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MOHAMMAD ABU-HAMDIYYAH
The Effect of Urea on the Structure of Water and Hydrophobic Bonding’
by Mohammad Abu-Hamdiyyah Departmat of C h e m d r y , University of Southern California, Los Angeles, California 90007 (Received March 6, 1966)
Nonpolar solutes, especially hydrocarbons and hydrocarbon moieties, dissolve interstitially in water and in aqueous urea solutions. Physicochemical properties of aqueous urea solutions indicate that urea does not influence the structure of water the way ions do, but actively participates in the formation of clusters. Weakening of hydrophobic bonding upon addition of urea is due to the increased ease of cluster formation. However, since urea and water molecules have different geometries, strengthening of hydrophobic bonding occurs in the case of methane and ethane (and presumably the rare gases) at low temperatures because of the more exacting geometrical requirements for the formation of the interstices that accommodate these molecules.
Introduction Hydrophobic bonding is the term coined to describe the phenomenon in which nonpolar residues or moieties (usually hydrocarbon) in the presence of water associate together to form intramolecular aggregates, as in the native protein solutions, intermolecular aggregates such as micelles in surfactant solutions, or, in the extreme case, the precipitation of a separate phase, ie., the limited solubility in water, as is the case for hydrocarbons. 2-6 Schematically the process may be represented by nA=
A,
where A is the nonpolar residue and A,, the aggregate of the residues. This association has been explained in the past by energetic (enthalpic) considerations alone. Nowadays, however, it is believed that entropy plays a dominant role in this p r o c e s ~especially ~ ~ ~ ~ ~for not very long chain compounds’ and at low temperatures, and furthermore that the process is largely controlled by the water structure. Recently, several authors have called attention to the effect of urea upon various facets of hydrophobic bonding and the purpose of this paper is to arrive at a coherent interpretation of the available evidence. The structure of water has been discussed by many authors and up-to-date reviews of the present situation of water structure theories exist.8-12 All the modern theories are refined versions of the model proposed by Bernal and Fowler,13the main feature of which is the The Journal of Physical Chemistry
tetrahedral hydrogen bonding of water molecules to each other in three dimensions, giving rise to an extensive open structure. Whether the hydrogen bonding involves all the molecules in the liquid at a given time (the uniformist approach) 14-16 or whether only part of the molecules are hydrogen bonded, the rest being nonbonded monomers (the mixture approach),17-19 is im(1) Based on a proposition submitted in partial fulfilment of the r e quirements for the Ph.D. degree at the University of Southern California. (2) W. Kauzmann, Advan. Protein Chem., 14, 1 (1959). (3) G. NBmethy and H. A. Scheraga, J. Chem. Phys., 36, 3401 (1962). (4) A. Wishnia, J . Phys. Chem., 67, 2079 (1963). (5) E. S.Hand and T. Cohen, J. Am. Chem. SOC.,87,133 (1965). (6) E. D. Goddard, C. A. J. Hove, and G. C. Benson, J . Phys. Chem., 61,593 (1957). (7) L.Benjamin, %%id., 68, 3575 (1964). (8) F. 8. Feates and D. G. Ives, J . Chem. SOC.,2798 (1956). (9) G. NBmethy and H. A. Scheraga, J . Chem. Phys., 36,3382 (1962). (10) 0.Ya. Samoilov, Zh. Struld. Khim., 4, 459 (1963). (11) H. S. Frank in “The Proceedings of the Conference on Desalination Research,” Publication 942 of National Academy of Sciences, National Research Council, Washington, D. C., 1963,p. 141. (12) R. P.Marchi and H. Eyring, J . Phys. Chem., 68, 221 (1964). (13) J. D. Bernal and R. Fowler, J. Chem. Phys., 1, 515 (1933). (14) J. A. Pople, Proc. Roy. SOC. (London), A205, 163 (1951). (15) D. N. Glew, J. Phys. Chem., 66, 605 (1962). (16) J. A. Barker, Ann. Rm. Phys. Chem., 14, 245 (1963). (17) L.Hall, Phys. Reu., 73,775 (1948). (18) G. Wada, Bull. Chem. SOC.Japan, 34, 955 (1961). (19) L. Pauling in “Hydrogen Bonding,” D. Hadsi, Ed., Pergamon Press Ltd., London, 1959,p. 1.
EFFECT OF UREAON THE STRUCTURE OF WATERAND HYDROPHOBIC BONDING
material for our purposes since both approaches agree on the formation of interstices and cavities as a result of the hydrogen-bonded structures which are essential for our treatment. Therefore we may use the language of either approach without any loss. When nonpolar solutes dissolve in water, they increase the mutual ordering of water molecules around them as shown by thermodynamic studies.6t20 Thus the idea of “icebergs” being formed around nonpolar solutes was born, 2o and such solutes have been called structure formers.’l The effect of urea on the strength of the hydrophobic bond21-28has recently been investigated and, according to all recent studies,21-n the effect of urea is to reduce the strength of the hydrophobic bonding in all cases except for CH, and C2Hs molecules and hence CH3 and C2Hs moieties at low temperatures where the reverse is true.24p25No explanations were offered for this exceptional strengthening of hydrophobic bonding by urea. The reasons given for the usual weakening of hydrophobic bonding are varied and sometimes clearly confused. It was suggested that the weakening of hydrophobic bonding may be due to the favorable interaction of urea with the ‘(iceberg” region or the formation of urea-hydrocarbon clathrate-like aggregates in solution,21or may be due to the reduction of the cooperative structure of water that is responsible for the solvent structure effects, i.e., to the destruction of the “icebergs” which form around the hydrocarbon m o i e t i e ~ . ~Others21 ~ ? ~ ~ suggested that the increased solubility of hydrocarbons in aqueous solutions may either be due to the facilitation of hydrocarbon solvation by water in the presence of urea or to the solvation of the hydrocarbon by both urea and water molecules. Recently it was hintedz5126 that ordered structures may be formed around the solute by both urea and water molecules. Nevertheless, it was concluded26 that urea behaves as a structure breaker. On the other hand it was concludedz8that urea disrupts the structure of water and solubilizes the hydrocarbons since “the solubility properties of hydrocarbons in water result in large part from an ordering of the solvent.” More specifically, the role of urea was likened to that water moleof an i ~ n i e ,. , orienting ~ ~ ~ neighboring ~ ~ cules, restricting their participation in hydrogenbonded water clusters and leading to a region of disorder around the solvated solute molecule, thus “ions and urea may modify the iceberg structure around the hydrocarbon chains of the single ions and consequently affect micelle formation of ionic detergent^."^^ In this paper the mechanism of the dissolution of hydrocarbon solutes in water is first reviewed and then an attempt is made to analyze the way in which urea
2721
affects the structure of water and hence hydrophobic bonding. It will be pointed out (i) that hydrocarbon solutes always dissolve interstitially in pure water as well as in aqueous urea solutions, (ii) that urea is able to participate in cluster formation with water molecules, and (iii) that the formation of interstices to accommodate relatively large hydrocarbon moieties becomes easier in aqueous urea solutions than in pure water, but the converse holds for small interstices that would accommodate short-chain compounds (CHI and C2Ha) at low temperatures. In order to avoid confusion and because the word “structure forming” is used in discussing the role of ions,29of nonpolar and of urea (in this paper), it is worthwhile to point out the differences between their effects. The structure-forming ions destroy the hydrogen-bonded structures of water and draw the water molecule to themselves to form a compact structure (of water molecules around the ions). 11,3o The nonpolar solutes shift the thermodynamic equilibrium between the molecules of liquid water towards the open (icelike) structure components. 31 The structure-forming properties of urea differ from both kinds of solutes just described by the fact that (as will be shown later) it takes active part in the formation of the open structure in the solution (i.e., the formation of clusters). Thus, urea contributes to the open structure in the solution by the same mechanism as water molecules, instead of filling interstices in the open structure like nonpolar solutes, or of substituting for water molecules in the “lattice” sites and drawing the water molecule to itself to form a new different structure at the expense of the original one like a structure-forming ion. z9t
Discussion The Solubility of Nonpolar Solutes in Strongly Polar Solvents. In the condensed state two extreme types (20) H. S. Frank and M. W. Evans, J. C h m . Phys., 13, 507 (1945). (21) W. B-ing and A. Holtzer, J. Am. Chem. Soc., 83,4865 (1961). (22) P.Mukerjee and A. Ray, J . Phys. Chem., 67, 190 (1963). (23) P.Mukerjee and A. I(.Ghoah, &id., 67, 193 (1963). (24) D. B. Wetlaufer, S. I(. Malik, L. Stoller, and R. L. Coffin, J . Am. C h m . Soc., 86, 508 (1964). (25) Y. Nozaki and C. Tanford, J. Biol. Chem., 238, 4074 (1963). (26) G. C. Kresheck and L. Benjamin, J. Phys. C h m . , 68, 2476 (1964). (27) M.J. Schick, ibid., 68, 3585 (1964). (28) J. A. Rupiey, ibid., 68, 2002 (1964). (29) H.S. Frank and W. Y. Wen, Discussions Faraday SOC.,24, 133 (1957). (30) I. G. Mikhanilov and Yu. P. Syrnikov, Zh. Strukt. Khim., 1, 12 (1960). (31) V. I. Yashkichev and 0. Ya. Samoilov, %%id., 3, 195 (1962).
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of solutions are possible: (a) substitutional, where the solute particle occupies a “lattice” site normally occupied by a solvent molecule and thus competes with the solvent molecules for the sites, as in. the case of substances whose molecules have comparable sizes and force fields; and (b) interstitial, where the solute particle occupies a cavity in the structure of the solvent and thus does not compete with the solvent molecules for the “lattice” sites, as in the case of nonpolar solutes and strongly polar solvents.3 2 Ample support for this comes from thermodynamic calculations (e.g., Frank and Evans20 and D’Orazio and Woods3) and recently from nuclear magnetic relaxation measurements in aqueous and nonaqueous solutions of a large number of substances containing alkyl groups.34 The results of the last reference clearly indicate the change in the type of solution as the polarity of the solvent decreases, showing interstitial solution in water on the one hand and substitutional in acetone on the other. Therefore, in the case of hydrocarbons (or hydrocarbon residues) and water substitutional solution is not to be expected. This can also be roughly demonstrated using regular solution theory35taking methane as the hydrocarbon. From the solubility of methane in cyclohexane at 25’ and 1 atm.,35the substitutional solubility of methane in water at the same temperature and pressure has been estimated using the solubility parameters 2.6, 8.2, and 16.8 for methane, cyclohexane, and water, respectively. It has been found that the mole fraction of methane in water expected on this basis is 3.5 ;< 10-lo which is about one-millionth of the actual solubility (2 X 10-5).36 In contrast, the interstitial solubility of methane (and other hydrocarbons) has been calculated by NBmethy and Scheragaa and the calculated value is in reasonable agreement with experiment. The stability of the interstitial solution must depend on the sizes of the solute particles and of the interstices. For example, a small particle in a comparatively large cavity cannot stabilize the structure and thus the system is unstable. However, a large solute particle may draw upon more than one cluster in order to accommodate It is most probably for this the nonpolar m ~ i e t y . ~ reason that no clathrates of helium, neon, or hydrogen by themselves are known,36 whereas hydrates of innumerable cornplex molecules have been prepared. 37 Formation of Urea-Water Clusters. It can be expected on structural grounds that urea enters into cluster formation in aqueous solutions since it was inferred29 that hydrogen-bonded solutes (e.g., “3) or groups (e.g., OH or NHZ)“don’t alter water structure much, if at all” because these solutes or groups “should be able to enter clusters with only slight distortion and The Journal of Physical Chemistry
MOIIAMMAD ABU-HAMDIYYAH
to transmit both cluster-forming and cluster-disrupting tendencies.” Urea falls into this group of solutes and the carbonyl group should modify the structure only slightly. Evidence to support this point of view comes from two main sources, from the physical properties of pure aqueous urea solutions and from the fact that this is the only structure capable of accounting for the effects of urea upon the “solvent” power of aqueous solutions towards nonpolar solvents including hydrophobic bonding. (a) Physical Properties of Aqueous Urea Solutions. (1) Solubility. It is a well-known fact that urea has an extremely high solubility in water (about 20 M at 25’). The ease of mixing of urea with water means that urea is able to compete with water molecules for the hydrogen bonds. (2) Heat Capacity. The partial molal heat capacity of urea at infinite dilution, (which is indicative of the environment of the urea molecule at this dilution)21a2is of the order of +20 cal./mole deg., which is very close to that of the solid.38 This shows that a similar environment exists in the solid state and in aqueous urea solutions. As it is known that in the solid the urea molecule is hydrogen bonded,39the same must be true in water except that here, at infinite dilution, the molecules surrounding the urea molecules are no longer urea but only water molecules. Any analogy of urea to ions, whether structure forming or breaking, is ruled out completely by the partial molal specific heats of both structure-breaking and structure-forming ions at infinite dilution which are large and negative29in contrast to the large and positive value for urea. (3) Viscosity. If the extensive hydrogen bonding indicated in aqueous urea solutions were fundamentally different from that existing in pure water and responsible for its cooperative structure, i.e., if water molecules hydrogen bonded to urea were not able to hydrogen
cures
(32) R. W. Gurney, “Ionic Processes in Solution,” McGraw-Hill Book Co., Inc., New York, N. Y., 1953,p. 54. (33) L. A. D’Orazio and R. H. Wood, J . Phys. Chem., 67, 1435 (1963). (34) H. G. Hers and M. D. Zeidler, Ber. Bunsenges., 68,821 (1964). (35) J. H. Hildebrand and R. L. Scott, “The Solubility of NonElectrolytes,” Reinhold Publishing Corp., New York, N. Y., 1950, p. 134. (36) J. H. van der Waals and J. C. Platteen, Advan. Chem. Phys., 2, 1 (1959). (37) G. A. Jeffrey and R. K. McMullan, J . Chem. Phys., 37, 2231 (1962),and references therein. (38) E. J. Edsall and J. T. C o b , “Proteins, Amino Acids and P e p tides as Ions and Dipolar Ions,” Reinhold Publishing Corp., New York, N. Y., 1943,pp. 163, 169. (39) A. K. Kitaigordskii, ” Organic Chemical Crystallography,’ ’ Consultants Bureau, New York, N. Y.,1961,p. 154.
EFFECT OF UREAON THE STRUCTURE OF WATERAND HYDROPHOBIC BONDING
bond with other neighboring molecules, then this would also be reflected in the value of the B coefficient in Dole-Jones viscosity equation.a2 In fact, the B coefficient for urea in water is a small positive quantity indicating only a weak structure-forming tendency when compared to the corresponding values of structureforming ions in water.a2 On the other hand, if the urea molecules would tend to destroy the structure of water and produce disorder, then the B coefficient would be a negative quantity.a 2 (4) Dielectric Constant. Further evidence which points towards the increase of structure in the aqueous urea solution as a result of hydrogen bonding comes from dielectric constant measurements which show that addition of urea to water raises the value of the dielectric constant.40 A high dielectric constant is indicative of the existence of strong cooperative structures which give rise to a strong orientation polarizatie@ and flickering c1usters.l’ That this large positive increment in the dielectric constant of water is not due to the large dipole moment of urea ( p = 4.56 D.40,41) but due to a structural effect as explained above, may be seen from the fact that Ae/Ac for urea equals 2.72, whereas for symmetrical dimethylurea, which would not be able to enter into cluster formation and which has a larger dipole moment than urea ( p = 4.8O4Os4l),Ae/Ac is approximately zero.41 Furthermore, if urea were to destroy the cooperative structure of the solution, as is the case of structure-forming and -breaking ions,2gthen the dielectric constant would decrease and not increasie as found by e ~ p e r i m e n t . ~Thus ~ again it is clear that urea does not behave like an ion. (5) Finally, if urea is indeed very similar to water in hydrogen-bonding ability and is thus able to participate in cluster formation as indicated in the previous paragraphs, then aqueous urea solutions should be nearly ideal. The activity of water in urea solutions up to the saturation point (20 M at 25’) has been m e a s ~ r e dand ~ ~ ,led ~ ~to the conclusion that the soluThis also rules tions were indeed “nearly out any association of urea molecules to themselves in the solution, since any significant amount of selfassociation would produce a large nonideality. (b) The Eg’ect of Urea on the Strength of the Hydrophobic Bond. Recent investigations which have already been cited21-n show that addition of urea to aqueous solutions weakens hydrophobic bonding, except2426for the CH, and CzHe molecules which will be discussed separately. Since, as has already been pointed out, hydrocarbon solutes dissolve interstitially in water on one hand and addition of urea to water does not diminish the polarity of the solvent on the other,22it is expected that hy-
2723
drocarbon solutes would also dissolve interstitially in aqueous urea solutions. However, in view of the fact that urea is capable of forming clathrates in the solid state, something similar might be taking place in aqueous urea solutions, which would account for the increased solubility of hydrocarbons in the latter solvent. That this does not happen in aqueous urea solutions is indicated by the near ideality of urea-water mixtures and the following: firstly, it has been shown45that the formation of urea clathrates is restricted to the solid state only; secondly, addition of urea to water increases the solubility of substances that have branchedhydrocarbon moieties24which are incapable of forming urea clathrates in the solid state46; and thirdly, if pure urea clathrates were to form in aqueous urea solutions, the increase in the solubility of the hydrocarbon24or the increase in the c.m.c. of the surfaceactive agent22would be a nonlinear function of the urea concentration and the solubility would rise steeply because of the cooperative nature of cage formation as has been pointed out by Tanford, et aLZ5 However, it is found experimentally that the increase in the solubility of the hydrocarbon or the increase in the c.m.c. is almost a linear function of the urea concentration.21sz2s25 Since interstitial solution seems, therefore, to be the mechanism by which the hydrocarbon residues stay in solution, their increased solubility in urea solution means that the number of possible interstices which can accommodate the molecularly dispersed nonpolar moieties has increased (or more precisely, that the formation of cavities to accommodate the hydrocarbon molecules has become easier in aqueous urea solutions). It seems that hydrophobic bonding is weakened also in other cases where the three-dimensional hydrogen bonding in solution increases.47 Urea, which is a very good hydrogen-bonding agent, actively participates in the formation of “mixed” clusters in aqueous urea solutions. There is evid e n ~ e that ~ ~ these * ~ ~ clusters are less thermally stable (40) J. T. Edsall and J. Wyman, “Biophysical Chemistry,” A c e demic Press,Inc., New York, N. Y.,1958,p. 323. (41) J. Wyman, Chem. Rev.,19, 213 (1936). (42) R. A. Robinson and R. H. Stokes, “Electrolyte Solutions,” Academic Press, Inc., New York, N. Y., 1955,p. 18. (43) G . N. Lewis.and G . H. Burrows, J. Am. Chem. Soc., 34, 1515 (1912). (44) G. Scatchaid, W. J. Hamer, and S. E. Wood, &id., 60, 3061 (1938). (45) J. A. A. Ketelaar and B. A. Loopstra, Reo. t m v . chim., 74, 113 (1955). (46) J. A. A. Ketelaar, “Chemical Constitution,” 2nd Ed., Elsevier Publishing Go., Amsterdam, 1958,p. 367. (47) M.Abu-Hamdiyyah and K. J. Mysels, J . Phys. Chem., 69,1466 (1965).
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MOHAMMAD ABU-HAMDIYYAH
than the pure water clusters. This indicates that the enthalpy is higher in the aqueous urea solutions and therefore the driving force for the spontaneous transfer of nonpolar solutes from pure water to aqueous urea solution is entropic in origin. According to the concepts developed here, the increase in entropy can be attributed to the increase in the number of available cavities or interstices for the solute molecules in aqueous urea solutions (again because their formation becomes easier). If the above conclusions are correct, then it is possible to predict the effect of alkyl-substituted ureas on the strength of the hydrophobic bond. First, an alkylurea derivative will have a decreased ability to hydrogen bond with water compared to that of the unsubstituted urea, and this tendency will increase with the number of the substituent groups, and, as a result of this, the ability to participate in cluster formation will also be reduced. Secondly, the alkyl (or the nonpolar) substituent groups themselves, when they are introduced in the solution (as part of the urea derivative), can only accommodate themselves in cavities in the aqueous solution with the result that the number of “would be” cavities in the solution decreases. We see that both of these properties of alkylurea deriv* tives work against the weakening of the hydrophobic bond and it is expected, therefore, that these compounds would have an effect opposite to that of the unsubstituted urea. There are no experimental data regarding the effect of these derivatives on simple systems involving hydrophobic bonding, but recently a study of the effect of several urea derivatives on the denaturation of serum albumin was reported.48 It was found that the alkyl-substituted ureas have an effect opposite to that of urea itself and this action increases with the length and the number of alkyl substituents. It must be pointed out, however, that denaturing of proteins seems to be a complex process which is incompletely understood and therefore other factors besides hydrophobic bonding may be involved in this process. However, if the point of view that the native structure of globular proteins (of which serum albumin is a member) in aqueous solution is primarily the result of hydrophobic bonding as suggested by Tanford, et a1.,49-5ais accepted, then the action of alkylurea derivatives on the native structure of serum albumin is readily understandable. The results of Schick27are also similarly explained. The explanation that the increase in the c.m.c. of nonionic surfactants is due to the increased hydration of the ether oxygens because urea destroys the water structure and thus makes water molecules more free The J
O U T o ~j
Ph.ysica1 Chemistry
and therefore more available for hydration is very unlikely for the reasons already stated. Again, if urea were to play the same role as an ion that draws the water molecules towards itself and prevents them from participation in hydrogen bonding with other neighboring water molecules as was suggested,27 then the c.m.c. of these nonionic surfactants should decrease, and not increase as found by experiment, on addition of urea. Since urea is a nonionic solute its addition to a colloidal association electrolyte solution affects mainly the structure of the solution. Thus the findingz7 that urea increases the c.m.c. of various dodecyl sulfates in the order R4N+ > Na+ > Li+ seems to be indicative of the effect of these ions on the structure of the solution. Li+ has the most disturbing effect, by polarizing the neighboring water molecules to itself and preventing them from participation in hydrogen-bonded clusters; Na+ effect is less than Li+ and RdN+ has the least effect. In addition, the tetraalkylammonium ion has a tendency to leave the surface of the micelle for the bulk solution when urea is added because the formation of clusters (or cavities) to house such a hydrophobic ionz9becomes easier in the presence of urea as has already been pointed out. Therefore, we have two effects which are acting in the same direction in the case of the tetraalkylammonium ion to enhance the increase in the c.m.c. and indeed the effect is much larger when the tetraalkylammonium ions are the counterions than when the counterion is Na+ or Lif. According to these ideas, urea should increase the c.m.c. of RS04M according to their counterions in the following order: R4N+ > Cs+ > Rb+ > K+ > Na+ > Li+. In the absence of urea the effect of these ions is mainly electrostatic as governed by the charge and the effective size of the ion (the structural effect is small compared to the electrostatic) and, therefore, the explanation of Mysels, et u Z . , ~ ’ J ~ still holds. The Anomalous Solubility of Methane and Ethane. Water and urea molecules have different geometries and, since both partake in the formation of clusters, it is expected that these would not be as symmetrical (48) J. A. Gordon and W. P. Jencks, Bwchembtry, 2, 47 (1963). (49) C. Tanford, P. K. De, and V. G. Taggart, J.Am. Chem. Soc., 82, 6028 (1960). (50) C. Tanford, &id., 84, 4240 (1962). (51) C. Tanford and P. K. De, J. Biol. Chem., 236,1711 (1961). (52) P. L. Whitney and C. Tanford, &id., 237, 1735 (1962). (53) J. F. Brandt, J . Am. Chem. SOC.,86, 4302 (1964). This paper gives a good discussion of Tanford’s views. (54) K. J. Mysels, Final Report, Project NR 356254, O5ce of Naval Research Contract NONS274 (00).
EFFECTOF UREAON THE STRUCTURE OF WATER AND HYDROPHOBIC BONDING
as those of pure water and that the sizes and symmetries of the interstices would also be different. This is indicated by the sol~bilities~~ of methane (in the temperature range 5-45’) and of ethane (below 20’), which are reduced upon addition of urea. Thus here, in contrast to all previous cases, the molecularly dispersed hydrocarbon is destabilized, i e . , hydrophobic bonding is strengthened. As was pointed out earlier, two factors govern the solubility of a nonpolar solute in the interstitial solution: the size of the solute molecule and that of the cavity in the solution. Methane is a small molecule and the cavities that accommodate it in aqueous solutions are small. Glew15 found that the structure of water around the methane molecule is a broken-down methane gas-hydrate type polyhedron and that the size of the interstitial hole occupied by the methane molecule has very nearly the same size as that occupied by the methane molecule in the gas-hydrate lattice, namely the 20-coordinated site. Methane molecules do not occupy the comparatively larger cavities which are approximately equal to the 24-coordinated sites. In other words, methane molecules are incapable of stabilizing a cluster that forms a comparatively large cavity. Since addition of urea to an aqueous methane solution decreases the solubility of methane (in the given temperature range), some of the clusters which originally housed the methane molecules must have been destroyed and it became more difficult to form such small cavities. The same thing happens in the case of ethane, which occupies in pure water an interstice of approximately the 24-coordinate site size,15 viz., addition of urea decreases the solubility of ethane in the range of temperature 5-20’. Hence, it may be concluded that urea destroys some of the “polyhedra” or “icebergs” which are formed around methane and ethane molecules, and that the cavities formed by the mixed ureawater clusters are less favorable for housing the small molecules than those of water with the result that a decrease in solubility occurs. However, as the temperature rises, other factors operate which tend to normalize the solubility of methane (or ethane) in aqueous urea solutions. Firstly, the number of cavities formed by pure water, the polyhedra housing the methane (or ethane) molecules, will tend to decrease with increasing temperature as is evident from the decrease of the solubility of methane (and ethane) i n pure water in the temperature range
2725
5-50’. Secondly, although the mixed clusters are less favorable for housing the methane (or ethane) molecules than those of pure water, they tend to persist at higher temperatures, which is related to the fact that the enthalpy of the mixed clusters is higher than that of pure water clusters (the polyhedra). The first, which is the larger, factor tends to decrease the solubility as a whole; the second factor tends to increase the solubility in urea solutions. Moreover, the second factor favors the larger molecule, ethane, more than methane, the smaller molecule, and so the normalizing temperature for the former tends to occur at a lower temperature than that of the latter. This tendency is very clearly seen in the solubility curves of methane and of ethane24 in aqueous urea and in water. The curves for ethane intersect at 20’ and those for methane-the smaller molecule-approach each other in the investigated range and would intersect probably just above 50’. I n this connection it would be very interesting to find out experimentally the effect of urea on the solubility of the rare gases in water. It is expected that the solubility of the rare gases in aqueous urea would be less than their solubility in pure water at low temperatures.
Summary and Conclusions (1) Nonpolar solutes dissolve interstitially in strongly polar solvents. (2) Addition of urea to water increases the hydrogen bonding in solution. (3) Urea partakes in the formation of urea-water clusters that are responsible for the formation of interstices in the solution which accommodate the nonpolar moieties of the solute and therefore in this sense may be looked upon as a structure former. (4) Likening urea to an ion does not fit the facts and must therefore be rejected. (5) It is easier to form a large cavity in aqueous urea solution than in pure water and the converse is true for small cavities that house methane or ethane molecules and presumably also the rare gases at low temperatures.
Acknowledgments. The author is very grateful to Professor Karol J. Mysels for his help and advice throughout the preparation of this manuscript. This work was supported by a fellowship sponsored by the Continental Oil Company and by Public Health Service Research Grant GM 10961-01 from the Division of General Medical Sciences, Public Health Service.
Volume 69,Number 8 August 1866
