment of dipoles with changing composition. I n the alcohol
0
0.2
04
0.6
08
10
VOLUME FRACTION ORGANIC COMPONENT
Figure 3. Dielectric constant deviations (NMP-water). (14) I. M. Klotz and J. S. Franzen, J . Amer. Chem. Soc., 84, 3461 (1962). (15) D. N. Clew, J . Phys. Chem., 66, 605 (1962). (16) M. E. Friedman and H. A. Scheraga, ibid., 69, 3795 (1965). (17) K. Nakanishi, N. Kato, and M. Maruyama, ibid., 71, 814 (1967). (18) D. Decroocq and J. C. Jungers, Compt. Rend., 2 5 2 , 1454 (1961).
THEK-J~ETHYLPROPIONAIVIIDE-WATER SYSTEM 6
I
I
I
I
I
5
4
9 3
2
I
0 )
I
I
0.2
0.4
I
0.6
I
0.8
I
I.o
M O L E FRACTION ORGANIC COMPONENT
Figure 4. Correlation parameters: 0, NMP (30°),this work; A,-NMA (30°);5- - 1, ethanol ( 2 5 " ) (J. H. Hall and H. 0. Phillips, West V u . Univ. Bull., Eng. Expt.Sta. Research Bull., 26, 26 (1934)); 0 , acetic acid (25');*5 0, acetic acid (Z5").24 systems the nearest analog is in the small dip shown by t-butyl alcohol containing 15 mol % water.lg The water-K-butylacetamide system investigated by Reynaud20 is qualitatively similar to water-NMP and the sharp decrease in E as small amounts of water were added to the amide was ascribed to the breaking of amide-amide hydrogen bonds by the water. The degree of orientation of dipoles in pure liquids is indicated by the Kirkwood correlation parameter, g, but there is no simple theoretical treatment that is valid for mixtures of polar molecules. (See Figure 4.) The variation of g with composition was estimated by an adaptation of the equation given by Colez1 (eq 5), where p is the permanent dipole moment, N / V is
61
sociation in the pure liquids, but differ sharply from ethanol and acetic acid. As the concentration of the organic component increases, the water structure is broken down and there is less cooperative interspecies association. Above 50 mol % amide the dielectric properties of the solution are dominated by the amide. The near constancy of g for mixtures of ethanol and water has been pointed out bef0re.l' The acetic acidwater system is difficult experimentally because of the high conductances encountered. Campbell and GieskesZ4were not able to carry their measurements of the dielectric constant above 45 mol % water. Borovikov and Fialk0v,~5using a resonance technique, reported values for the entire composition range. Their results suggest that acetic acid is very similar to ethanol in the essentially aqueous region. At more than 90 mol % acetic acid, g becomes less than 1, as a result of the antiparallel coordination of the dimers.'s
Conclusions With respect to the volumetric behavior in dilute solutions in water, ATMPresembles its homolog, XMA, qualitatively but shows more pronounced effects. At less than 0.1 mol fraction of NMP there is evidence for very strong reinforcement of the water structure around the organic molecules, leading to a deep minimum in AV,. At higher concentrations the greater structuremaking influence of N N P relative to NMA follows, at least qualitatively, the predictions based on the greater length of the hydrocarbon moiety. The dielectric data reveal a much closer resemblance between the two amides than between any other pair of the organic solutes. The two sets of evidence, density and dielectric constant, show quite clearly that NMP and water each have a strong effect on the structure of the other. I n the primarily aqueous solutions this effect is shown by increased density, while in the organic-rich solutions it is revealed by a sharp decrease in the dielectric constant.
Acknowledgment. The author is indebted to Professor H. s. Frank for his comments and general suggestions concerning the interpretation of these results. the number density (molecules per cubic centimeter), and eo and E, are the low- and high-frequency dielectric constant limits, respectively. I n the mixtures it was assumed that the appropriate values of p 2 and E, in eq 5 are linear functions of the mole fractions of the components. Gas-phase values were used for / L ,and ~ ~E, for water and acetic acid components were taken as 10% greater than the square of the refractive index;23 the values of E, used for the other components were the following: ethanol, 2.20;23NMA, 2.51j2a and NMP, 2.54.2& On the basis of the empirical g factors, the two amides resemble each other, differing mainly in the extent of as-
(19) A.
C. Brown and D. J. G. Ives, J. Chem. SOC.,1608 (1962).
(20) R. Reynaud, Compt. Rend., Ser. C,266, 489 (1968). (21) R. H. Cole, J . Chem. Phys., 27, 33 (1957). (22) R. D. Nelson, Jr., D. R. Lide, Jr., and A. A. .l/aryott, "Selected Values of Electric Dipole Moments for Molecules in the Gas Phase," National Bureau of Standards Report NSRDS-NBS10, U. S. Department of Commerce, Washington, D. C., Sept 1967. (23) W. Dannhauser and L. W. Bahe, J. Chem. Phys., 40, 3058 (1964). (24) A. N. Campbell and J. M. T. M. Gieskes, Can. J. Chem., 42, 1379 (1964). (25) Yu. Ya. Borovikov and Yu. Sa. Fialkov, Elektrokhimiya, 1, 1106 (1965). (26) C. P. Smyth and (1930).
H. E. Rogers, J, Amer. Chem. SOC.,52, 1824
Volume 73,Number 1 January 1969
