M. V. Kauigud and K. J . Patil
71 4
Volumetric and Isentropic Compressibility Behavior of Aqueous Amine Solutions. I M. V. Kaulgud* and K. J. Patil Department of Chemistry, Nagpur University, Nagpur-10, lndia
(Received September 14, 1973)
Partial molal volumes ( 7 2 ) and apparent molal adiabatic compressibilities (I&) at 20" have been obtained at low concentrations (0-25 mol %) for the following amines from measurements of density and sound velocity: MeNH2, EtNH2, n-PrNH2, n-BuNHz, (Me)2NH, (Et)2NH, Et(NH212, and BzNH2. The Vz(x2) curves exhibit minima for all amines but BzNH2, and are analogous to the behavior of alcoholwater systems. The & values are generally negative at low concentrations but assume positive values at higher concentrations. For MeNH2, (Me)2NH, and (possibly) EtNH2 the apparent adiabatic compressibility goes through a minimum. The results are shown to be consistent with the stabilizing influence of the amines on the water structure. The &(x2) curves indicate a possible distinction of the way amine molecules exert their stabilizing influence. The lower members seem to dissolve substitutionally, whereas the higher members occupy the cavities forcing water into an ordered arrangement (hydrophobic hydration).
Introduction In recent years there has been considerable interest in the studies in dilute aqueous solutions of nonelectrolytes and a special class of electrolytes (tetraalkylammonium halides) because these solutes exhibit abnormal behavior at low concentrations. Thus, the work of Glewl showed enhanced stabilization of water structure around the solute molecules in the form of clusters or cages for ethylene oxide-water system from partial molar volume and pmr studies. Anomalies in many thermodynamic properties of alcohol-water systems, including the behavior of partial molal volume, have been explained by Franks and Ives2 on the basis of flickering cluster model of Frank and Wen.3 Partial molal volumes of some tetraalkylammonium salts go through a minimum at a particular concentration indicating the enclosure of these ions in water cagesS4These conclusions have been confirmed by Kay, et a1.,5 who studied the variation of the viscosity B coefficient and its' temperature variation for these alkyl ions. Depending upon the size of the alkylammonium ion, the B values are large, positive, and decrease with increasing temperature. Thus, information available in dilute solutions of many nonelectrolytes indicates (Franks9 formation of clathrate hydrate-like structure in solution. In a recent communication7 from this laboratory, it was shown that sound velocity in some aqueous-aliphatic amine solutions go through a pronounced maximum at a certain low concentration, which is mainly governed by the geometry of the molecules. Measurements8 of the viscosity B coefficient for some of these amines in solution and their temperature variation showed a behavior, similar t o those of tetraalkylammonium ions in water, indicating that a similar ordering effect on water molecules must be exerted by the amine molecules as well. Additional evidence for this behavior of amines is derived from the work of Jeffrey, et a1.,9 who have obtained solid hydrates of some of the amines and have also established their structure by X-ray study. In order to throw more light on the interaction of ahphatic amines with water, we have undertaken a detailed study of partial molal volumes as well as apparent compressibility of the following amines in water at different concentrations and at 20": methyl-, ethyl-, dimethyl-, diThe Journal of Physical Chemistry, Vol. 78, No. 7, 1974
ethyl-, n-propyl-, and n-butylamine. To this list were added ethylenediamine and benzylamine, to see if the presence of an additional interacting -NH2 group or the presence of a benzene ring in the amine molecule has any significant influence on the interactions. Experimental Section Methylamine (Fluka, 40% in water), ethylamine (Fluka, 70% in water), and dimethylamine (Riedel-de-Haen, 40% in water) were directly used. The strength of these solutions were obtained by titrating them with standard hydrochloric acid solutions volumetrically as well as by pH titrations. Dilutions were made by adding a weighed amount of water to weighed amount of solutions and the concentrations were obtained in terms of mole fraction (x2) of amine. n-Butylamine (Fluka, purum), n-propylamine (Fluka, practical), and benzylamine (Riedel-de-Haen) were dried over potassium hydroxide pellets and distilled twice. Ethylenediamine (CP grade) was purified by the standard method. The refractive indices and densities of these purified liquids agreed well with literature values. All solutions were prepared fresh before experiment with double distilled water in stoppered conical flasks. The densities at constant temperature (20 f 0.02') were found by using a calibrated 10-ml density bottle suspended in a U-10 ultrathermostat. The densities are considered to be accurate to i=5 units in the fifth decimal place. Sound velocity measurements for these systems were reported p r e v i ~ u s l y . ~ Results and Discussion Results of densities and sound velocities measurements for all the amines studied between 0 and 100 mol % (except methyl-, ethyl-, and dimethylamine, where the measurements could not be done at higher concentrations) are available as supplementary material. The plots of density us. weight fraction of amine result in curves which are convex upward indicating a contraction in volume after mixing in all cases. From the density data, the apparent molal volumes (4") were calculated using the expression
Volumetric Behavior of Aqueous-Amine Solutions
71 5 c
where c is the concentration in molarity, M2 is the molecular weight of solute, and do and cl are the densities of solvent and solution, respectively. The corresponding partial us. molmolal volumes ( 7 2 ) were evaluated by plotting ality ( m ) , finding the slopes a&/am, and using the expression
62
-102
60
100
The estimated error in & and V2 at the lowest concentration is about k0.2 ml, but is much smaller a t higher concentrations. The apparent molal compressibility of solute were determined from
where p and PO are the adiabatic compressibilities of solution and solvent, respectively. The estimated error in & a t the lowest concentration is kl x cm2 dyn-1. In Figure 1 we have plotted V2 and the excess partial molal volume (inserts) V2E = V 2 - V2 (V2 = molar volume of the pure solute) us. the mole per cent of amine. Figure 2 shows the variation of the apparent molal adiabatic compressibility with concentration for all the amines. As the concentrations are sufficiently low and the temperature (20") not very high, the difference between the isothermal and adiabatic compressibility is thought to be small enough to be ignored (cf. results of & at 20" for n-PrOH of Alexander and Hilllo). Smooth extrapolations of the curves in Figures 1 and 2 yielded Vz0 and +ko, the corresponding property at infinite dilution. In view of the work of Franks and Smithll on the extrapolation of &, curves to zero concentration in order to obtain 7 2 0 values, the procedure followed by us might be questioned. In order to test the correctness of our procedure, the Vl(x2) values for methyl- and dimethylamine (where sufficient data points below the mimimum in V 2 are available) were examined for their dependence on xz2 and x z 3 ( x g = mole fraction of amine). It was found that V l ( x 2 ) exhibited a satisfactory dependence on xz2 rather than x23 indicating the absence of an inflection point in the V2(x2) curves,ll thus justifying smooth extrapolation. It was assumed that other amines also possess no inflection points in the V2(x2) curves. This assumption is not wholly unjustified as the difference in VzOfor the homologous amines, which represents the limiting partial molal volume of a -CH2 group, turn out in our case to be 17.1, 16, and 14.8 ml for the monoamines respectively ( c f . 'Table I) in fair agreement with the value of 15.0 ml obtained by Alexander12 from measurements on alcohols. From Figures 1 and 2 the following observations can be made. (1) V'2 and hence VzE ( = V 2 - V Z )go through a minimum in all the amines except benzylamine. (2) The concentration at minimum in and VzE decreases as the chain length of the amine molecule increases. (3) The limiting partial molal volumes are smaller than the molar volumes of the pure solutes suggesting loss in volume of amine in solution. Again the limiting partial molal excess volumes VzoE is governed by the chain length, being more negative for longer chains. Introduction of a second amine group in EtNH2 to form Et("& reduces I7z0E. Similar observations were also made for the partial molal volumes of n-alcohols and glycols by Alexander12 and Nakanishil3 and seem to be a common feature of monosubstituted f &
Mol '10 amine Figure 1. Partid molal volume ume (inserts) v ~ E(= 02 - ~
(02) and excess partial molal vol-
2 at ) 20" as a function of mole per cent of amine for the aqueous solutions of: (a) (the concentration axis for this frame is from 0 to 50 mol YO)ethylenediamine ( 0 ) and benzylamine (A,right-hand scale), (b) n-propylamine ( 0 ) and n-butylamine ( A ,right-hand scale), (c) ethylamine ( 0 )and diethylamine ( A ,right-hand scale), (d) methylamine ( 0 )and dimethylamine (A,right-hand scale). alkyl derivatives. (4)The apparent molal compressibility +k are generally negative at low concentration but become positive at higher concentration after passing through zero. The concentration at which & passes through zero is about the same as that at which V2 goes through minimum. Remarkably in methyl- and dimethylamine (and probably in ethylamine) the &'s are weakly positive at low concentration and undergo minima before assuming more positive values at higher concentrations. The magnitude of the limiting apparent molal compressibility f$ko also appears to be governed by chain length, becoming more negative for longer amines. For benzylamine the $k values are positive throughout the concentration range. ( 5 ) The slope aP2/dxz of the V z ( x 2 ) curves before minima is least for methyl amine and highest for n-butylamine, other amines having intermediate values. Slopes of apparent compressibility d1p~/ax2also show similar characteristic differences. In Table I are collected values of Vzo, 1720E, &O, 8 V 2 / 8x2, and a&/dx2 at 20" as also the viscosity B coefficients at 25"* for all the amines. Examination of Figures 1 and 2 shows at the outset that benzylamine behaves like a normal solute with both Vz and (positive) & increasing monotonously with concentration. The abnormal behavior of V2 and VZE for the amines is similar to those of alcohols2 and tetraalkylammonium halides4 indicating a similarity in the solute-solvent interactions in the case of these classes of solutes. The negative
The Journalof Physical Chemistry, Vol. 78, No. 7, 1974
71 6
M. V. Kaulgud and K. J. Patii
TABLE I Mole fraction of amine at minimum gy
Solute
Methylamine Ethylamine n-Propylamine n-But ylamine Dimethy lamine Diethylamine Ethylenediamine Benzylamine a
0.18 0.083 0.05 0.02 0.085 0.045. 0.22
(+KO)
mI
?so,
40.9 58.0 74.0 88.8 60.8 90.8 62.1 104.7
V$,
ml
-5.90" -8.05 -8.5 -10.1 -8.05
-12.1 -5.0 -4.2
x
1010, cmz dyn-1
+4.5 -2.5 -9.5 -16.0 $2.5 -10.0 -6.5 +8.5
3??/DX2,
ml
- 20 - 69 - 159 -163
- 102 - 153 - 30
(bgiilax?) x 10'0, cmzdyn-1
Viscosity B coefficient (at 2 5 ' ) , 44-1
48 200 357 1150 181 416
0.11 0.23 0.27 0.35 0.19 0.51
45
b
The molar volume (Vz)for these solutes have been obtained from the density data taken from ref 14. Positive;
1
I
-20 0
I
5
I
10 15 Mol % amine
I
20
-10
25
Figure 2. Apparent molal compressibility (fJk at 20" as a function of mole per cent amine for the aqueous solutions of: methyl-
amine ,).( ethylamine (m), n-propylamine ( O ) , n-butylamine (A),dimethylamine (0,right-hand scale), diethylamine ( V , right-hand scale), ethylenediamine ( 0 ,right-hand scale), benzylamine ( X ) . values of & for some amines at lower concentrations are similar to those for n-PrOH observed by Alexander and Hilllo at infinite dilution and a t temperatures lower than 30". This was interpreted by them as being due to the loss of structural compressibility of water on account of the increase in the population of four bonded water molecules in the vicinity of the solute molecules. Critical evaluation of all the available physicochemical data for alcohol-water systems led Franks and Ives2 to the conclusion that there must be reinforcement of water structure in the neighborhood of alcohol molecules in dilute solutions. Results of V ~ ( X ZI&,) , and viscosity B coefficients for amines8 also seem to be consistent with this view of structure-forming effect by the solute. Additional support for this view comes from the analysis of the entropies of hydration ( A s h ) values for some dialkylamines by Franks and Watwho have shown that A s h values are negative even after making allowance for the loss of rotational entropy on dissolving. Negative entropies of hydration must of necessity arise out of an ordering effect exerted by the amines on the neighboring water molecules. The relative magnitudes of VzoE, &, and the B coefficients throw interesting light on the differences. Thus VzoE for methyl to butyl alcohol are of the order of -2 to -6 ml,I6 the corresponding amines show higher values of The Journal of Physical Chemistry, Vol. 78, No. 7, 1974
-5.9 to -12 ml (Table I). Although no reliable estimate can be made for this quantity for alkylammonium ion, these are presumably much larger than for amines.17 Likewise, +ro for tetramethyl- to tetrabutylammonium ion shows maximum negative values of -9.1 to -25.5 x 10-10 cm2 dyn-1 (obtained by subtracting for Br- = +2 X from that of salt18), while those for amines are little less negative (Table I), and those for alcohols are least negative: e.g., (fJko for n-PrOH in water is -3 x 10-10 cm2 d y n - l l 0 and for EtOH -2 x cm2 dyn-l.19 A more negative 6 k o means a greater loss of structural compressibility of water implying a greater ordering effect by the solute on the solvent. Values of viscosity B coefficients for amines (Table I, column 8) are generally higher than for alcohols (0.087 to 0.3 M-I 20) but lower than those for tetraalkylammonium ions5 (-1.4 M - I for (n-C4H9)4N+). It can thus be inferred from these observations that amines have a stronger ordering effect on water structure than alcohols but less than the alkylammonium ions. The slopes - aV#/axz show revealing differences similar to those for alcohols in water,2,21 increasing with the hydrophobic character of the solute. Higher slopes mean that the interactions causing negative VzOE become dominate as more solute molecules dissolve in water. Following Franks, et u1.,10,22,23it appears to be reasonable to conclude that for n-PrNHz, n-BuNH2, and (Et)2NH the stronger solute-solute interactions are responsible for high slopes - a V 2 E / d ~ 2 .Conversely, the smaller slopes for other amines, especially MeNHz and EtNH2, would indicate that the solute-solute interactions are much weaker, which would amount to stronger solute-solvent interactions. Substitution of one more amine group in EtNHz to give Et( NH2)2 should cause stronger solute-solvent interaction lowering -aV2E/ax2, as is also observed. The slopes &&/ax2 are also revealing. n-BuNH2, nPrNH2, and (Et)2NH show very high values. If a strong solute-solute interaction for these solutes is accepted, rapid increase in (fJk to positive values can be understood as resulting from a superposition on the negative structural contribution (resulting from structural stabilization of water) a rapidly increasing positive contribution due to solute-solute interaction (pure amines possess much higher compressibility than water). The steeper the increase, the stronger the assumed interaction. Higher values of -dV2/ax2 should thus result in higher values of a(fJk/axz, which is acutally found to be the case (Table I). The concentration dependence of (fJk for MeNH2 and (Me)2NH is rather startling and appears to be the first cases of & showing extrema at low concentration^.^^ EtNHz also shows signs of a possible minimum, but the same cannot be established for want of sufficient data at
Volumetric Behavior of Aqueous-Amine Solutions
717
TABLE I1 Solute
Methylamine Ethylamine n-Propylamine n-Butylamine Dimethy lamine Diethylamine
Activity coefficient
0.186 0.494 2.07
7.70 0.505 4.70
Free energy of solution AG, cal/mol
- 1023 - 296
+431
+1158 - 250 1150
+
lower concentrations. A possible clue to understanding these features can be had from the activity coefficients and free energy of solution obtained by Christie and Crisp25 in dilute solutions of a number of amines at 25”, which are given in Table 11. It is noteworthy that only MeNH2, EtNH2, and which have activity coefficients less than unity (negative deviations from Raoult’s law) and negative free energies of solution (indicating affinity for water molecules), exhibit minima in &. Whereas other amines showing activity coefficients greater than unity have no minimum in & and, moreover, the & values are large and negative at low concentrations. It is helpful at this point to invoke the hypothesis of substitutional dissolution for the lower amines (MeNH2, EtNH2, and (Me)aNH) and interstitial dissolution for others (an idea proposed by Franks and Ives2 in their review). Accordingly, the interstitially dissolved amines can be thought of as occupying either the cavities existing in the open water structure, or else such suitable cavities are “created on demand” (Franks and Ives) to suit the size and shape of the solute. The latter case will lead to strengthening of water structure in the vicinity of solute molecules giving negative values of &. In substitutional dissolution, an amine molecule at infinite dilution occupies one of the so called “framework sites” (Franks and Ives), displacing one water molecule into the interstitial site (Frank and Wen). Formation of one or two hydrogen bonds with amines thus leads to breakdown of a few others formed originally by the oxygen atom of the displaced water molecule. This leads to an incipient breakdown of water structure leading to small positive values of &. Addition of further amine molecules can then be thought of as undergoing preferentially hydrogen bonding with the displaced water molecules, thus creating a fresh lattice site. This would decrease &. This decrease in & with increasing concentration may continue to a point, where accommodation of the solute on the framework site is compatible with the native structure of pure water. Any addition of solute beyond this point must lead to a gradual break down of the lattice leading to an increase in 4 k again. MeNHz which shows a broad flat minimum at a comparatively higher concentration happens to be the most compatible with the water structure because of its small size. The bulky with a sharp minimum at lower concentration appears to be less SO. The minimum in & for the longer EtNH2 molecules lies probably at concentration below 1 mol 70.Such an explanation based upon substitutional dissolution would be in conformity with the magnitude of activity coefficients,
smaller slopes for aYzE/ar2 and smaller values for viscosity B coefficients. In conclusion, we can say that the volumetric and compressibility properties of amines, which in dilute aqueous solution show a strong similarity to those of alcohols and other monofunctional nonelectrolytes, are consistent with amine molecules having a stabilizing influence on water structure. Results of apparent compressibility measurements further indicate a possible distinction of the way amine molecules exert their stabilizing influence. The lower members of the homologous series appear to dissolve predominantly substitutionally and thus strengthen the native water structure whereas the higher members dissolve by occupying cavities and forcing water into an ordered arrangement (hydrophobic hydration).
Acknowledgments. We thank Professor R. H. Sahasrabudhey for his keen interest in this work. One of us (K. J. P.) is thankful to the University Grants Commission, India, for the award of a Junior Research Fellowship. Supplementary Materials Available. Tables of concentrations, densities, sound velocities, partial molal volumes, and apparent molal isentropic compressibilities data for all the eight amines at 20” will appear following these pages in the microfilm edition of this volume of the journal. Photocopies of the supplementary material from this paper only or microfiche (105 x 148 mm, 24X reduction, negatives) containing all of the supplementary material for the papers in this issue may be obtained from the Journals Department, American Chemical Society, 1155 16th St., N.W., Washington, D. C. 20036. Remit check or money order for $3.00 for photocopy or $2.00 for microfiche, referring to code number JPC-74-714. References and Notes (1) D. N. Glew and N. S.Rath, Can. J. Chem., 45,3058 (1967). (2) F. Franks and D. J. G. Ives, Quart. Rev., Chem. SOC., 20, 1 (1966). (3) H. S. Frank and W. Y. Wen, Discuss. Faraday Soc., 24, 133 (1957). (4) W. Y. Wen and S. Saito. J. Phys. Chem., 68, 2639 (1964). (5) R. L. Kay, T. Vituccio, C. Zawoyski, and D. F. Evans, J. Phys. Chem., 70,2336 (1966), (6) F. Franks in “Physico-Chemical Processes in Mixed Aqueous Solvents,” F. Franks, Ed., Heinernann, London, 1967. (7) M. V. Kaulgud and K. J. Patil,Acustica, 28, 130 (1973). (8) R. I. Patel, K. J. Patil, and M. V. Kaulgud, Z. Phys. Chem. (Frankfurt a m Main), 86, 67 (1973). (9) R. K. McMullan, J. H. Jorden, and G. A. Jeffrey, J. Chem. Phys., 47, 1218 (1967). (10) D. M. Alexander and D.J. T. Hill, Aust. J. Chem., 18,605 (1965). (11) F. Franks and H. T. Smith, Trans. Faraday Soc., 64, 2962 (1968). (12) D. M. Alexander, J. Chem. Eng. Data, 4, 252 (1959). (13) K. Nakanishi, N. Kato, and M. Maruyarna, J. Phys. Chem., 71, 814 (1967). (14) J. Tirnrnermans, “Physico-Chemical Constants of Pure Organic Compounds,” Elsevier, Amsterdam, 1950. (15) F. Franks and B. Watson, Trans. FaradaySoc., 65, 2339 (1969). (16) L. Benjamine, J. Phys. Chem., 70,3731 (1966). (17) F. Franks and H. T. Smith, Trans. Faraday Soc., 63, 2589 (1967). (18) 8. E. Conway and R. E. Verrall, J. Phys. Chem., 70, 3953 (1966). (19) Obtained by analyzing the sound velocity data of R. Kuhnkies, Dissertation No. D-83. Technical University, Berlin, 1962. (20) T. HerskovitsandT. M. Kelly, J. Phys. Chem.. 77, 381 (1973). (21) K. Nakanishi, Bull. Chem. SOC.Jap., 33, 793 (1960). (22) F. Franks and M . A. J. Quickenden, Chem. Commun., 388 (1968). (23) F. Franks, M. A:J. Quickenden, D. S. Reid, and B. Watson, Trans. Faraday SOC., 66, 583 (1970). (24) We are grateful to Dr. F. Franks for drawing our attention to this point. (25) A. 0. Christie and D. J. Crisp, J. Appl. Chem., 17, 11 (1967).
The Journal of Physical Chemistry, Vol. 78, No. 7, 1974
