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The effects of 1-propanol, 2-propanol, NaCl, urea, and [beta]-D-fructose on the molecular organization… To, Eric Chun Hin 1999

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THE EFFECTS OF 1-PROPANOL, 2-PROPANOL, NaCl, UREA AND (3-D-FRUCTOSE ON THE MOLECULAR ORGANIZATION OF WATER by Eric Chun Hin To B.Sc, The University of British Columbia, 1997 A THESIS SUBMITTED IN PARTIAL FULLFILLMENT OF THE REQUIREMENT FOR THE DEGREE OF MASTER OF SCIENCE in THE FACULTY OF GRADUATE STUDIES (Department of Chemistry) We accept this thesis as conforming to the required standard THE UNIVERSITY OF BRITISH COLUMBIA April 1999 © Eric Chun Hin To, 1999 In p resen t ing this thesis in partial fu l f i lment of the requ i rements for an a d v a n c e d d e g r e e at the Univers i ty of Brit ish C o l u m b i a , I agree that the Library shall m a k e it f reely avai lable for re fe rence a n d study. 1 further agree that p e r m i s s i o n for ex tens ive c o p y i n g o f this thesis fo r scho lar ly p u r p o s e s may be granted by the head o f m y d e p a r t m e n t o r by his o r her representat ives . It is u n d e r s t o o d that c o p y i n g o r p u b l i c a t i o n of this thesis for f inancial gain shall n o t be a l l o w e d w i t h o u t my wr i t ten p e r m i s s i o n . T h e Univers i ty of Brit ish C o l u m b i a V a n c o u v e r , C a n a d a D e p a r t m e n t D E - 6 (2/88) A B S T R A C T Excess partial molar enthalpies of 1-propanol, HIPE, were measured directly and accurately in four ternary aqueous systems by titrating small increments of 1 -propanol into mixtures comprised of 1-propanol, H2O, and a "third component" at 25°C. Firstly, in the 1-propanol - H2O - 2-propanol mixtures, the enthalpic interaction function between 1-propanol molecules, Hi P ip E , were evaluated. Based on the knowledge accumulated in our laboratory on the mixing behaviour of each solute on the hydrogen bond network of water, the result indicates that 1 -propanol and 2-propanol modify the molecular organization of water in an identical and additive manner. Secondly, in the 1 -propanol -H2O - NaCl mixtures, the enthalpic interaction function suggests that though NaCl acts as a structure breaker, it also modifies the hydrogen bond network of water in a positively cooperative manner with 1-propanol, a relatively hydrophobic solute. However, the results also suggest that a NaCl molecule binds to 7 to 8 water molecules, which were made unavailable for 1-propanol to interact with. The remaining bulk water away from solute NaCl would interact with 1-propanol as it is pure water. Thirdly, in the 1-propanol-H20-Urea mixtures, the total vapour pressures of this ternary system were measured in addition to HipE. The Boissonnass method based on the Gibbs-Duhem relation were employed to calculate the partial pressures, excess chemical potentials of each component. Excess partial molar entropies of 1-propanol were then evaluated. Both enthalpic and entropic interaction functions imply that urea molecules blend themselves into the hydrogen bond network of water keeping the connectivity of the hydrogen bonds intact but concurrently the water mediated solute-solute interaction is reduced. Finally, the interactions of IP-IP in terms of enthalpy and entropy, HIPIPE and SIPIPE , respectively, in l-propanol-H20-(3-D-Fructose mixtures suggest that fructose and 1-propanol modify the hydrogen bond network of water in a positively cooperative manner, much like the l-propanol-HiO-Glycerol systems studied previously. However the different appearances in the interaction functions of fructose as compared to other solutes suggests that the solvation of fructose might involve different mixing mechanisms which are yet to be determined. T A B L E O F C O N T E N T Page Introduction i - o Experimental 16-22 Results and Discussions 1P-H 20-2P 23-30 lP-H 2 0-NaCl 31-37 lP-H 20-urea 38-50 lP-H20-B-D-Fructose 51-65 Conclusion 66-68 References 69-71 Appendix Raw Data Tables 72-82 Derivation of equations 83-84 Numerical evaluation of partial pressures and chemical potentials 85-86 LIST O F T A B L E S AND FIGURES Table Descriptions Page 1 Thermodynamic Quantities (p, T, ni) system 4 A l H I P E in 1P-H 20 -2P at 25°C. 72 A2 H . P E i n l P - H 2 0 - N a C l a t 2 5 ° C . 73 A3 H , P E i n l P - H 2 0 - U r e a a t 2 5 ° C . 74 A4 Vapour Pressures in lP-H 20-Urea at 25°C. 75 A5 Partial Pressures in lP-H 20-Urea at 25°C. 76-77 A6 T S i P E in lP-H 20-Urea at 25°C. 78 A7 H 1 P E at lP-H 20-Fruc at 25°C. 79 A8 Partial Pressures in lP-H 20-Fruc at 25°C. 80 A9 U I P E and T S i P E in lP-H 20-Fruc at 25°C. 81 A10 GC analysis: Peak ratio of IP to H 2 0 in lP-H 20-Fruc at 25°C. 82 Figure Descriptions Page 1 Measurement of an excess partial molar enthalpy 4 2 H I P E in binary aqueous IP at 25°C. 8 3 Hi PI P E in binary aqueous IP at 25°C. 9 4a IB A-IB A enthalpic interaction 10 4b DMSO-DMSO enthalpic interaction 10 5 V B E E in binary aqueous BE at 25°C. 11 6 Relative positions of transition points of some small hydrophobic solutes 13 7 Calorimeters HP3457A and L K B Bromma 8780 17 Figure Descriptions Page 8a A typical exothermic reaction 19 8b A typical endothermic reaction 19 9 Vapour Pressure measurement apparatus 20 10 Excess partial molar quantities of 2P in 2P -H 2 0 at 25°C. 24 11 H I P E in 1P-H 20-2P of various x 2 P ° at 25°C. 26 12 H I P I P E in 1P-H 20-2P of various x 2 P ° at 25°C. 28 13 1P-2P Mixing Scheme Transition Diagram 29 14 NaCl-IP Mixing Scheme Transition Diagram 32 15 H 1 P E in lP -H 2 0-NaCl of various x N a C i ° at 25°C. 33 16 H I P I P E in lP -H 2 0-NaCl of various x N a ci° at 25°C. 35 17 H 1 P E in 1 P-H 20-Urea of various x u r e a 0 at 25°C. 39 18 H I P I P E in lP-H 2 0-Urea of various x u r e a 0 at 25°C. 41 19 p u r e a E in lP-H 2 0-Urea of various x,P at 25°C. 43 20 P I P E in lP-H 2 0-Urea of various x u r e a ° at 25°C. 44 21 T S I P E in lP-H 2 0-Urea of various x u r e a ° at 25°C. 45 22 T S , P I P E in lP-H 2 0-Urea of various x u r e a ° at 25°C. 46 23 Enthalpic and entropic interaction between IP and urea 47 24 Mixing Scheme Transition boundary in lP-H 2 0-Urea at 25°C. 48 25 H I P E in lP-H 2 0-Fruc of various x f r u c ° at 25°C. 52 26 H I P I P E in lP-H 2 0-Fruc of various x f r u c ° at 25°C. 53 27a Mixing Scheme and Phase transition of lP-H 2 0-Fruc at 25°C. 55 27b Mixng Scheme boundary of lP-H 2 0-Fruc at 25°C. 56 Figure Descriptions Page 28 H I P E in lP-H 20-Fruc of various x f r u c ° at 25°C. 58 29 T S I P E in lP-H 20-Fruc of various Xfruc 0 at 25°C. 59 30 T S I P I P E in lP-H 20-Fruc of various x f r u c ° at 25°C. 60 31 H 1 P E in lP -H 2 0-Gly of various x g , y ° at 25°C. 62 32 H I P I P E in lP -H 2 0-Gly of various x g i y ° at 25°C 63 33 Mixing Scheme Transition boundary of lP -H 2 0-Gly at 25°C. 64 ACKNOWLEDGEMENT I would like to thank the following individuals for their participation in some of the data measurements presented in this thesis. IP-H2O-2P system: Dr. Jianhua Hu, Wesley M . Chiang, Daniel H.C. Chen, Jason S.Y.Sze, Dr. Charles A. Haynes, and Dr. Y. Koga. lP-H 20 -NaCl system: Hiroshi Matsuo, Denise C.Y. Wong, Dr. Seji Sawamura, Dr. Yoshihiro Taniguchi, and Dr. Y. Koga. lP-H 20-Urea system: Dr. Jianhua Hu, Dr. Charles A. Haynes, and Dr. Y. Koga. I would also like to thank my co-supervisor Dr. G. Patey and Dr. Y. Koga for their continuous guidance and support throughout this research project. This project was supported by the National Sciences and Engineering Research Council of Canada. INTRODUCTION "Water's most remarkable feature is deception, for it is in reality a substance of infinite complexity, of great and inaccessible importance, and one that is endowed with a strangeness and beauty sufficient to excite and challenge anyone making its acquaintance..." Owen R. Fennema University of Wisconsin - Madison Editor of "Food Science, 3 r d edition" When the first living cell was formed on this planet about 4 billion years ago, it had already been going through a long process of evolution. It all started from individual atoms, to the formation of small molecules like H2O, CO2, and N 2 . These molecules, along with other ones, were blasted into the air due to volcanic eruption. Bombarded by ultraviolet radiation and lightning from intense storms, they collided to form stable chemical bonds of larger molecules such as amino acids and nucleotides. These are the building blocks of proteins, nucleic acids, and fatty acids, which in turn are the raw material for a living cell. But what special "element" actually initiates, nurtures, sustains, and inevitably, terminates life? Water, uniquely, is the only substance on this planet that occurs abundantly in all three physical states of matter. In a solid state, or ice, it can be found as glaciers and ice caps, on water surfaces in winter, as snow, hail, frost, and cloud formed of ice-crystals. In a liquid state, it appears as rain cloud formed of water droplets, as dew on vegetation, and as swamps, lakes, rivers, oceans on 3A of the surface of the earth. In a gas state, it occurs as water vapour.(1) Being the major constituent of living matter, water serves as a solvent, transporting, combining, and chemically breaking down substances such as fats, carbohydrates, proteins, salts, and similar chemicals. Blood in animals and sap in plants also consist largely of water, which serves to transport food and remove waste materials. 1 Out of the hundreds of thousands of common chemical compounds in our biosphere, why is water the most important of all? By comparing the properties of water with those of molecules of similar weight and atomic composition (CH 4 , N H 3 , H2S, etc.), it is possible to determine if water is unique. Based on this comparison, water is found to melt and boil at unusually high temperatures; to exhibit unusually large values for surface tension, thermal conductivity, permittivity, heat capacity, and heats of phase transition; to have a moderately low value for density; to exhibit an unusual attribute of expanding upon solidification. These peculiar physical characteristics ( 2 ) of water suggest its uniqueness. In addition, water is a universal solvent that is able to interact with polar, non-polar, charged, and amphiphilic molecules. These interactions, indeed, raise a lot of interests among scientists. Among those, Franks et. al. ( 3"5 ) were considered one of the greatest contributors. Consequently, aqueous solution chemistry received much attention in the past as well as in present time. Most biological systems are in aqueous phase. It is almost an obligation to investigate the interactions among water and other biomolecules since they lead to familiar events such as protein folding/unfolding, nucleic acid coiling/supercoiling, structural positioning of carbohydrate in water, etc. However, one must be careful about the usage of the word "interaction", which could carry two definitions in biophysical chemistry. The first kind is a direct binding of a solute onto sites of a biopolymer. ( 6 7 ) The second, however, is an indirect interaction among solutes within a water medium (8~ n \ in which, the hydrogen bond network (HBN) of water plays a vital role. It is the latter kind of interaction that this thesis is focused on. The approach employed in this research project is purely thermodynamical. In biochemistry and biophysics, thermodynamic data could be used for relating molecules, small and large, to equilibrium properties of systems involving such molecules. The absolute advantage of this approach is that thermodynamic properties of aqueous solution are directly and readily measurable with the help of sophisticated instruments that are both precise and accurate. The resulting data could then be used to elucidate intermolecular interactions in a macroscopic point of view. A vast amount of work has been devoted to this subject and studies up to the early 1980's were comprehensively reviewed. ( 3 " 6 ) The Koga group has measured the excess partial molar quantities of various aqueous systems^  1 3 ' 2 1 ) extensively for the last several years. These quantities have served to deepen the insights into the nature of aqueous solutions. The excess partial molar enthalpy of solute i, Hj E , is defined as (3HE / 5ni)nw, p, T-This exhibits the enthalpic contribution of species i to the entire system where the moles of water, pressure and temperature are kept constant. In other words, this is the actual * E enthalpic situation that solute i is experiencing in the system. The measurement of Hi is illustrated in Figure 1. So how would this enthalpic situation change when another infinitesimal amount of i is added to the system? By differentiating H E by moles of i, we hence define HjjE = N (5HjE / 5nj) (1) where N is the total moles of solvents and solutes. The stability criterion states that, if this H j E quantity is negative, solutes in the systems enjoy the addition of more solutes, or they simply attract each other. Namely, a favourable thermodynamic interaction(21) takes place. However, if H j E is positive, solutes repel or would try to stay apart from one 3 Figure 1 : Measurement of an excess partial molar enthalpy. + 8 rij • rij + 5rii n w 8q = H (nj + 8nj, nw) - H (nj, nw) - H° 5n; lim 5a, = lim [ H (n, + Sn,, n»,) - H ("nu 1 - H° 8nj->0 8nj 8nj->0 8nj = (SH/ani) n w > p,T - Hi 0 Hj E = (9H E / dni) „W,P,T Excess partial molar enthalpy of solute i Where 8q = change of enthalpy of the system. H = enthalpy of the system consisting of components in ( ) n; = moles of solute n w = moles of solvent H° = enthalpy of pure solute i E E Sj , Vj can also be defined in a similar fashion. Table 1: Thermodynamic Quantities (p, T, nj) system Oth 1st Derivative Derivative 2nd Derivative 3rd Derivative 4th Derivative G H;{T} C p ; {T,T} 5CP/5T; {T,T,T} dCp(i)/dp;{T,T,ni,p} S: {T} Hi; {T,ni} aC p /5 p ; {T,T,p} 92Hi/ani2;{T^ii,ni,ni} V; {p} S;; {T,ni} Cp(i); {T,T,ni} S2Si/ani2;{T,ni,ni,ni} Vi; {p,ni} dHJdni; {T,ni,nj} K T ; {p,p} 5Si/5ni; {T,ni,nj} a p ; {p,T} dVi/dm; {p,ni,nj} <(Axi)2>;{ni,ni} K T(i); {p,p,rii} ap(i); {p,T,ni} another in the system. In other words, there is a thermodynamic unfavourable interaction among the solutes. For entropy, however, the signs are opposite. The excess partial molar entropy of i , S E is defined as (5S E / dn\). Also , the entropic interaction is defined as follows: S i i E = N ( 5 S i E / a n i ) (2) If Sjj E > 0, a favourable interaction takes place while S j E < 0 represents an unfavourable interaction in terms of entropy. Consequently, the interaction functions H j E and S ; E become extremely important tools to study the thermodynamic situations in an aqueous solution. Table 1 ( 2 2 ) is a list o f the derivatives of Gibbs Free Energy, G . The variables in parentheses denote the quantities that G is differentiated with respect to. For instance, isobaric thermal expansivity, a p {p,T} = [(3G/<3P)T/3T]p , is to first differentiate G with respect to pressure at constant temperature and then secondly with respect to temperature at constant pressure. In most conventional thermodynamical approaches, partial molar quantities (2 n d derivatives of G) have been rarely measured. Nonetheless, H j E , S j E , V j E , and the rest of the 3 r d derivatives of G seem to be quite useful as explained above. We would like to differentiate the 2 n d derivatives of G to arrive at the 3 r d derivatives of G because we believe that the higher order o f G , the more detailed the information/ 1 3 ) Indeed, Hepler used (SC P / dp); {T,T,p}, a 3 r d derivative to discuss the degree of "structure" in H 2 0 and D 2 0 . ( 2 3 ) However, by differentiating, the uncertainty of the data would easily be magnified 10 fold or even more. However as long as we measure the 2 n d derivatives of G 5 very carefully in small increments of i with high reproducibility, the quality of the corresponding 3 r d derivatives should be of sufficient accuracy to achieve our purposes. For the sake of convenience, equation (1) can be re-written in (N, XJ) variable system as HiiE = N (dHiE/ani) = (1 -xO (dHiE/axOnw (3) where N is the total moles of solution and x; is the mole fraction i. A detailed derivation is included in the appendix. Our investigation into the intermolecular interaction begins with a simple two component system. The purposes are to see how the solute contributes thermodynamically (ie. enthalpically, entropically, volumetrically, etc.) to the entire system and how this contribution changes as the amount of solute increases. Furthermore, the hydrophobicity/hydrophilicity of the solutes certainly has a gigantic impact on the way they mix with the water molecules (we named MIXING S C H E M E S ( 1 3 ) , which will be explained later) and on the hydrogen bond network of water in the system. In order to obtain information of the above interaction functions, the thermodynamic quantities of H j E and Sj E are to be determined. First, we measured directly the excess partial molar enthalpies of the solute, H E , by adding a small amount into water. Then taking the slope of H E at intervals of 8XJ = 0.002, the enthalpic interaction functions, HjEwere obtained. For SjjE, the total pressure of the mixture is measured as a function of mole fraction X ; . Secondly, the partial pressures of each component (solute and solvent) are evaluated. A numerical analysis (see appendix) based on the Gibbs-Duhem relation is employed to calculate partial pressures of each component and subsequently, the excess chemical potential of i, U j E was calculated, using 6 HiE = H E - TSi E = RT ln( Pi / X i Pi°), (4) where Pj is the partial pressure of i, Xj is the mole fraction of solute i, and P ° is the partial pressure of pure i at standard state. Knowing both u,E and H E at a fixed temperature will yield S E . After graphical differentiation, Sj E , the entropic interaction function is determined. Out of more than a dozen different solutes studied by the Koga group, the interaction among 1-propanol molecules and that between 1-propanol and water molecules are the best understood and characterized.(21)The excess partial molar enthalpies of 1-propanol, HipE and the IP-IP enthalpic interaction function, H I P I P E in a IP-H2O binary system at 2 5 ° C are shown in Figure 2 and 3 , respectively. In the H I P I P E vs X I P plot, a peak shaped relationship was observed. This 3 R D derivative of G anomaly is the most common type observed among hydrophobic solutes. Other types of anomaly were also observed in aqueous binary isobutyric acid (IBA) ( 1 4 ) and aqueous binary dimethylsulfoxide ( D M S O ) ( 1 4 ) systems are shown in Figure 4. In spite of the different shapes of anomalies observed, it is possible to divide up each curves into at least two separate regions indicating the presence of two qualitatively different ways the solutes and the solvent (water) molecules are mixed within the solution. All the data obtained are consistent with the following interpretation. Mixing Scheme I is operative in the most water-rich region, in which the HBN of water remains percolated(24). Namely, at any instance, the HBN is connected throughout the entire region of solution, as suggested by Stanley et al. ( 5 ) in the 1980's. Starting from pure water, when the first solute molecule enters into water environment, the HB probability of water in the immediate vicinity increases substantially, resulting in a formation of a clathrate or an "ice-like" structure(13) 7 H 1 P E in 1P-H20 system at 25°C o -4 H - 6 H -8 H 1 0 1 2 0 . 0 0 0 . 0 2 0 . 0 4 0 . 0 6 0 . 0 8 0 . 1 0 0 . 1 2 0 . 1 4 0 . 1 6 Figure 2: Excess partial molar enthalpies of 1P, H 1 P E in binary aqueous 1P system at 25.00°C. Each injection was 0.6869 + 0.005mL. H 1 P. 1 P E in 1P-H20 system at 25°C 250 200 150 * 100 tu 50 0 -50 -X — ! M o o "1 Y I IT i i III i i i i 0.00 0.02 0.04 0.06 0.08 0.10 0.12 0.14 0.16 0.18 0.20 Figure 3: Enthalpic 1-propanol -1-propanol interaction at 25.00 C Point X is the start of Mixing Scheme Transition from Region I to U; point M is the Nominal transition point; point Y is the end of Mixing Scheme Transition. 9 0 2 4 6 8 10 12 14 'xDMSO Figure 4b: DMSO-DMSO Enthalpic Interaction Excess Partial Molar Volume of BE in 2-Butoxyethanol (BE) - H20(W) o B S w w CO -4 -6 - 8 -10 T T 1 ' ^ • ' • • • D D ° D D D • • • • • • ••• 0 J l _ 001 J L J L 002 X B E 0.03 0.04 Figure 5: Excess partial molar volume of 2-Butoxyethanol in 2BE-H2O system at 25°C. around the solute. At the same time, the HB probability of bulk H2O away from the ice-like structures seems to decrease. The evidence for this is shown in the initial decrease in V B E E in Figure 5. When the second solute molecule enters the system, similar events (i.e., ice-like structure formation) happen to the second molecule except the changes are less drastic than the first time. This trend continues as the composition of solute increases. The HB probability in bulk H2O continues to decrease but is still high enough to have the entire network connected infinitely. From the study of 2-butoxyethanol -JJ 2Q(26-27) D j n a r y Sy Stem, the results suggested that the physical purpose of Mixing Scheme I was to prepare for the formation of an addition compound ( 2 8 )(a clathrate of solute-(H20)m type) at a low enough temperature. However, once the HB probability drops below the percolation threshold(25), the entire HBN is no longer infinitely connected. This marks the transition point form Mixing Scheme I to Mixing Scheme II. The transition itself is spanned over a small range in solute composition(21), starting when the weakest hydrogen bond (HB) is broken (Point X in Figure 3) and finishing when the strongest hydrogen bond is destroyed (Point Y in Figure 3). After this point, the solution is interpreted as in Mixing Scheme II where the HBN is no longer connected throughout. The width and the locus of the transition region depend primarily on the hydrophobic moiety of the solute molecule(21). Figure 6 shows the relative transition loci (Point M = nominal transition point) among some small solutes. This shows that, as a general trend, the larger the hydrophobic chain of the solute, the transition from Mixing Scheme I to Mixing Scheme II takes place at smaller values of solute mole fraction. Once the transition is complete, the system is in Mixing Scheme II. The microscopic physical interpretation of this region is the formation of two different kinds 12 of clusters(21). Namely, water-rich and solute-rich clusters are present. X-ray ( 2 9 ' 3 0 ) and light scattering(31) studies provided evidences for this claim from previous works. It was argued that the purpose of region II was to have the molecules in the mixture organized in a way such that it prepares for a liquid-liquid phase separation at the right temperature ( 2 6 ' 32) X i 0.1 2 B E | | | | 2P | | T B A | I B A | | I I Butanone | IP Longest Hydrophobic Chain 0.2 0.3 E t O H M e O H D M S O Shortest Hydrophobic Chain Figure 6. The relative positions (in terms of mole fraction) of transition points of some small hydrophobic solutes. Based on previous work on 2-Butoxyethanol (2-BE) and D M S O ( J J } a third mixing scheme may be generally present in these aqueous systems. As the solute composition * E E increases, Mixing Scheme II would no longer be operative. Instead, Hjj and SH show zero or very small values, suggesting that the thermodynamic environment for the solute 13 is equivalent or almost equivalent to that in the pure state. Clusters of pure solute molecules, perhaps existing as micellar forms or with the H2O molecules adsorbed on the surfaces of these clusters are the physical interpretation of Mixing Scheme III. This further implies that water molecules only act as single gas-like molecules rather than as a mass of HB network in this region (13 ). While the titration calorimeter that was used in this study is convenient for a liquid solute, titrating a solid solute, such as NaCl or urea, could be problematic. To avoid this difficulty, ternary (3-component) systems were studied. These systems consist of water (solvent), 1-propanol, and a third component which is the target of study. Recalling from above, the data were extensively obtained for 1 -propanol- H2O binary system ( 2 1 ). The interaction function for this binary system is the most familiar and the easiest to be analysed among other 2-component systems. Our approach is first to have this "3rd component" already present in H2O. Then partial molar quantities of 1-propanol are measured when 1-propanol is added to the system. By varying the composition of the "3rd component", the effect on the mixing behaviour of 1-P in the system is studied. Namely, the change (if any) in H i P i P E and also that of the width and locus of the Mixing Scheme transition region in the system is observed. Therefore, the thermodyamic impact or the contribution of the "3rd component" can be determined indirectly by studying the thermodynamic behaviour of 1-propanol as a probe. The present work is concerned with 4 different ternary systems. 2-propanol, sodium chloride, urea, and p-D-fructose were selected as the "3rd component" in those systems. Each of these solutes, representing its own special type, is abundant and is of physico-biological significance. Firstly, sodium chloride, a salt and an ionic electrolyte, is 14 a typical non-hydrophobic solute. In fact it is so hydrophilic that it ionizes and dissolves in water readily(1). The mixing behaviour of NaCl is expected to be completely different from that of 1-propanol. Secondly, urea, a non-electrolyte, is also a non-hydrophobic solute. It is commonly found in fertilizer(34). Physiologically, urea is well known for its protein denaturing capability(35). It would be interesting to investigate the thermodynamic reasons behind urea's special capability as well as its effect in water. Thirdly, P-D-fructose is a common sugar found in fruits(1). It is ambiguous in terms of hydrophobicity. The five-membered carbon ring suggests a degree of hydrophobicity whereas the OH group present on each carbon exhibits signs of hydrophilicity. How p-D-fructose mixes with water and with 1-propanol can tell us more about the aqueous nature of this sugar and how it would work together with 1-propanol to affect the HBN of water. Finally, 2-propanol was included into our ternary system for one purpose. The individual impact, thermodynamically and structurally, of 2-propanol was already understood from a binary 2P-H2O system studied previously and was found to be the same as that of IP. Whether the same effect of 2-propanol can be observed in a ternary system with 1-propanol is significant to our strategy used to study the three other solutes. Hence it is worthwhile to include the study of 2-propanol in the beginning of our discussion as a good test for our methodology. 15 E X P E R I M E N T A L ACS reagent grade NaCl (99+%) and urea (99%) were purchased from Aldrich and used as supplied. D(-)-fructose (99.95%) was obtained from Sigma and also used as supplied. 1-propanol and 2-propanol (both 99.5%, HPLC grade) from Sigma-Aldrich was opened in dry nitrogen atmosphere to avoid contamination from air moisture (due to their high hygroscopic property). 1-propanol was subsequently stored in a tightly closed container under nitrogen atmosphere. H2O was triply distilled, the last two distilled in a Pyrex glass still immediately before use. The excess partial molar enthalpies of 1-propanol in IP-H2O-2P, lP-H20-NaCl, lP-H20-urea, and lP-H20-P-D-fructose were measured directly at 25.00°C by titration calorimetry (36'37) using the two equipment setups as shown in Figure 7. The cell initially contained a mixture of H2O and the 3 r d component (i.e., 2-propanol, NaCl, urea, or P-D-fructose) of a known concentration. Then a fixed amount of 1-propanol was titrated into the cell through a syringe, in which the average volume of each shot was determined to be 0.6869+0.0005mL by calibration. The cell contains typically 1 OOmL of solvent. A water bath, with very sensitive temperature controllers, was maintained at 25.00+0.01°C throughout the entire experiment to minimize any variation in heat exchange between the mixture inside the cell and the outside environment. In addition, a mechanical stirrer was employed to ensure the homogeneity of the solution. A small temperature change was detected by a very sensitive semi-conducting thermistor. In the HP 3457A calorimetry system, a standard DC voltage and a known resistance in the circuit were applied. Then measuring the voltage across the 16 HP 3457A upply Thermistor-Water Bath (25°C) — Computer Switch Heater Stirrer 1P addition tubing 1-propanol Syringe O oo 0 0 O O Q Q o LKB Bromma 8780 z \ A 71 ii Figure 7: Calorimeters HP3457A and LKB Bromma 8780 in our laboratory standard resistance and the thermistor would yield the thermistor resistance in the following relationship: Rtherm = (Vsample / V s t (j) X Rstd (5) The thermal response of the system was therefore determined indirectly by measuring the values of the thermistor resistance and the changes in resistance over a period of time. The resistance, R can be related to temperature, T in the equation: R = Aexp(B/T) (6) where A and B are constants. The changes in resistance and temperature are related as the following: AR / AT = -R m * B / T m 2 (7) where the subscript m denotes a mean value. Hence, it follows that AT = AR * T m 2 / R m (8) The L K B Bromma 8700 precision calorimetry system is very similar to the HP system except that the value of resistance is determined by a Wheatstone bridge. For calibration, two heating sequences were performed. Namely, an accurately known amount of energy was supplied to the cell before the addition of 1-propanol. Another one was supplied afterwards. Two examples of R vs Time are shown in Figure 8. From this resistance vs. time trace, the heat evolved in the titration is calculated by comparing it with the two heating sequences. The excess partial molar enthalpies were then calculated as Hj E - 8q /8ni, where 8q is the heat evolved with the amount of titrant 8nj. The uncertainties of each data point were estimated to be less than 0.1 kJmol"1. 18 Figure 8a: A typical Exothermic Reaction a Q) O c ro - i—• w w <D Ron H e 3 t 00 1 Reaction Time Figure 8b: A typical Endothermic Reaction G <u o c -«-» c/) in 0 Time 2^ 100 Torr Gauge 1 Torr Gauge A X 1 \J Source container (eg 1-propanol) Inner bath (25°C) Outer bath («25°C)-Stopcock Gk^ ss vacuum lin& Heater Wire Cell (Only the RHS cell was used| b Vaccum To vacuum A Source container Charged with H20 + 3rd component Figure 9 : Vapour pressure measurement apparatus so Vapour pressure measurements of IP-H2O-2P, lP-HfeO-NaCl, lP-HbO-urea, and lP-H20-P-D-fructose systems were performed using the "Static Method" ( 1 8' 2 0 ). Figure 9 exhibits a schematic diagram of the apparatus used. A glass cell, approximately 20mL, was connected to a gas handling glass manifold and a high vacuum line capable of reaching 10"6 Torr. A heating wire (approx. 20°C higher than the cell) was wrapped around the narrow connecting tube of glass to prevent condensation. The gas handling manifold, along with an MKS Baraton pressure gauge (100 Torr full scale with a sensitivity of +0.001 Torr) were encased in a wooden box in which the temperature was maintained at approximately 40°C. Outside the box, the glass cell was immersed in a water bath, which in turn was placed in a larger bath with an air space between them. The temperature of the outer bath, about 5°C lower than the inner bath, was controlled by a mercury regulator to within 0.1°C. The inner bath temperature was controlled at 25.00°C by a Fisher proportional controller with uncertainty of +0.005°C within a few hours of measurement, and determined by the vapour pressure of pure H2O. A Beckman thermometer was used to monitor the day to day temperature fluctuation of the inner bath, which is approximately within 0.01°C. To measure the total pressure of the mixture, a solution of the 3rd component in H2O of a known concentration was first charged into the cell. The opening was then immediately sealed and the whole system was put under vacuum. After a few Freeze-pump-thaw operations to remove all the unwanted air in the cell, a known amount of 1-propanol was vacuum transferred from a source into the vacuum lines. After thermal equilibrium was reached, the 1 -propanol was transferred into the cell, which was frozen with a liquid nitrogen bath. Once the transfer was completed, the cell was closed off, 21 immersed into the 25°C water bath, and stirred for one hour. Finally, the stirrer was switched off and the cell was opened to the gauge for approximately one hour or until the total pressure reading was stable. The vapour was then retrieved back into the cell by freezing it. Only a negligibly small amount of vapour was pumped away. Hence, the new composition in the cell could be calculated and thus another 1-propanol sample was ready to be charged into the system. The procedure and experimental set up for the external Gas Chromatography employed in the lP-H20-Fruc systems is as follows. About 3mL of a solution with known composition was transferred to an 8mL vial. The vial was subsequently hermetically sealed with a silicon septum and allowed for thermal equilibrium at 25.00+0.01°C for one hour in an water bath. The upper portion of the vial, above the level of the solution and that of the water bath, was heated to ca. 40°C to avoid condensation. A 100+5 mL aliqout of vapour was removed from the head space by a warm (45°C) gas tight syringe and immediately injected in a Carle Analytical Gas Chromatographer 311. Peaks of air, water and IP were printed out. The ratio of the peak area of IP to that of water represents the ratio of the partial pressure of the two components in the gas phase. This information, along with the total pressure measurements, assists us in determining the partial pressure as well as the excess chemical potential of each of the components. 22 R E S U L T S AND DISCUSSIONS I) l-propanol-HUQ^-propanol In a binary aqueous 2-propanol system, the excess partial molar enthalpies, H2pE and the total vapour pressure of 2P-H2O were measured. By the Boisonnass mefhod ( 2 0 ), a numerical analysis ( 1 8' 2 6 ) based on the Gibbs- Duhem relation, the partial pressures and hence the excess chemical potentials of 2-propanol, p2pE were evaluated. Thus, the excess partial molar entropies of 2-propanol, S2pE were determined, using U2 P E and H2pE that were determined by calorimetry. These three thermodynamic quantities of 2-propanol, as a function of 2-propanol mole fraction, are plotted in Figure 10. The composition dependence of these excess partial molar quantities is almost a replica of that for aqueous 1- propanol systems. At the lower end of X2P, the large enthalpy gain and the large entropy loss clearly compensate for one another resulting in a relatively low chemical potential loss. This compensation persists throughout the composition range measured and it contributes to the small change in chemical potential in the system. The enthalpic interaction function of 2-propanol, H 2 p 2 p E , was very similar to H i P i P E (Figure 3), suggesting the transition region between Mixing Scheme I and II for the binary aqueous 2- propanol solution starts from about X2p=0.05 and ends at about 0.11 with 0.08 being the nominal transition point ( 2 1 ), where on average, the HBN of water loses connectivity in the entire system. Based on this knowledge about the mixing behaviour of 2-propanol in the binary aqueous solution, a three-component system comprised of 1-propanol, water, and 2-propanol was studied. Our main objective was to determine whether the same behaviour 23 Excess Partial Molar Quantities of 2P in 2P-H20 system at 25°C 10 o LU CM CO A A A A LU CL CM LU CL CM •10 H •15 H • • • • • A • • 1 A • • • A -20 ~~T~ I I I I I I I I I 0.00 0.02 0.04 0.06 0.08 0.10 0.12 0.14 0.16 0.18 0.20 2P Figure 10: • Excess chemical potentials of 2P, n 2 P E • Excess partial molar enthalpies of 2P, H 2 P E ^ Excess partial molar entropies of 2P, S 2 P E at 25.00°C. 24 of 2-propanol could be recovered and observed through measuring the thermodynamic quantities of 1-propanol in this ternary system. Consequently, the excess partial molar enthalpies of 1-propanol and the total vapour pressure of IP-H2O-2P were measured. However, while the vapour pressure data were analysed as usual for any 3-component systems^17'18,50), it was realized that there was a ridge of maximum on the X1P-X2P field. This local maximum defies the numerical analysis based on the Gibbs-Duhem relation^ ' 2 6 \ Due to this problem, the excess chemical potentials of each of the component in this ternary system cannot be evaluated unless the ratio among IP, 2P, and water in the gas phase is determined. Currently, a new experimental method is being developed, whereby a gas chromatographic analysis device is incorporated within the vapour pressure measuring apparatus. The presence of the excess chemical potential data will complement the excess partial molar enthalpies, but until pipE and u.2PE are evaluated, only HipE will be examined. The excess partial molar enthalpies of 1-propanol, HipE, at various initial 2P compositions as a function of xjp are plotted in Figure 11. (The raw data are listed in Table A l in the appendix.) In the limit xi P-> 0, the excess partial molar enthalpies of infinite dilution of IP represent the enthalpic contribution of the very first 1-propanol molecule entering the system. In the first series, where the initial mole fraction of 2P is 0, the first IP molecule contributed about -lOkJ/mol to the system. As x 2p° increases, this enthalpic contribution becomes less, suggesting that 2-propanol has already contributed to the HBN enthalpically, prior to the addition of 1-propanol. The most striking appearance of this graph has to be the fact that series #5 (x2p° = 0.1260) does not follow the same trend found in the other series. In fact, the addition of 25 H 1 P E in 1P-H 20-2P systems of various x 2 P ° at 25°C 0 H - 2 - 4 H - 6 -8 - 10 •12 \ f * • ° v A D O A • _ s ° A A 0 ° ° ° A O O o I l I I ' I I 0 . 0 0 0 . 0 2 0 . 0 4 0 . 0 6 0 . 0 8 0 . 1 0 0 . 1 2 0 . 1 4 X 1 P F i g u r e 1 1 : E x c e s s partial molar enthalpies, H 1 P E in 1 P - H 2 0 - 2 P with different initial mole fractions of 2 P , x 2 P°. ° x 2 P ° = 0 , D x 2 p ° = 0.0140, A x 2 P ° = 0.0280, V x 2 p ° = 0.0500, O x 2 p ° = 0.1260. 26 1-propanol to this mixture of 2P-H2O was an endothermic process from the beginning. So why is this series so special? Recalling from above, the transition region for aqueous 2-propanol from Mixing Scheme I to II starts from about X2P=0.05 and ends at x2p=0.11. In other words, prior to the addition of IP, the solution of series #5, having a x 2p° of 0.1260, was already in Mixing Scheme II with water behaving as an ordinary liquid without its infinite connectivity. The other 4 series, all with a X 2 P ° < 0.05, would first be in region I and region II would come into place as x i P increases. This accounts for the completely different behaviour of series #5 as there was already enough 2-propanol to drive such a transition in the system. Focusing on the other 4 series, it is apparent that the dependence of HipE on X I P is similar for all X2p° except for parallel shifts to the left as X2p° increases. This phenomenon suggests that the presence of 2P has already modified the molecular organisation and the HBN of water to an extent that the new coming IP only modifies what was left by 2P in the same manner. To support this claim, H I P I P e (equation 1), were evaluated and plotted against X ] P in Figure 12. Four curves with similar shape were observed. By the same methodology as described in Introduction, the X M Y values for each curve were determined. These values represent the width and the locus of the transition region. Despite the uncertainties being as high as +1 OkJ/mol in H I P I P , two typical trends can be observed. First, as X 2 P ° increases, H I P I P E values start off higher, meaning the more 2P present, the more unfavourable enthalpic interaction among IP molecules. Second, the dependence of H I P I P E on X I P remains parallel except for an apparent shift to the left as X 2 P increases. Consequently, X M Y values also shift left (lower x i P values), suggesting the amount of IP required to drive the aqueous system from Mixing Scheme I to II is 27 H 1 P 1 P E in 1P-H 20-2P systems at various x 2 P ° at 25°C 0.00 0.02 0.04 0.06 0.08 0.10 0.12 0.14 Figure 12: Enthalpic 1-propanol -1-propanol interactions in 1P-H20-2P at various initial mole fraction of 2P, x2P°at 25.00°C. O x 2 p° = 0, • x 2 p° = 0.0140, A x 2 P° = 0.0280, V x 2 p° = 0.0500 Point X, M, and Y inidcate the region of Mixing Scheme Transition. 28 1P - 2P Mixing Scheme Transition Diagram 0 . 1 2 0 . 1 0 -i 0 . 0 0 0 . 0 2 0 . 0 4 0 . 0 6 0 . 0 8 0 . 1 0 X 1 P Figure 13: 1-propanol - 2-propanol Mixing Scheme transition boundary created by the transition points obtained from Figure 12. 29 less proportionally to the amount of 2P already present. To examine this point more clearly, a 1P-2P transition diagram is plotted in Figure 13. The Y-axis represents the initial composition of 2P and the X-axis shows the amount of IP required for a transition. Apparently, a straight line separating the two regions was observed. This suggests that 2P molecules must have modified the molecular organisation and the HBN of water in the same way as 1-propanol. Therefore, these two solutes work together cooperatively in an additive manner. This conclusion is certainly not surprising. Our binary aqueous IP and 2P systems have already shown the capability of the 2 solutes, individually, modifying the HBN of water in the same way. The fact that an expected conclusion is now reached provides confidence in the methodology presented in this thesis for the study of the mixing behaviour of a solid solute in a ternary aqueous system. 30 II) l-propanol-H2Q-NaCl A ternary system comprised of 1-propanoic water, and sodium chloride was considered in order to study the mixing behaviour of our third component and the interactions among the solutes through the medium of water. Sodium chloride in water, as being probably the most studied and famous aqueous electrolyte of all, is a well-known example of a structure breaker as described by Hepler ( 2 3 ). Breaking up the formation of ice-like structures in water is completely opposite to the behaviour of small hydrophobic molecules such as 1-propanol. (Numerous detailed studies(38"40) of NaCl have been published over the years and these will not be repeated here.) If the addition of IP results in the formation of "icebergs"(3"5) which further reduces the HB probability in bulk water, one would expect that such effects would be reduced or opposed by the counteraction of sodium chloride. Namely, the ability of IP to drive a transition from Mixing Scheme I to II in the solution should be restricted somewhat under the influence of NaCl. This hypothesis was quickly shown to be untrue when the excess partial molar enthalpies of IP in a three-component system involving NaCl were measured. Our study is limited to the water-rich and one phase region as shown in the phase diagram of the ternary system in Figure 14. The excess partial molar enthalpies IP at various XNaci° (initial mole fraction of NaCl) were measured and are plotted as a function of x i P in Figure 15, with uncertainties estimated to be +0.07kJ/mol. (The raw data can be found in Table A2 in the appendix.) According to Figure 15, it is observed that as x N a c i° increases, HipE basically shifts upward. This suggests that the enthalpic contribution of IP is reduced when the 31 NaCl - 1P Mixing Scheme and Phase Boundaries 0 . 1 2 0 . 1 0 0 . 0 8 0 . 0 6 0 . 0 4 0 . 0 2 0 . 0 0 0 . 0 0 0 . 0 2 0 . 0 4 0 . 0 6 0 . 0 8 0 . 1 0 0 . 1 2 X 1 P Figure 14: The Mixing Scheme and Phase boundaries of 1P-H20-NaCI system at 25.00°C. • Mixing Scheme Boundary A Phase Boundaries, S-L: solid-liquid phase, S-L-L : solid-liquid-liquid phase, L-L : liquid-liquid phase 32 H 1 P E in 1P-H20-NaCI at various x N a C | ° at 25°C o - 4 -6 H -8 10 H 12 7 V A D o v A ° O VoDo° V A V A • V v A n O A A n ncP r i i i i 0.00 0.02 0.04 0.06 0.08 0.10 0.12 X I P Figure 15: Excess partial molar enthalpies of 1-propanol in 1P-H20-NaCI systems at various initial mole fraction of NaCl, x N a C |° at 25.00°C. ° x N a c , ° = ° . D x N a C |° = 0.01092, A x N a C |°= 0.01962, V x N a C I° = 0.03503. 33 composition of NaCl increases. This observation is indeed identical as in the IP-H2O-2P system as described earlier. However, in contrast to the IP-H2O-2P system, the initial slopes of the curves in Figure 15 appear to be almost the same. When H I P I P E was evaluated by drawing smooth curves and taking the slope at intervals of 5xip = 0.002, the results are quite interesting. Figure 16 shows a plot of H I P I P E V S X I P . Lines are drawn to indicate the locus of X M Y values for the first series. It is clearly shown that the X M Y locus moves to the left as X N a c i ° becomes larger. This result was surprising because it suggests that it takes less I P to drive the transition from Mixing Scheme I to II with the help of NaCl. In other words, I P and NaCl work together cooperatively in removing the infinite connectivity of water's HBN. Does this mean the hydrophilic NaCl mixes with water in the same manner as the hydrophobic solutes such as I P and 2P? Not exactly, as is discussed below. Figure 16 shows that the enthalpic I P - I P interactions appear to have all originate from the exact same point when X I P = 0 , regardless of the initial sodium chloride composition. This suggests that in spite of different compositions of NaCl, the enthalpic interactions that are felt among the first few I P molecules are almost identical. Our interpretation can be summarized in the following points: 1) I P molecules dissolves into the bulk water away from the Na+Cl". 2) The I P - I P enthalpic interaction remains the same as in pure water. Thus the molecular organization in the bulk water away from the Na+Cl" ions remains the same as in pure water. 3) As NaCl increases, the nominal transition point M occurs progressively at a smaller value of X I P , indicating Na+Cl" binds to a certain amount of H2O, making it unavailable for I P to dissolve into, but leaving the bulk water away from Na+Cl" almost the same as pure water. However, this claim of ours 34 H 1 P 1 P E in 1P-H20-NaCI systems of various x N a C I ° at 25°C ' I P IP NaCl 250 200 H 150 o E ^ 100 CL T -• Q. 50 0 0.00 0.02 0.04 0.06 0.08 X IP l i r 0.10 0.12 0.14 0.16 Figure 16: Enthalpic 1-propanol -1-propanol interaction in 1P-H20-NaCI systems at various initial mole fraction of NaCl, x N a C I° at 25.00°C. O x ° = 0, NaCl • xNaC,° = 0.01092, A x N a C |° = 0.01962, V x N a C I° = 0.03503. 35 is not yet conclusive due to the lack of supporting evidence from the excess partial molar entropy data. The total vapour pressure of this 3-component system was measured. Once again, the same problem with the numerical analysis of excess chemical potentials was encountered as in the IP-H2O-2P system. Once the vapour pressure measuring apparatus is improved such that GC analysis can be performed at the same time, the u-iPE can be evaluated and hence the S I P E values. So what role does NaCl play in this ternary system? When the nominal transition points were compared among the 4 series, it was determined that as XNaci° was increased from 0 to 0.035, point M moves from xiP=0.062 to XIP=0.044. In terms of number of molecules, this translates to a ratio of 15 H2O : 1 I P for XIP=0.062 at XN a c i ° = 0 and a ratio of 21 H 2 0 : 1 I P : 0.8 NaCl for x i P = 0.044 at xNaci o =0.035. Since this ratio is taken at point M where the transition occurs as an average, (21-15)= 6 water molecules become unavailable for 1-propanol due to the presence of this 0.8 molecules NaCl. It follows that (6/0.8)=7.5 H2O molecules on average are being "held hostage" by each pair of Na+Cl" ions and not being "freed" for iceberg formation by the I P solutes. Hence the amount of bulk water molecules available for I P to serve its purpose (ie. to form iceberg structure and then modify the H B N of water) is reduced and thus less I P is required. In a recent study of gas phase clusters of X'(H20)mtype molecules (X" = Cl", Br", I") by pre-dissociation vibrational spectroscopy^''42), when m = 4 or 5 (depending on the anion) the spectra become almost identical to that of pure liquid H2O. This result supports our hypothesis that one pair of Na+Cl" ions retains 7.5 molecules of H2O such that the bulk water shows the same mixing behaviour as pure liquid water. 36 Another interesting finding from our results is that the 7.5:1 ratio could be closely related to the 9:1 ratio of the solubility of NaCl in water as explained in the following. In figure 14, the boundary for Mixing Scheme transition is also shown. As described in the Introduction, the idea is that the two qualitatively different Mixing Scheme regions seem to serve two distinct purposes. In region I, the molecules in the solution are organized in a way that addition compounds(32) can be formed at a low enough temperature. In region II, they are organized to anticipate a phase separation at a high enough temperature. If a straight line is drawn through the Mixing Scheme transition points toward the Y-axis, not only does it merge into the phase boundary line between the one-phase region and the liquid-liquid 2 phase region on the phase diagram, it also intersects the Y-axis at XNaCi° =0-12, which translates into a 7.5 :1 ratio. Although not seen in a real situation, this suggests that when there is no IP present, the mixing scheme transition would occur at a ratio of 1 NaCl to 7.5 water. This suggests that at saturation XNaci(sat) = 0.1, the rest of the (9 -1.5) 1.5 "extra" bulk water would not have a high enough HB probability to have the whole network connected infinitely. At the smaller values of XNaci° < XNaci(sat), the same 7.5:1 ratio would still persist. But the remaining water molecules would have a high HB probability to retain percolation until enough IP enters the system and reduces the HB probability below the percolation threshold and thus switch into region II. This interpretation, though consistent with the results, is not conclusive. Some other factors, such as ionic dissociation of NaCl in a ternary system, may have a great influence on our elucidation of the aqueous system. In addition, excess partial molar entropy data and possibly, excess partial molar volume data are needed for a more comprehensive discussion. 37 Ill) l-propanol-BUO-urea First discovered by H.M.Roulle in 1773, urea is a colourless crystalline solid, found mostly in urine of humans and with smaller quantities in blood, liver, lymph, and serous fluid. It is produced in the liver as an end product of protein metabolism. Due to its high nitrogen content, commercially prepared urea is used as agricultural fertilizers^ \ The most profound biochemical awareness of urea has to be its denaturing effect on proteins. Many believe that the mechanism behind this effect is the change of water's HBN induced by urea which in turn changes the solubility and the hydrophobic effect of the protein (35). In a recent paper by Vanzi et al . ( 3 5 ) , it was determined that urea increases the solubility of hydrocarbons with chains longer than 2 carbons. They inconclusively proposed two possible mechanisms: 1) urea changes the HBN of water and helps the hydration of hydrocarbons. 2) both urea and water solvate the hydrocarbon. Specifically, urea hydrogen bonds with the protein peptide groups. To add deeper insights into this problem, our thermodynamic method seems to be a legitimate approach. A ternary system involving urea, water, and 1-propanol was examined and the mixing behaviour of urea and 1 -propanol was elucidated. Since our experimental apparatus was not designed for titrating solid into solution, the excess partial molar enthalpies of urea cannot be measured directly in a binary system. Instead, the excess partial molar enthalpies of IP were measured in the presence of different initial compositions of urea and they are plotted in Figure 17. The 38 H 1 P E in 1 P-H 2 0-Urea systems of various x u r e a° at 25°C o - 2 H - 6 -8 •10 0 . 0 0 0 . 0 2 0 . 0 4 0 . 0 6 0 . 0 8 0 . 1 0 0 . 1 2 0 . 1 4 Figure 17: Excess partial molar enthalpies of 1-propanol in 1 P-H20-Urea systems of various initial mole fraction of urea, x u r e a° at 25°C. Each injection of 1P was 0.6869+0.005mL. O V • 0 _ urea 0.03040, A x ° = 0.04985, ° = 0.1020, O x „ ° = 0.1501, 0 x 0 = 0.2035. urea 39 raw data can be found in Table A3 of the appendix. The uncertainty was estimated to be +0.1kJ/mol. Six series were carried out. As the initial mole fraction of urea increases from 0 to 0.2035, the enthalpic contribution of IP becomes much less (from about -lOkJ/mol to —4kJ/mol). The curvature of each series is quite different. It is striking that all 6 curves overlap and crossover at a single point at XJP=0.06. After passing through this point, their relative positions are completely reversed. This kind of excess partial molar enthalpic behaviour was never observed in any three-component system studied so far. One can certainly sense that there is something unique about the way urea mixes with water and IP. The enthalpic IP-IP interaction, H I P I P E , were evaluated by taking the slope of each curve at intervals of 8XIP=0.002 and they are plotted in Figure 18. In this enthalpic interaction plot, as x u r e a 0 increases, the overall HjpipE significantly decreases, suggesting that the intermolecular interactions among IP molecules are reduced when there are more urea molecules in the system. The effect that most attracts our attention is that the peaks of each curve of H I P I P e appear to be located at almost the same position even if the uncertainties are taken into account. In addition, when the X M Y values of each curve were evaluated, it was determined that although the values of X and Y, in terms of X I P , vary in some small extent, M , the nominal transition point, is the same for each series. This further suggests that although the presence of urea lowers the enthalpic interactions among IP's, it has no influence on the locus of the Mixing Scheme transition(13). The measured total vapour pressure data of lP-H20-urea mixture corrected for 25.00°C are listed in Tables A4 and A5 of the appendix. There was no local extremum in 40 1P1P in 1 P - H 2 0 - U r e a systems of various x u r e a° at 25°C 200 1 5 0 H 1 0 0 H 5 0 H 0 0 . 0 0 IP Figure 18: Enthalpic 1-propanol -1-propanol interaction in 1P-H20-Urea systems of various x u r e a° at 25.00°C. Points X, M and Y represents the region of Mixing Scheme Transition. O x 0 = 0 A urea w ' • x , ° = 0.03040, A x °= 0.04985, urea ' urea ' V x ° = 0.1020, O 0 _ 0.1501, O x, ° = 0.2035. 41 the total pressure. Thus, a numerical method (18) was used to evaluate the partial pressures. The excess chemical potentials of urea and that of 1 -propanol were than calculated and (see appendix for numerical analysis) plotted in Figure 19 and 20, respectively. Both figures indicate that in this particular ternary system, u.E (i being either urea or IP) decreases on increasing X J . This further indicates that the interaction in terms of chemical potential, pj E , is negative, exhibiting a net attraction in chemical potential between the solutes. The data points in Figure 20 for the 4 series performed overlap one another which suggests the effect of the 3r d component (urea) on the excess chemical potentials of IP is almost negligible(45). p i P E remains unchanged upon the increase of urea concentration, once again suggesting the special mixing behaviour of urea, which is quite different from IP and other hydrophobic solutes. From the graphically interpolated values of H I P E and the values of u,ipE, the excess partial molar entropies of 1-propanol, S i P E , were calculated. The values of T S I P E are listed in Table A6 of the appendix and are plotted in Figure 21. This plot and Figure 17 show great resemblances. When the corresponding T S I P I P E were evaluated, an extremely similar x ] P dependence was observed as in the H I P I P E plot. Figure 22 , although with more scattered data points, still manages to exhibit the same effect as in the H I P I P E plot. One more supporting piece of evidence for this claim is the plot of IP-urea interactions in Figure 23. H I P . T J E decreases upon the increase of urea and the curvature changes suddenly at about the boundary between region I and II, where this locus does not change on increasing urea content. Figure 24 provides a graphical illustration of the loci of transition. All of the above findings lead us to ponder about the role of urea in aqueous solutions. The question of whether urea acts as a structure breaker(23) (like NaCl) has 42 urea in 1 P - H 2 0 - U r e a of various x 1 p at 25 C o E —> HI 4 . 4 -4 . 2 4 . 0 3 . 8 H TO 5 3 . 6 i t 3 . 4 3 . 2 H 3 . 0 0 . 0 0 0 . 0 5 0 . 1 0 0 . 1 5 0 . 2 0 vUrea Figure 19: Excess chemical potentials of urea in 1P-H20-Urea systems of various mole fraction of 1-propanol, x 1 p at 25.00°C. O X1p=0.02, O X1p=0.03, • X1p=0.04, V X1P=0.05, O X1P=0.07, A X1P=0.09, O X1P=0.11. 43 | n 1 p E in 1P-H 20-Urea of various x u r e a at 25°C o E -> LU Q. i t 7 . 0 6 . 5 -A 6 . 0 5 . 5 5 . 0 4 . 5 H 4 . 0 0 . 0 0 0 . 0 2 0 . 0 4 0 . 0 6 0 . 0 8 0 . 1 0 0 . 1 2 IP Figure 20: Excess chemical potentials of 1-propanol in 1P-H20-Urea systems of various mole fraction of 1-propanol, x1p, at 25.00°C. O *urea= 0-07, O V xu r e a = 0.19. x =011 A A urea u - ' 1 • ^ X urea = 0 - 1 5 ' 44 TS 1 P E in 1P-H20-Urea of various x u r e a at 25°C o E -8 H i -10 UJ Q. tn H- -12 •14 H -16 H •18 0.00 0.02 0.04 0.06 x 1 P 0.08 0.10 0.12 Figure 21: Excess partial molar entropies of 1-propanol in 1P-H20-Urea systems of various mole fraction of urea, xurea, at 25.00°C. The temperature term is included to keep the units of excess partial molar entropies the same to that of excess partial molar enthalpies. O O * u r e a = 0.15. X u r e a = ° . A X u r e a = 0 0 5 . V X urea = 0 - 1 . 45 TS 1 P 1 P E in 1P-H20-Urea of various x u r e a at 25°C 200 ~ 150 o E XL UJ Q . 0. to" 100 H 50 0 /°°\ / • \ Error / n ° 9 A ° \ / • A A b A A A X / A A \ \ 0 6 \ A • 0.00 0.02 0.04 0.06 0.08 0.10 0.12 0.14 K 1 P Figure 22! Entropic 1-propanol - 1-propanol interactions in1P-H 20-Urea systems of various mole fraction of urea, x u r e a , at 25.00°C. Points X, Y, M represent the region of Mixing Scheme Transition. ° X u r e a = 0 - 0 1 5 . D > < u r e a = 0 0 4 0 . A X u r e a = ° - 1 2 5 -46 Enthapic and Entropic Interaction between 1P and Urea 60 o E —> x. UJ => t Q . co" UJ I a. 50 .40 30 20 10 0 •10 --20 0.00 0.02 0.04 0.06 0.08 0.10 0.12 k1P Figure 23: Interaction functions between 1-propanol and urea in 1P-H20-Urea systems of various mole fraction of urea, x u r e a at 25°C. H 1 P _ u r e a E is shown by shaded symbols and T S 1 p . u r e a E is shown by black symbols. X urea = 0-015, -urea = 0-125. x = 0.040, urea ' A x =0.075, urea ' 47 Mixing Scheme Transistion Boundary in 1P-H20-Urea 0 . 2 5 0 . 2 0 -\ 0 . 1 5 CD 0 . 1 0 0 . 0 5 H 0 . 0 0 0 . 0 0 0 . 0 2 0 . 0 4 0 . 0 6 0 . 0 8 0 . 1 0 0 . 1 2 1P Figure 24: Mixing Scheme Transition Boundary in 1P-H20-Urea systems at 25.00°C. The left side of each locus represents point X; the midpoint represents point M; the right side represents point Y. 48 been raised. Frank and Franks ( 4 6 ) once suggested that the water molecules around urea are less hydrogen bonded than in bulk water. But other evidences showed that the H B N of water is maintained or even enhanced around the amino groups, while it is distorted by the carbonyl carbon ( 4 7 ). When the structure of urea is examined, it is reasonable to believe that both the carbonyl oxygen and the amine hydrogen provide acceptable sites for hydrogen acceptor and donor bonds. This implies that urea could readily insert itself into the HBN of water and explains the ability of large amounts of urea to dissolve in water with minimum disruption of the overall hydrogen bonding of aqueous solutions. From all these information and the experimental results, urea may not be either a structure breaker or a maker. Since the HB between urea and water is so similar to those between water molecules, when urea molecules dissolve in water, they are locked into the HBN of H2O by replacing some of the water molecules. So they do not favour the loss of overall HB probability, nor do they prevent this from occurring. Their indifference towards the modification of the HBN of water is because they have become a part of the HBN. Hence, with or without the presence of urea, it requires the same the amount of 1 -propanol to drive the H B N system to Mixing Scheme I I , where there is no infinite HBN connectivity. This suggests that urea does not alter the nature of the connectivity, and perhaps the HB probability, of H2O. However, there is a difference between the hydrogen bonds between water molecules and those between urea and water. This difference is significant enough to be detected by our solute-solute interaction functions (equation 2). The decrease of H I P I P E on increasing urea suggests the reduction of interaction among IP molecules. Since our "interaction" is a water mediated process in Mixing Scheme I , if urea is inserted into the whole HBN, the degree of fluctuation of the HBN would 49 decrease due to the more rigid structure of urea. So i f a high degree of fluctuating H B N which exists in pure H2O is a required element for H2O mediated interaction, increasing urea content would diminish the IP- IP interaction. The close relation between the degree of fluctuation and the strength of solute-solute interaction is suggested in a recent paper by Tamura et. a l . ( 4 8 ) . To answer Vanzi ' s question posed at the beginning of this section, our results seem to agree with his first idea that urea changes the H B N o f water and affects the hydration of hydrocarbons ( 3 5 ). The change on H B N by urea is on the fluctuation but not the connectivity, or the H B probability. In other words, urea does not interact with the hydrophobic solute directly as in his second proposed mechanism. Instead, urea changes the structure of H B N around the hydrophobic groups in protein, thereby increasing their solubility and weakening the'hydrophobic effect. This in turn destabilizes the native state of protein and induces a conformational change in the protein and denaturation results. 50 IV) 1-propanol - H,Q - B-D-Fructose The excess partial molar enthalpies of 1-propanol were measured in a ternary system comprised of 1-propanol, H2O, and P-D-fructose (called fructose from now on) in order to study the interactions among 1 -propanol, water, and fructose, the mixing behaviour of fructose, and the influence of fructose on the molecular organization of water in the aqueous system. Figure 25 is a plot of H I P E vs xjp of various initial mole fractions of fructose at 25°C. Once again, the trends developed were in our familiar territory. As the initial amount of fructose increases, the enthalpic contribution of IP to the entire system is decreased accordingly. This behaviour is similar to what we have already observed in the cases of the IP-H2O-2P and lP-HbO-NaCl systems. Initial composition of fructose was used as high as 0.0633 (i.e. 67.6g fructose in lOOg H2O) in order to investigate whether any special changes would occur at higher fructose content. The results indicate that the same trend persists even at such high fructose composition. As in the three previous ternary systems, the enthalpic solute-solute interaction functions, H I P I P E were evaluated by taking the slope of Figure 25 at small intervals of X I P and the results are plotted in Figure 26. A familiar peak shaped 3 r d derivative of G anomaly was again observed in this system. As X F I - U C 0 increases, the position of the peaks, along with the locus of the X M Y values, moves to the left. Concurrently, the overall H I P I P E values are reduced. This pattern was not observed in the three previous systems. In the IP-H2O-2P system, Hi P I pE values at the peaks are identical but the values at X I P = 0 was progressively less. In thelP-H20-NaCl system, not only the peak values but also the initial H I P I P E values are the same. Finally in the lP-H20-urea system, the peaks locate at the same locus in terms of xip. This, undoubtedly, hints at a different mixing behaviour 51 H 1 P E in 1 P-H 20-Fruc of various x ^ 0 at 25°C o -4 -8 H -10 U . U . U 7 . n 7 / t o - , • A • • o o v v A D D o O O V A • 0 0 O V A O 0 0 0 V A Q ° 0 V A o A ° v A " ° $ V o o 0 0 A O V A D O V o V A V A O A O V A • A ° • • O • • ° • D C O o o o o I I I I I 0.0.0 0.02 0.04 0.06 0.08 0.10 0.12 X 1 P Figure 25: Excess partial molar enthalpies of 1-propanol in 1P-H 2 0-Fruc systems of various initial mole fraction of p-D-fructose, x ^ 0 at 25.00°C. Each injection contained 0.6869+0.005mL of 1-propanol. ° x f r u c ° = 0, D x f r u c ° = 0.0149, * x f r u c ° = 0.0315, v X f r u c ° = 0.0509, O x f m c ° = 0.0579, o x f r u c ° = 0.0633. H 1 P 1 P E in 1P-H20-Fruc of various x f r u c ° at 25°C o E ID Q . CL 2 0 0 1 5 0 A 1 0 0 5 0 A 0 o o • A no A o V A oco% -<> o A Q Error 0 . 0 0 0 . 0 2 0 . 0 4 0 . 0 6 0 . 0 8 0 . 1 0 0 . 1 2 k 1 P F i g u r e 26: E n t h a l p i c 1 -propano l - 1 -propano l interact ion in 1 P - H 2 0 - F r u c s y s t e m s of v a r i o u s initial m o l e f ract ion of p -D - f ruc to se , x, at 2 5 . 0 0 ° C . O true >c ° = 0 • x •^ fruc w ' x 1 A o x f r u c ° = 0 .0315, V x f r u c ° = 0.0579, ° fruo° = 0 -0149, 0 = 0 .0509, ^fruc xfruc° = 0 -0633. 53 of fructose from the other three solutes mentioned. The presence of fructose not only brings the M i x i n g Scheme transition to an earlier stage, it also weakens the IP- IP via-water interactions ( 8" 1 2 ) in terms of enthalpy. The phase and M i x i n g Scheme diagram for the lP-H20 -Fruc system is shown in Figure 27a. The phase boundary was determined by turbidity titration. For the M i x i n g Scheme boundary, a non-linear curve connecting the two axes suggests that the two solutes work together cooperatively but non-additively in modifying the H B N o f water. Namely, as the aqueous system initially contains more fructose, it takes less 1-propanol to drive a transition from M i x i n g Scheme I to M i x i n g Scheme II. But the degree of influence on the H B N of water by the fructose is somewhat different since the M i x i n g Scheme boundary is a curve rather than a straight line as in the case o f IP-H2O-2P as shown in Figure 27b. 2-propanol has a structural similarity to 1-propanol. One might expect their degree of influence on the H B N o f water be, i f not the same, very similar. However, in the IP-H2O-Fruc case, the shape of the H I P I P E curves are different from those of IP-H2O-2P, - N a C l , and -Urea . Their different molecular structures should have already suggested that their individual mixing behaviour would not be the same. The above suggestions were further supported when the total vapour pressures of the lP-H20-fructose systems were measured. However, the same problem with the numerical analysis of the partial pressures was initially encountered as in the IP-H2O-2P and l P - H 2 0 - N a C l systems. But for this system, a preliminary G C analysis was performed and the results were used to calculate the partial pressures of IP and H2O. The excess 54 Mixing Scheme and Phase transition of 1 P-H20-Fruc 0 . 1 6 0 . 1 4 0 . 1 2 H 0 . 1 0 E 0 . 0 8 0 . 0 6 0 . 0 4 0 . 0 2 0 . 0 0 0 . 0 0 0 . 0 5 0 . 1 0 0 . 1 5 i i r 0 . 2 0 0 . 2 5 0 . 3 0 0 . 3 5 0 . 4 0 ~1P Figure 27a: The Mixing Scheme boundary • and the Phase transition boundary • between 1-propanol and P-D-fructose shown by the initial mole fraction of p-D-fructose, x f r u c ° a 9 a i n s t t r , e m o ' e fraction of 1-propanol, x1p. ^ is from reference (49). 55 Mixing Scheme Transition Boundary in 1P-H20-fruc at 25°C 0 . 1 2 -, 1 0 . 0 0 0 .01 0 . 0 2 0 . 0 3 0 . 0 4 0 . 0 5 0 . 0 6 0 . 0 7 0 . 0 8 0 . 0 9 ~1P Figure 27b: Mixing Scheme Transition boundary in 1P-H20-Fruc systems at 25°C. The point on the Y axis is from literature ( 4 9 ). 56 chemical potentials of IP, pipE, were evaluated and are plotted against X I P in Figure 28. It was observed that the pipE of all 6 series decrease smoothly as x i P increases. Also, u.ipE increases with XFruC°- T S I P E values were calculated using U I P E and H I P E and are plotted in Figure 29. A greater fructose content resulted in a less entropy loss by 1-propanol suggesting less randomness in the solution as more ordered molecular organizations(3"5) have been formed in water surrounding the fructose molecules. So most of the entropy loss was involved with fructose such that the entropic contribution by 1-propanol became less. When the entropic interactions, T S I P I P E were evaluated by taking a derivative with respect to X I P , the same anomaly as appeared in the enthalpy case was recovered. Figure 30 is a plot of T S I P I P E against xip. The same trends as observed in H I P I P E are also exhibited. The exact same Mixing Scheme transition loci as from the enthalpy data were determined. Returning to Figure 27a, the Mixing Scheme transition boundary seems to move towards the Y-axis. Unfortunately, for the binary aqueous fructose, we were not able to determine the transition point nor to see i f any transition exists at all. As explained in the Introduction, our apparatus was not designed for titrating solids and the noise for titrating a solid-liquid mixture into water was extremely high. Galena et. a l . ( 4 9 ) suggested the hydration number of P-D-fructose was 8.8 (by density and ultrasound measurements), equivalent to X f m c = 0.100. If this is true, as shown Figure 27a, the Mixing Scheme boundary seems to extrapolate smoothly (with a curvature) to the point of x ^ 0 = 0.10. Whether this is purely coincidence is uncertain. A more sensitive microcalorimetre has been built in our laboratory and is currently being calibrated. The Mixing Scheme transition point of binary fructose H2O will be determined in the future and the transition 57 E 0 ILi.jp in 1 P-H 20-Fruc of various x f r u c at 25 C 9.0 8.5 8.0 V 7 - 5 E 3 7.0 LJJ Q_ £ " 6.5 6.0 -5.5 -5.0 O o v 0 o v v A C A o • o 0 ^ O 0 V X7 O 0 v o V O A V A 0 A 9 o v O V A O A D 9 V A O A • 8 S v v V A 9 A B 0 V V 0 A V e A A O S T T T T 0.00 0.01 0.02 0.03 0.04 0.05 0.06 0.07 0.08 X 1 P Figure 28: Excess chemical potentials of 1-propanol in 1P-H20-Fruc systems of various mole fraction of p-D-fructose, at 25.00°C. u xfruc=0, U x f r u c = 0.015, A X f r u c = 0.030, V x f r u c = 0.055, O x f m c = 0.060, 0 x ^ = 0.065. T S 1 P E in 1 P-H20-Fruc of various x f r u c ° at 25°C o £ Q. LU CO I-0 . 0 0 0 .01 0 . 0 2 0 . 0 3 0 . 0 4 0 . 0 5 0 . 0 6 0 . 0 7 0 . 0 8 L 1 P Figure 29: Excess partial molar entropies of 1-propanol in 1P-H20-Fruc systems of various initial mole fraction of p-D-fructose, x f r u c° at 25.00°C. The temperature term is included to keep the units of excess partial molar enthalpy and those of excess partial molar entropy the same.O = 0, ^ x f r uc° = 0.0509, ° = 0.0149, A x ° = 0.0315, O x ° = 0.0579, • O ~fruc >W0 = 0-0633. 59 T S 1 P 1 P E in 1P-H 20-Fruc of various x f r u c ° at 25°C o E LU 2 4 0 2 2 0 2 0 0 1 8 0 1 6 0 1 4 0 1 2 0 1 0 0 8 0 6 0 4 0 2 0 O o o o o • • e ° A D • v A D N V S 6 V y A • Error o o 9 g 0 A o ° O A A A o o A • o O V • 0 V A • 0 A ft 7> A 0 V V 0 0 o o o • o 0 I I I : I I I I 0 . 0 0 0 .01 0 . 0 2 0 . 0 3 0 . 0 4 0 . 0 5 0 . 0 6 0 . 0 7 0 . 0 8 K 1 P Figure 30: Entropic 1-propanol - 1-propanol interaction in 1P-H20-Fruc systems of various initial mole fraction of (3-D-frcutose, x f r u c° at25.00°C. O x f r u c° = 0, • x ^ 0 = 0.0149, V X f r u c ° = 0.0509, O x f r u c° = 0.0579, 0 x ° = 0.0633. ^fruc 60 boundary will be determined if it exists. The thermodynamic contribution of fructose and that of glycerol in aqueous solution show some resemblance. Figures 31, 32, and 33 show the H I P E , H I P I P E , and the transition boundary between glycerol and IP, respectively(50). The experimental results for glycerol and fructose are very similar. As a matter of fact, one of the reasons fructose was selected as part of this project is the fact that glycerol was previously studied and the structures of the two solutes are closely related. The structure of glycerol contains a backbone of 3 carbons and there is a hydroxyl group on each of the carbons. Similarly, fructose also possesses 6 carbons and all but one of them is attached to a hydroxyl group. Carbon number 2 (Fischer's numbering system) is attached to a keto-oxygen. In addition, the 6 carbons form a 5-member or 6-member ring. For glycerol, whether the hydrophobic carbon backbone is dominant over the hydrophilic OH groups or vice versa is yet to be determined. But the present experimental results suggest that glycerol is a possible structure maker similar to 1-propanol since there is a non-additive but positive cooperativity between glycerol and 1-propanol in modifying the HBN of water. Therefore, glycerol can be considered as a link of three short alcohols (methanol, in particular). In the case of fructose, the 5 or 6-member ring would make it appear that its structure is more hydrophobic than normal acyclic alcohols and we would expect it to hydrate less comfortably in water. But the fact that it dissolves in water so easily already proves that it does not create a lot of disruption in water. Many hydration studies of sugars had been carried out in the last few years ( 4 9 , 5 1" 5 7 ) but no definite conclusions have 61 H 1 P E in 1P-H 20-Gly of various x , ° at 25°C 1 0 H 0 . 0 0 0 . 0 5 0 . 1 0 0 . 1 5 0 . 2 0 k1P Figure 31: Excess partial molar enthalpies of 1-propanol in 1P-H20-Gly systems of various initial mole fractions of glycerol, xgly° at 25.00°C. Each injection contained 0.6869+0.005ml_ 1-propanol. O V 0 _ vgiy = 0, • x ° = 0.024, A x o _ gly gly = 0.50, 0 _ 0.085. H 1 P 1 P E in 1P-H 20-Gly of various x a l v ° at 25°C giy 2 5 0 2 0 0 1 5 0 - A V 1 o o E 1 0 0 -LU A ^ 7 Q_ n 5 0 -0 -- 5 0 o o o • o • A A • 7 A • O • A O Error O 0 0 . 0 0 0 . 0 5 I 0 . 1 0 X 1P 0 . 1 5 0 . 2 0 Figure 32: Enthalpic 1-propanol - 1-propanol interactions in 1P-H20-Gly systems of various initial mole fractions of glycerol, x, at 25.00°C. O 0 _ A 0 _ vgiy 0.050, Ngiy V = 0, • giy xg ly° = 0.024, 0 _ vgiy = 0.085. 63 Mixing Scheme Transition Boundary of 1P-H 20-Gly 64 been drawn. But most have agreed that a 6-member ring hydrates better than a 5-member ring, different stereoisomers have no different thermodynamic consequences but changing the ratio of axial to equatorial OH's changes many of its physical properties. Our experimental approach allows us to study the mixing action of fructose in a 3-component system. The results contain a lot of useful information about the influence of fructose on the other solute-solute interaction and on the molecular interaction as a whole. However, the details of interaction between fructose and water were not yet clear. Questions such as whether fructose is a structure breaker or maker still need to be answered. Other additional data such as excess partial volume, isobaric thermal expansivity, and entropy-volume cross fluctuation would fill in some of the gaps that still exist at present time. 65 C O N C L U S I O N The "third components" selected in this research project were each unique in some way. The mixing behaviour of 2-propanol, NaCl, urea, and fructose within an aqueous solution are expected to be different. Our methodology of employing thermodynamic quantities proportional to the second and third derivatives of G, and of using the thermodynamic behaviour of 1-propanol as a probe for studying the effects of a third component on water appears successful as demonstrated in this thesis. In the 1P-H2O-2P mixtures, the enthalpic interaction functions of 1-propanol, H I P I P E , for various initial composition of 2-propanol show that 2-propanol modifies the hydrogen network of water cooperatively and additively with 1-propanol. The fact that 2-propanol affects the molecular organization of water in the same manner as 1-propanol has been already determined in studies of a binary aqueous 2P system. This finding of an additive effect by IP and 2P from the present work provides confidence in the methodology we employed. In the lP-H20-NaCl system, the enthalpic interaction functions show that each pair of Na+Cl" binds to 7.5 water molecules, which are made unavailable for 1-propanol to interact with. At the same time, the thermodynamic behaviour of 1-propanol indicates that the remaining bulk water is the same as in pure water. Consequently, with the presence of sodium chloride, a lesser amount of 1-propanol is required to modify the molecular organization of water to the point where the hydrogen bond connectivity is destroyed. In the lP-F^O-Urea system, our results suggest that urea is non-cooperative with 1-propanol in the modification of the HBN of water. The amount of 1-propanol required to drive the Mixing Scheme transition remains the same with various 66 composition of urea in the mixture. However, the enthalpic and entropic interaction functions among 1 -propanol molecules are reduced progressively as the urea concentration is increased. This suggests that urea molecules are locked into the HBN of water keeping the HBN connectivity unchanged, but it does have the effect of reducing the fluctuation of the HBN. Thus, the water-mediated solute-solute interaction is weakened. In the final ternary system, lP-H20-Fruc, the enthalpic and entropic interaction functions indicate that fructose and 1-propanol work cooperatively to reduce the connectivity of the HBN of water. The loci of the Mixing Scheme transitions move towards smaller values of X I P , as in the NaCl and 2P systems. However, the values of the I P - I P interactions are reduced as in the urea system. It is interesting to observe that three Mixing Scheme transition diagrams (Figures 13, 14, 27b) of the four systems studied show a line (linear or curved) joining the two axes. Urea is the only exception where a vertical line is observed. The Mixing Scheme transition diagram of urea (Figure 24) suggests that there is no transition region for the binary aqueous urea system. Is it possible not to have a Mixing Scheme I for any binary aqueous mixture where Mixing Scheme II is operative even at infinite dilution? Could this phenomenon be related to the hydrogen bonding capability of urea? These questions, in addition to many others posed in the Discussion, cannot be answered at present but rather lead to ideas for future research. Overall, the process of data acquisition in the experiments conducted in this research requires the precision and the accuracy of good instruments, plenty of patience by the experimentalist, and it is undoubtedly time consuming. However, the results often bring forth very informative and compelling evidences of many of the physical 67 phenomena in aqueous solution that are not able to be measured or detected by other means. 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Data 1990, 35, 41-43. 54) Barone, G.; Cacace, P.; Castronuovo, G.; Elia, V. Carbohydrate research, 1981, 91, 101-111. 71 APPENDIX Table A l : Excess partial molar enthalpies of 1-propanol, HipE, in IP-H2O-2P systems at 25°C. There were 5 systems with X2p° varied, X I P is the mole fraction of 1 -propanol. X2p° is the initial mole fraction of 2-propanol. = 0\ X 2 p° =0.0140 X2p° '¥=0.028 X 2 p° . =0.0500 X2P° =0.1260 Xip HIP kJ/mol Xip Hjp kJ/mol Xip Hip k J/mol Xip Hip kJ/mol Xip H | P kJ/mol 0.004 -10.00 0.0011 -9.06 0.0012 -6.68 0.002 -2.94 0.0042 2.73 0.008 -9.51 0.0034 -8.78 0.0036 -6.75 0.004 -2.70 0.0070 2.81 0.012 -9.11 0.0056 -8.42 0.0059 -6.35 0.006 -2.39 0.0098 2.89 0.016 -8.64 0.0100 -8.26 0.0083 -5.938 0.008 -1.98 0.0126 2.77 0.020 -8.14 0.0122 -7.41 0.0129 -5.06 0.010 -1.60 0.0153 2.71 0.024 -7.63 0.0144 -7.23 0.0153 -4.59 0.012 -1.26 0.0180 2.66 0.028 -7.04 0.0166 -6.87 0.0199 -3.71 0.014 -0.94 0.0208 2.63 0.032 -6.39 0.0187 -6.45 0.0222 -3.28 0.016 -0.63 0.0235 2.52 0.036 -5.68 0.0209 -5.91 0.0244. -2.70 0.018 -0.31 0.0262 2.54 0.040 -4.91 0.0252 -5.32 0.0267 -2.37 0.020 -0.02 0.0288 2.37 0.044 -4.12 0.0273 -4.91 0.0289 -1.93 0.022 0.21 0.0315 2.29 0.048 -3.22 0.0294 -4.48 0.0312 -1.54 0.024 0.61 0.0341 2.09 0.052 -2.54 0.0315 -4.10 0.0334 -1.21 0.026 0.81 0.0368 2.12 0.056 -1.82 0.0336 -3.65 0.0356 -0.83 . 0.028 0.98 0.0394 2.00 0.060 -1.20 0.0357 -3.24 0.0401 -0.10 0.030 1.10 0.0420 2.02 0.064 -0.73 0.0378 -2.83 0.032 1.20 0.0446 1.93 0.068 -0.34 0.0399 -2.31 0.0444 0.364 0.034 1.27 0.0471 1.90 0.072 -0.04 0.0420 -2.50 0.0466 0.474 0.036 1.34 0.0523 1.78 0.076 0.02 0.0440 -1.67 0.0488 0.554 0.038 1.39 0.0573 1.91 0.080 0.33 0.0461 -1.33 0.0509 0.752 0.040 1.43 0.084 0.46 0.0501 -0.76 0.0531 0.839 0.042 1.46 0.088 0.55 0.0521 -0.50 0.0552 0.920 0.044 1.45 0.092 0.63 0.0562 -0.08 0.046 1.49 0.096 0.70 0.0582 0.12 0.048 1.49 0.100 0.74 0.0601 0.21 0.050 1.50 0.104 0.77 0.0621 0.33 0.052 1.48 0.108 0.79 0.0641 0.45 0.054 1.43 0.112 0.83 0.0661 0.56 0.056 1.43 0.116 0.85 0.0680 0.63 0.058 1.44 0.120 0.86 0.0700 0.66 0.124 0.88 0.0719 0.79 0.128 0.89 0.0738 0.77 0.132 0.90 0.0754 0.84 0.136 0.91 0.0796 0.0815 0.0834 0.0852 0.91 1.00 0.91 0.95 72 Table A2: Excess partial molar enthalpies of 1-propanol, HipE, in lP-FbO-NaCl systems at 25°C. 4 series, each with a different XNaci°> were performed, X I P is the mole fraction of 1-propanol and XNaci° is the initial mole fraction of NaCl. x N a C I ° = 0 0 X N a C I = 0.01090' • ' X N a C I " ' = 0.01962 x N a C l ° = 0.03507 Hip X | P Hip X i p Hip X i p H 1 P kJ/mol kJ/mol kJ/mol kJ/mol 0.0016 -10.09 0.0015 -9.06 0.0016 -8.41 0.0049 -6.98 0.0048 -9.90 0.0046 -8.88 0.0049 -8.26 0.0081 -6.79 0.0082 -9.59 0.0076 -8.52 0.0082 -7.82 0.0127 -6.470 0.0112 -9.22 0.0106 -8.26 0.0114 -7.55 0.0144 -6.06 0.0144 -8.87 0.0136 -7.95 0.0147 -7.22 0.0176 -5.67 0.0157 -8.35 0.0166 -7.67 0.0179 -6.73 0.0207 -5.18 0.0175 -8.43 0.0196 -7.13 0.0211 -6.31 0.0269 . -4.50 0.0206 -7.76 0.0227 -6.68 0.0242 -5.81 0.0300 -3.43 0.0237 -7.65 0.0259 -6.08 0.0274 -5.26 0.0331 -2.74 0.0276 -7.04 0.0291 -5.61 0.0305 -4.71 0.0361 -2.10 0.0311 -6.34 0.0322 -4.90 0.0336 -4.09 0.0391 -1.50 0.0347 -5.88 0.0354 -4.33 0.0367 -3.44 0.0451 -0.58 0.0382 -5.11 0.0385 -3.67 0.0397 -2.79 0.0483 -0.31 0.0417 -4.54 0.0416 -3.16 0.0431 -2.13 0.0518 -0.12 0.0451 -3.78 0.0448 -2.47 0.0466 -1.45 0.0552 0.44 0.0486 -3.17 0.0482 -1.82 0.0502 -0.95 0.0520 -2.42 0.0516 -1.29 0.0537 -0.52 0.0554 -1.93 0.0549 -0.86 0.0572 -0.23 0.0587 -1.33 0.0582 -0.62 0.0607 -0.01 0.0618 -0.98 0.0615 -0.18 0.0642 0.19 0.0651 -0.61 0.0647 0.023 0.0676 0.27 0.0683 -0.32 0.0710 0.37 0.0716 -0.27 0.0744 0.42 0.0748 0.15 0.0777 0.45 0.0812 0.37. 0.0851 0.53 0.0844 0.48 0.0927 0.59 0.0884 0.52 0.1001 0.62 0.0923 0.72 0.1074 0.63 0.0963 0.72 0.1002 0.79 0.1041 0.75 73 Table A3: Excess partial molar enthalpies of 1-propanol, H I P E , in lP-FJ^O-urea systems at 25°C. 6 series were performed with a different x u r e a 0. X I P is the mole fraction of 1-propanol and x u r e a 0 is the initial mole fraction of urea. V "= Aurt..i 0 A III I'll 0 o.KM Aurca 0 0499 Ainf.i 1) 1020 y " _ 0.1501 x ° = Aurea * ( 0.2035 X) P HipK X I P Hip E Xip Hip b Xip HipE Xip Hip E Xip H I P K kj/mol kJ/mol kJ/mol kJ/mol kJ/mol kJ/mol 0.0016 -10.09 0.0017 -8.59 0.0017 -7.92 0.0019 -6.13 0.0019 -5.02 0.0021 -4.15 0.0048 -9.90 0.0050 -8.34 0.0051 -7.51 0.0056 -5.93 0.0058 -4.70 0.0063 -3.96 0.0082 -9.59 0.0083 -7.98 0.0085 -7.34 0.0094 -5.80 0:0097 -4.67 0.0104 -3.70 0.0112 -9.22 0.0116 -7.50 0.0119 -6.94 0.0131 -5.45 0.0135 -4.48 0.0145 -3.56 0.0144 -8.87 0.0148 -7.46 0.0153 -6.65 0.0168 -5.26 0.0173 -4.34 0.0185 -3.34 0.0157 -8.35 0.0181 -7.04 0.0186 -6.36 0.0204 -4.97 0.0248 -3.83 0.0266 -3.04 0.0175 -8.43 0.0213 -6.71 0.0219 -6.i6 0.0276 -4.36 0.0285 -3.52 0.0305 -2.71 0.0206 -7.76 0.0245 -6.32 0.0252 -5.95 0.0312 -4.05 0.0322 -3.38 0.0344 -2.60 0.0237 -7.65 0.0276 -5.92 0.0285 -5.39 0.0347 -3.60 0.0358 -2.95 0.0383 -2.27 0.0276 -7.04 0.0339 -5.21 0.0317 -5.00 0.0383 -3.24 0.0394 -2.63 0.0422 -2.18 0.0311 -6.34 0.0370 -4.72 0.0349 -4.48 0.0418 -2.95 0.0430 -2.36 0.0460 -1.95 0.0347 -5.88 0.0401 -4.14 0.0381 -4.07 0.0452 -2.68 0.0466 -2.31 0.0536 -1.57 0.0382 -5.11 0.0435 -3.70 0.0413 -3.71 0.0487 -2.29 0.0545 -1.57 0.0577 -1.43 0.0417 -4.54 0.0470 -3.20 0.0445 -3.20 0.0521 -1.99 0.0585 -1.46 0.0620 -1.08 0.0451 -3.78 0.0506 -2.61 0.0479 -2.98 0.0557 -1.68 0.0625 -1.26 0.0664 -0.76 0.0486 -3.17 0.0575 -1.78 0.0513 -2.33 0.0635 -1.14 0.0664 -1.06 0.0706 -0.71 0.0520 -2.42 0.0610 -1.63 0.0552 -1.95 0.0673 -0.94 0.0703 -0.953 0.0749 -0.51 0.0554 -1.93 0.0678 -0.72 0.0588 -1.74 0.0712 -0.63 0.0780 -0.39 0.0791 -0.47 0.0587 -1.33 0.0712 -0.44 0.0660 -0.98 0.0749 -0.45 0.0818 -0.30 0.0832 -0.30 0.0618 -0.98 0.0746 -0.20 0.0696 -0.64 0.0787 -0.33 0.0894 -0.03 0.0873 -0.21 0.0651 -0.61 0.0779 -0.02 0.0731 -0.41 0.0824 -0.14 0.0931 -0.06 0.0914 -0.08 0.0683 -0.32 0.0812 0.09 0.0766 -0.17 0.0861 -0.04 0.0968 0.07 0.0955 -0.02 0.0716 -0.27 0.0878 0.33 0.0800 -0.06 0.0897 0.01 0.1004 0.16 0.0995 0.04 0.0748 0.15 0.0910 0.36 0.0835 0.06 0.0970 0.11 0.1041 0.11 0.1035 0.09 0.0812 0.37 0.0945 0.43 0.0869 0.21 0.1005 0.16 0.1080 0.20 0.1074 0.14 0.0844 0.48 0.0982 0.44 0.0903 0.23 0.1041 0.25 0.1208 0.27 0.1113 0.15 0.0884 0.52 0.1020 0.53 0.0969 0.31 0.1079 0.27 0.1250 0.29 0.1152 0.16 0.0923 0.72 0.1057 0.55 0.1006 0.50 0.1119 0.32 0.1291 0.34 0.1190 0.24 0.0963 0.72 0.1129 0.67 0.1044 0.41 0.1160 0.29 0.1228 0.24 0.1002 0.79 0.1165 0.61 0.1083 0.49 0.1200 0.29 0.1266 0.25 0.1041 0.75 0.1201 0.64 0.1121 0.51 0.1317 0.40 0.1303 0.26 0.1237 0.65 0.1159 0.48 0.1272 0.68 0.1197 0.56 0.1307 0.67 0.1271 0.57 0.1341 0.67 0.1307 0.54 0.1344 0.57 74 Table A4: lP-H 20-urea Vapour Pressures at 25.00°C. Series 1 Series 2 Series 3 v Nil. Xme.i P (Torr) XIP ^urea P (Torr) HIP •Vire.l P (Torr) 0 0 23.756 0 > 0.0306 23.071 0 0.0441 22.784 0.0037 0 24.741 0.0023 0.0306 23.689 0.0027 0.0440 23.557 0.0079 0 25.795 0.0050 0.0305 24.338 0.0055 0.0440 24.196 0.0135 0 27.023 0.0082 0.0305 25.100 0.0088 0.0439 24.931 0.0206 0 28.484 0.0121 0.0304 25.956 0.0123 0.0438 25.720 0.0291 0 30.029 0.0172 0.0303 27.013 0.0162 0.0437 26.545 0.0396 0 31.608 0.0238 0.0301 28.278 0.0204 0.0436 27.376 0.0503 0 32.928 0.0321 0.0299 29.719 0.0250 0.0434 28.214 0.0619 0 33.891 0.0414 0.0296 31.071 0.0297 0.0433 29.046 0.0748 0 34.459 0.0515 0.0293 32.230 0.0352 0.0431 29.899 0.0872 0 34.754 0.0603 0.0391 32.935 0.0411 0.0428 30.731 0.1036 0 34.986 0.0693 0.0288 33.420 0.0477 0.0426 31.519 0.1261 0 35.151 0.0786 0.0286 33.733 0.0552 0.0424 32.246 0.1561 0 35.276 0.0886 0.0283 33.951 0.0669 0.0414 33.020 0.1911 0 35.364 0.1011 0.0279 34.136 0.0787 0.0409 33.478 0.2241 0 35.427 0.1166 0.0275 34.281 0.0914 0.0404 33.748 0.1362 0.0270 34.391 0.1048 0.0399 33.924 0.1566 0.0265 34.463 0.1210 0.0393 34.062 0.1397 0.0387 34.157 0.1669 0.0378 34.251 0.2071 0.0368 34.314 Series 4 Series 5 Series 6 Xip u^rea P (Torr) XIP Xurea P (Torr) Xip i - < x. rea P(Torr) i 0 0.0746 22.138 0 0.1007 21.571 0 0 1921 19.745 0.0037 0.0743 23.163 0.0027 0.1005 22.281 0.0038 0 1917 20.869 0.0082 0.0741 24.281 0.0060 0.1003 23.085 0.0091 0 1911 22.045 0.0131 0.0738 25.390 0.0099 0.1000 23.969 0.0152 0 1903 23.346 0.0199 0.0733 26.810 0.0149 0.0996 24.880 0.0220 0 1893 24.663 0.0276 0.0728 28.209 0.0199 0.0992 25.883 0.0307 0 1880 26.173 0.0381 0.0721 29.851 0.0256 0.0988 26.946 0.0411 0 1864 27.683 0.0499 0.0713 31.286 0.0325 0.0982 28.086 0.0529 0 1844 29.031 0.0642 0.0702 32.880 0.0428 0.0972 29.548 0.0657 0 1823 30.107 0.0818 0.0390 33.350 0.0542 0.0962 30.781 0.0808 0 1798 30.915 0.1025 0.0675 33.641 0.0653 0.0952 31.619 0.0953 0 1774 31.339 0.1268 0.0658 33.783 0.0770 0.0941 32.161 0.1128 0 1746 31.620 0.0885 0.0930 32.522 0.1391 0 1707 31.832 0.0983 0.0922 32.686 ' 0.1154 0.0906 33.131 0.1439 0.0881 33.408 Table A 5 : Partial pressures, Pj (i = IP, W, Urea) (Torr) x u / Xip 0.0 0.005 0.01 0.015 0.02 0.025 0.03 0.035 0.04 0.045 0.05 0.055 P, 0.00 P w otu P.p 23.76 24.93 26.09 27.16 28.19 29.14 30.02 30.83 31.55 32.19 32.74 33:21 23.76 23.64 23.52 23.41 23.30 23.21 23.11 23.02 22.94 22.87 22.79 22.73 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 1.292 2.565 3.751 4.879 5.935 6.913 7.806 8.611 8.325 9.948 10.48 P, 0.01 P w a u P.p 23.52 24.73 25.94 27.05 28.06 28.99 29.83 33.59 31.26 31.85 32.38 32.84 -23.52 23.43 23.26 23.15 23.04 22.94 22.86 22.76 22.68 22.61 22.54 22.47 0.055 0.055 0.065 0.065 0.065 0.065 0.065 0.065 0.065 0.065 0.065 0.065 0.0 1.299 2.670 3.892 5.014 6.040 6.973 7.819 8.577 9.253 9.849 10.37 P, 0.03 P w a u P I P 23.10 24.32 25.52 26.63 27.64 28.56 29.40 30.16 30.83 31.43 31.96 32.42 23.10 23.01 22.86 22.75 22.64 22.54 22.45 22.36 22.27 22.20 22.12 22.06 0.147 0.147 0.152 0.152 0.152 0.152 0.152 0.152 0.152 0.152 0.152 0.153 0.0 1.312 2.662 3.879 4.997 6.022 6.954 7.799 8.559 9.236 9.836 10.36 P, 0.05 P w a u P . P 22.67 23.92 25.11 26.21 27.22 28.14 28.98 29.73 30.41 31.01 31.54 32.00 22.67 22.59 22.44 22.33 22.23 22.13 22.03 21.94 21.86 21.78 21.70 21.63 0.231 0.231 0.236 0.236 0.236 0.236 0.236 0.236 0.236 0.236 0.237 0.237 0.0 1.326 2.665 3.876 4.991 6.012 6.943 7.788 8.549 9.231 9.835 10.36 P. 0.07 P w a u P.p 22.25 23.51 24.70 25.79 26.80 27.71 28.55 29.30 29.97 30.58 31.11 31.58 22.25 22.17 22.02 21.91 21.81 21.71 22.61 21.52 21.43 21.36 21.28 21.21 0.312 0.312 0.316 0.316 0.316 0.316 0.316 0.316 0.316 0.316 0.317 0.317 0.0 1.340 2.671 3.877 4.988 6.007 6.937 7.783 8.545 9.226 9.832 10.37 P, 0.09 P w a u P.p 21.83 23.10 24.28 25.37 26.37 27.29 28.12 28.87 29.55 30.16 30.69 31.36 21.83 21.75 21.60 21.49 21.39 21.29 21.19 21.10 21.02 20.93 20.85 20.78 0.389 0.389 0.392 0.393 0.393 0.393 0.393 0.393 0.393 0.394 0.395 0.396 0.0 1.353 2.679 3.880 4.987 6.002 6.931 7.774 8.537 9.223 9.835 10.37 P. 0.11 P W a u P.p 21.41 22.70 23.87 24.95 25.95 26.86 27.69 28.45 29.12 29.73' 30.26 30.73 21.41 21.33 21.18 21.07 20.97 20.87 20.77 20.67 20.59 20.50 20.43 20.36 0.464 0.464 0.468 0.468 0.468 0.468 0.468 0.469 0.470 0.470 0.470 0.471 0.0 1.367 2.687 3.882 4.984 5.998 6.926 7.773 8.538 9.224 9.836 10.38 P, 0.13 P w a u P.p 20.99 22.29 23.46 24.54 25.53 26.44 27.27 28.02 28.70 29.30 29.84 30.31 20.99 20.91 20.76 20.65 20.54 20.44 20.34 20.26 20.17 20.09 20.01 19.93 0.537 0.537 0.541 0.542 0.524 0.542 0.542 0.542 0.542 0.542 0.543 0.544 0.0 1.381 2.696 3.885 4.986 4.997 6.923 7.765 8.527 9.216 9.833 10.38 P. 0.15 P W a u PIP 20.56 21.88 23.04 24.12 25.11 26.01 26.84 27.59 28.27 28.88 29.42 29.89 20.56 20.49 20.34 20.23 20.12 20.02 19.93 19.84 19.74 19.66 19.57 19.50 0.608 0.608 0.613 0.613 0.613 0.613 0.613 0.613 0.614 0.615 0.616 0.618 0.0 1.394 2.705 3.890 4.982 5.989 6.911 7.757 8.527 9.221 9.842 10.40 P. 0.17 P W a u P.p 20.14 21.48 22.63 23.70 24.68 25.59 26.41 27.16 27.84 28.45 28.99 29.47 20.14 20.07 19.92 19.81 19.70 19.60 19.49 19.39 19.31 19.22 19.14 19.06 0.678 0.678 0.683 0.683 0.683 0.684 0.684 0.686 0.684 0.684 0.688 0.690 0.0 1.408 2.713 3.890 4.982 5.991 6.920 7.767 8.534 9.229 9.853 10.41 P, 0.19 P W a u P.p 19.72 21.07 22.22 23.28 24.26 25.16 26.99 26.74 27.42 28.02 28.57 29.05 19.72 19.65 19.49 19.37 19.26 19.16 19.06 18.97 18.88 18.79 18.71 18.63 0.747 0.747 0.753 0.754 0.755 0.756 0.756 0.756 0.757 0.757 0.758 0.759 0.0 1.422 2.725 3.905 4.996 6.003 6.923 7.765 8.535 9.230 9.856 10.42 Table A 5 continues next page. 76 Table A5 (continues) x u / X i p 0.06 0.065 0.07 0.075 0.08 0.085 0.09 P, 33.60 34.16 34.46 34.60 34.68 34.85 34.79 0.00 Pw 22.68 22.58 22.51 22.47 22.47 22.47 22.47 a„ 0.0 0.0 0.0 0.0 0.0 0.0 0.0 P.p 10.93 11.57 11.95 12.13 12.12 12.38 12.38 P. 33.23 33.83 34.23 34.46 34.57 34.60 34.61 0.01 Pw 22.41 22.31 22.24 22.19 22.15 22.14 22:15 a u ,0.065 0.065 0.065 0.066 0.068 0.068 0.068 P.p 10.82 11.52 11.99 12.27 12.42 12.46 12.46 P, 32.81 33.42 33.83 34.07 34.19 34.23 34.25 0.03 Pw 22.00 21.89 21.81 21.74 21.73 21.69 21.65 a u 0.153 0.154 0.155 0.158 0.156 0.161 0.165 P.p 10.82 11.53 12.02 12.33 12.47 12.55 12.60 P, 32.40 33.02 33.43 33.69 33.82 33.87 33.89 0.05 Pw 21.57 21.46 21.37 21.32 21.24 21.22 21.22 a u 0.238 0.239 0.241 0.241 0.249 0.249 0.248 P IP 10.83 11.55 12.03 12.36 12.57 12.65 12.67 Pt 31.98 32.61 33.03 33.30 33.44 33.50 33.54 0.07 Pw 21.15 21.03 20.94 20.84 20.83 20.77 20.68 a u 0.318 0.320 0.322 0.329 0.325 0.332 0.344 P.p 10.83 11.58 12.10 12:46 12.61 12.74 12.86 P. 31.56 32.20 32.64 32.91 33.07 33.14 0.09 Pw 20.71 20.59 20.49 20.45 20.33 20.29 a u 0.397 0.399 0.402 0.400 0.412 0.415 P.p 10.84 11.60 12.14 12.46 12.74 12.85 Pt 31.14 31.79 32.24 32.53 32.69 0.11 Pw 20.29 20.17 20.06 19.92 19.90 a u 0.472 0.474 0.477 0.490 0.487 P.p 10.85 11.62 12.17 12.61 12.79 P, 30.72 31.38 31.84 32.14 0.13 Pw 19.86 19.73 19.61 19.61 a u 0.545 0.548 0.553 0.544 P.p 10.86 11.65 12.22 12.52 P, 30.31 30.97 31.44 31.751 0.15 Pw 19.42 19.29 19.18 18.91 a„ 0.619 0.622 0.626 0.655 P.p 10.88 11.68 12.26 12.84 P, 29.89 30.56 0.17 Pw 18.99 18.85 a u 0.691 0.694 P.p 10.90 11.71 P, 29.47 0.19 Pw 18.56 a u 0.761 P.p 10.91 77 Table A6! Excess partial molar entropies o f 1-propanol with the temperature factor, TSIPE in l P - E k O - U r e a systems at 25°C. x u r e a is the mole fraction o f urea and XIP is the mole fraction o f 1P A'uica -o X u r e a =0.05 i=0M0 ^ A u r e a -0.15 X i p TS.p* X i p T S ^ X , p TS I P K X i p TS.p* (kJ/mol) (kJ/mol) (kJ/mol) (kJ/mol) 0.005 -16.04 0.005 -13.82 0.005 -12.34 0.005 -11.29 0.010 -15.53 0.010 -13.49 0.010 -12.00 0.010 -11.02 0.015 -14.96 0.015 -12.96 0.015 -11.56 0.015 -10.69 0.020 -14.20 0.020 -12.36 0.020 -11.09 0.020 -10.32 0.025 -13.43 0.025 -11.74 0.025 -10.58 0.025 -9.86 0.030 -12.67 0.030 -11.06 0.030 -10.06 0.030 -9.37 0.035 -11.6 0.035 -10.42 0.035 -9.46 0.035 -8.85 0.040 -10.6 0.040 -9.55 0.040 -8.83 0.040 -8.35 0.045 -9.45 0.045 -8.77 0.045 -8.24 0.045 -7.87 0.050 -8.31 0.050 -8.03 0.050 -7.69 0.050 -7.44 0.055 -7.28 0.055 -7.38 0.055 -7.18 0.055 -7.05 0.060 -6.50 0.060 -6.77 0.060 -6.72 0.060 -6.67 0.065 -5.35 0.065 -5.76 0.065 -5.86 0.065 -5.99 0.070 -4.55 0.070 -4.96 0.070 -5.18 0.070 -5.37 0.075 -3.95 0.075 -4.40 0.075 -4.67 0.075 -4.88 0.080 -3.54 0.080 -4.01 0.080 -4.26 0.085 -3.24 0.085 -3.71 78 Table A 7 : The excess partial molar enthalpies of 1-propanol, HipE, in lP-FbO-Fruc systems at 25°C. X f r U c° is the initial mole fraction of fructose and X j p is the mole fraction of 1-propanol. X , ° = - 0 x ° = ^urea .0.01494.- *nirea 0.03151 x 0 = :< Aurea 0.05087 „'x 0 = ' ^ 'Virca 0.05781 x 0 = 0.06330 X i p H|P kJ/mol X i p H 1 P kJ/mol H 1 P E kJ/mol X,p H 1 P E kJ/mol X,p Hip E kJ/mol X1P HIP kJ/mol 0.0016 -10.09 0.0016 -8.40 0.0018 -6.53 0.0021 -5.48 0.0021 -4.22 0.0022 -3.89 0.0080 -9.61 0.0048 -8.15 0.0052 -6.28 0.0105 -4.55 0.0064 -3.99 0.0065 -3.66 0.0112 -9.22 0.0080 -7.86 0.0087 -6.13 0.0146 -4.07 0.0106 -3.45 0.0108 -3.24 0.0144 -8.87 0.0112 -7.45 0.0122 -5.78 0.0187 -3.47 0.0148 -3.20 0.0150 -2.68 0.0157 -8.35 0.0143 -7.09 0.0190 -5.08 0.0228 -3.02 0.0189 -2.81 0.0192 -2.18 0.0175 -8.35 0.0175 -6.75 0.0223 -4.60 0.0268 -2.43 0.0231 -2.34 0.0234 -2.60 0.0206 -7.70 0.0237 -5.93 0.0257 -4.10 0.0308 -1.92 0.0271 -1.82 0.0275 -1.45 0.0276 -7.04 0.0267 -5.64 0.0323 -3.11 0.0347 -1.54 0.0352 -0.89 0.0316 -1.12 0.0311 -6.34 0.0405 -3.32 0.0356 -2.76 0.0387 -0.97 0.0392 -0.53 0.0357 -0.62 0.0347 -5.88 0.0474 -2.15 0.0389 -2.16 0.0426 -0.54 0.0431 -0.21 0.0397 -0.25 0.0382 -5.12 0.0508 -1.97 0.0421 -1.71 0.0464 -0.11 0.0471 0.14 0.0437 0.15 0.0417 -4.54 0.0541 -1.13 0.0485 -1.00 0.0503 0.20 0.0509 0.43 0.0477 0.35 0.0451 -3.78 0.0575 -0.81 0.0520 -0.53 0.0541 0.40 0.0548 0.60 0.0516 0.67 0.0486 -3.17 0.0608 -0.38 0.0559 -0.15 0.0579 0.56 0.0586 0.66 0.0555 0.69 0.0520 -2.42 0.0641 -0.11 0.0597 0.13 0.0619 0.72 0.0624 0.73 0.0640 0.93 0.0554 -1.93 0.0674 0.09 0.0635 0.34 0.0661 0.83 0.0662 0.74 0.0683 1.16 0.0587 -1.33 0.0706 0.26 0.0710 0.58 0.0727 0.89 0.0768 0.98 0.0618 -0.98 , 0.0741 0.44 0.0747 0.74 0.0744 0.91 0.0810 0.96 0.0651 -0.61 0.0778 0.53 0.0783 0.73 0.0866 0.91 0.0852 0.95 0.0683 -0.33 0.0815 0.60 0.0820 0.78 0.0907 0.91 0.0893 0.89 0.0716 -0.28 0.0851 0.65 0.0856 0.81 0.0946 0.89 0.0748 0.15 0.0887 0.70 0.0927 0.82 0.0812 0.37 0.0923 0.75 0.1032 0.86 0.0844 0.48 0.0959 0.79 0.0884 0.53 0.0994 0.83 0.0923 0.72 0.1029 0.80 0.0963 0.73 0.1064 0.82 0.1002 0.79 0.1099 0.82 79 Table A 8 : Partial pressures (torr) of IP and water in IP-H2O-F1HC system at 25°C. Xfnjc / X i p 0.0 0.005 0.010 0.015 0.020 0.025 0.030 0.040 0.050 0.060 0.070 0.075 p. 0.0 P w P.p 23.76 24.95 26.14 27.29 28.36 29.30 30.11 31.70 32.96 33.76 34.26 34.44 23.76 23.05 22.57 22.25 22.00 21.78 21.58 31.44 21.39 21.30 21.25 21.25 0.0 1.902 3.566 5.044 6.358 7.517 8.529 10.26 11.57 12.46 13.01 13.19 Pt 0.005 P w P.p 23.65 24.95 26.20 27.40 28.45 29.43 30.24 31.86 33.09 33.94 34.42 34.60 23.65 23.01 22.58 22.32 22.11 21.98 21.84 21.86 21.91 21.68 21.48 21.39 0.0 1.941 3.616 5.077 6.341 7.455 8.402 10.00 11.18 12.26 12.94 13.21 Pt 0.01 P w p.p 23.58 24.95 26.25 27.46 28.55 29.55 30.40 32.02 33.22 34.08 34.53 34.71 23.58 23.00 22.63 22.40 22.24 22.15 22.08 22.15 22.22 21.90 21.58 21.40 0.0 1.949 3.621 5.063 6.307 7.397 8.325 9.871 11.00 12.18 12.95 13.31 p, 0.015 P w P.p 23.50 24.95 26.29 27.55 28.66 29.69 30.56 32.19 33.35 34.18 34.63 34.78 23.50 23.01 22.68 22.50 22.37 22.31 22.24 22.31 22.31 21.95 21.55 21.30 0.0 1.942 3.609 5.051 6.290 7.384 8.316 9.878 11.04 12.23 13.08 13.48 P. 0.02 P w P.p 23.41 24.95 26.33 27.61 28.77 29.82 30.74 32.35 33.48 34.27 34.71 34.83 23.41 23.01 22.72 22.55 22.45 22.38 22.34 22.33 22.22 21.86 21.40 21.10 0.0 1.937 3.607 5.058 6.317 7.435 8.403 10.02 11.26 12.41 13.31 13.73 P. 0.03 P w P.p 23.26 24.95 26.39 27.74 28.97 30.10 31.08 32.65-33.74 34.40 34.80 34.89 23.26 22.97 22.68 22.50 22.38 22.28 22.18 21.93 21.59 21.26 20.74 20.43 0.0 1.981 3.708 5.240 6.593 7.820 8.901 10.72 12.15 13.14 14.06 14.46 Pt 0.04 P w P.p 23.10 24.93 26.42 27.84 29.15 30.33 31.38 32.93 33.96 34.52 34.85 34.92 23.10 22.78 22.39 22.11 21.91 21.72 21.54 21.09 20.58 20.19 19.72 19.47 0.0 2.151 4.033 5.725 7.239 8.612 9.837 11.84 13.38 14.33 15.13 15.45 P, 0.05 P w P.p 22.96 24.88 26.43 27.94 29.31 30.55 31.63 33.16 34.14 34.62 34.88 34.93 22.96 22.41'21.82 21.41 21.08 20.79 20.52 19.94 19.42 18.81 18.44 18.32 0.0 2.470 4.611 6.530 8.235 9.764 11.11 13.22 14.72 15.81 16.44 16.61 P. 0.06 P w P.p 22.81 24.81 26.42 28.01 29.43 30.70 31.83 33.33 34.28 34.72 34.90 34.94 22.81 21.89 21.03 20.45 19.98 19.60 19.30 18.67 18.27 17.26 17.03 17.06 0.0 2.921 5.386 7.556 9.446 11.10 12.53 14.66 16.01 17.46 17.87 17.88 Pt 0.07 P w P.p 22.65 24.73 26.39 26.06 29.50 30.81 31.97 33.46 34.38 34.81 34.92 34.95 22.65 21.29 20.14 18.02 18.81 18.37 18.05 17.46 17.20 15.67 15.59 15.76 0.0 3.445 6.247 8.036 10.69 12.44 13.92 16.00 17.18 19.14 19.33 19.19 P. 0.08 P w P.p 22.46 24.64 26.33 38.03 29.80 30.89 32.06 33.56 34.44 34.86 34.94 34.97 22.46 20.68 19.28 18.40 17.90 17.29 16.95 16.42 16.28 14.10 14.17 14.48 0.0 3.957 7.052 9.634 11.90 13.60 15.11 17.14 18.16 20.76 20.77 20.49 P. 0.09 P w P.p 22.25 24.45 26.23 27.93 29.44 30.87 32.07 33.58 34.44 34.84 34.96 34.98 22.25 20.11 18.12 17.13 16.44 15.98 15.58 14.97 14.83 12.61 12.82 13.25 0.0 4.343 7.654 10.34 12.55 14.42 15.95 17.97 18.95 22.23 22.14 21.73 P. 0.10 P w p.p 21.98 24.15 25.99 27.73 29.28 30.72 31.91 33.50 34.38 34.73 21.98 19.67 18.12 17.13 16.44 15.98 15.63 15.10 14.83 11.23 0.0 4.482 7.870 10.60 12.84 14.74 16.28 18.40 19.55 23.50 p. 0.11 P w p.p 21.48 23.64 25.56 27.39 29.00 30.42 31.64 33.32 34.19 34.30 21.48 19.42 18.07 17.20 16.56 16.08 15.68 14.96 14.26 9.957 0.0 3.885 6.996-9.628 11.86 13.79 15.45 18.10 19.93 24.53 p. 0.12 P w P.p 20.62 22.90 24.96 26.95 28.61 30.02 31.27 33.01 33.88 34.11 20.62 19.53 18.74 18.21 17.68 17.14 16.61 15.34 13.78 8.804 0.0 3.373 6.223 8.737 10.93 12.88 14.66 17.67 20.10 25.31 80 Table A9: Excess chemical potential of IP, p i p E , and excess partial molar entropy of IP (including the temperature factor), T S I P E , in lP-FbO-Fruc systems at 2 5 ° C v 0 -Afruc 0 Xfhic 0.0149 xfruc 0.0315 0.0509 Y ° — Afruc 0.0579 Afruc ' 0.0633 M-IPE kJ/mol TS, P E kJ/mol HIPE kJ/mol TS 1 P E kJ/mol kJ/mol TS 1 P E kJ/mol HIPE kJ/mol TS, P E kJ/mol U-1PE kJ/mol TS 1 P E kJ/mol I^1PE kJ/mol TS 1 P E kJ/mol 0.000 7.52 -17.72 7.55 -16.10 7.61 -14.25 8.40 -14.13 8.63 -12.96 8.93 -12.91 0.005 7.33 -17.17 7.34 -15.44 7.50 -13.80 8.00 -13.15 8.28 -12.34 8.54 -12.26 0.010 7.17 -16.57 7.16 -14.76 7. 33 -13.35 7.83 -12.43 8.08 -11.81 8.32 -11.63 0.015 7.02 -15.84 6.99 -14.04 7.18 -12.75 7.68 -11.68 7.93 -11.16 8.15 -10.93 0.020 6.89 -15.03 6.83 -13.28 7.03 -11.99 7.55 -10.91 7.78 -10.45 7.98 -10.22 0.025 6.75 -14.12 6.68 -12.58 6.90 -11.14 7.41 -10.08 7.63 -9.70 7.82 -9.52 0.030 6.61 -13.17 6.53 -11.59 6.76 -10.23 7.28 -9.28 7.49 -8.98 7.66 -8.82 0.035 6.48 -12.21 6.38 -10.61 6.65 -9.40 7.19 -8.56 7.37 -8.32 7.49 -7.49 0.040 6.35 -11.16 6.26 -9.66 6.50 -8.55 7.00 -7.79 7.18 -7.65 7.32 -6.94 0.045 6.23 -10.05 6.12 -8.67 6.36 -7.71 6.85 -7.10 7.03 -7.03 7.14 -6.27 0.050 6.10 -8.94 5.99 -7.72 6.24 -6.95 6.71 -6.54 6.87 -6.52 6.75 -6.14 < 0.055 5.96 -7.90 5.88 -6.90 6.12 -6.33 6.57 -6.12 6.72 -6.13 6.81 -5.94 0.060 5.82 -7.05 5.78 -6.27 5.99 -5.85 6.44 -5.77 6.61 -5.93 6.74 -5.69 0.065 5.68 -6.30 5.67 -5.73 5.88 -5.48 6.30 -5.50 6.46 -5.73 6.58 -5.45 0.070 5.55 -5.74 5.57 -5.45 5.77 -5.21 6.15 -5.27 6.41 -5.26 0.075 5.42 -5.29 5.47 -5.05 5.67 -4.99 6.00 -5.08 6.23 81 Table A10: GC analysis: Ratio of peak area of IP to H 2 0 in the lP-H 20-Fruc system at 25°C. i f l l t H X | , I , C - ' 0.0258: : E*x^ £-^ jjy§ 0.0496 i&!;o:o9ii • V i n e 0.1221 X ) p 0.0039 1P/H20 0.0666 X i p 0.0045 1P/H20 0.0834 X i p 0.0050 1P/H20 0.1164 X i p 0.0048 1P/H20 0.1045 X i p 0.0066 1P/H20 0.4122 0.0083 0.1389 0.0091 0.1655 0.0093 0.1875 0.0120 0.4842 0.0191 0.5687 0.0123 0.1960 0.0136 0.2138 0.0143 0.2868 0.0181 0.7109 0.0321 1.0044 0.0203 0.2591 0.0225 0.2862 0.0243 0.4416 0.0287 0.7618 0.0504 1.4367 0.0319 0.4386 0.0356 0.4888 0.0339 0.5761 0.0356 0.8069 0.0511 0.5401 0.0517 0.5610 0.0650 0.8159 0.0572 1.2174 0.0685 0.6711 0.0681 0.6106 0.0821 0.8756 0.0880 0.6268 0.0819 0.7209 82 Derivation of Equation (3): H = PV + E dH = PdV + VdP + dE = TdS + VdP dH = (dH/dS)PdS + (5H/5P)sdP Maxwell Relationships Therefore T = (dH/dS)P V = (dH/dP)s G = H - TS dG = dH - TdS - SdT = VdP - SdT V = (dG/dP)T -S = (5G/5T)P = (G-H) / T H =-T (SG/5T)P + G H = - T 2 {3(G/T) / 5T}P H = { 5(G/T) / 5(1/T) }P Therefore, H is the first derivative of G. Hi E = dHE/arii =-T2{drGE/T)/STdni} Hence, H E is the 2nd derivative of G. dHE/ani = dW^m = -T2{a3(GE/T)/5T52n i} Define HiiE = N(aHiE/ani) = - T 2 N {tfiGp/T) I ST^ m} Thus, HjjE is the third derivative of G. 83 Converting (ni, nw) to (N, Xj) with n w , P, T constant: Hi E = HiE (ni, nw) = H , E (N, X i) dHjE = (aHiE/5ni)nw dni + (aHiE/dnw)ni dn w Since dn w = 0, (5HiE/5nw)ni dn w = 0 Therefore, dHiE = (aHE/5ni)nw (a) At the same time, dHjE = (aHiE/5nj)N dxj + (5HiE/5N)xi d N (b) Equating (a) and (b), (5HiE/ani)nw dns = (aHiE/dxi)Ndxi + (5HiE/aN)xi dN (aHiE/5ni)nw = (aHiE/axi)N (axi/ani)nw + (aH E / a N ) x i (aN/ani)nv Since (aN/anj)nw = 1, (axi/ani)nw = { a[ni/(ni+nw)] / anj}nw = n, + nw - n, = 1 - x, (m + n w) 2 " N •" Therefore, (aHiE/ani)nw=(i-xi / N) (aHiE/axi)N + (aHiE/aN)xi Since H E is an intensive quantity, (aHiE/aN)xi=0 Thus, (aHiE/ani)nw = O-xO (5HE/axi)nw N Hence, Define H n E = N (aHjE/anOnw = (1-Xj) (aHiE/axOnw (equation 3) Numerical evaluation of partial pressures From the total vapour pressure data, the partial pressures can be calculated as detailed in J. Phys. Chem. B, 1998,102, 5182. However, for the lP-H 20-Urea systems, the partial pressure of urea is negligibly small compared to those of 1 -propanol and water. Therefore the Gibbs-Duhem relation becomes: xu81nau + xip51nPip + xw81nPw = 0 P = P 1 P + Pw, where au(=xuYu) is the activity of urea. Except this, the data analysis method is the same as in the above reference paper. For urea, pu = Pu° + RT lnau = pu° + RT lnxu + RT lnyu where pu° is the chemical potential of pure solid urea. Therefore, at saturation, Pu = Pu°-Hence, au s a t = xu s a tyu s a t = 1- The "sat" superscript denotes saturation. The results are in Table A5. Thus the excess chemical potentials can be calculated: UjE = RT In {Pi/(Pi°Xi)} where i = IP and water. uuE = RTlnyu 85 For the lP-H20-Fructose system, the partial pressures were evaluated differently. GC analysis were performed twice. The data was accurate and highly reproducible. Since the ratio of P I P : Pw was obtained through peak areas determination, assuming the partial pressure of fructose is also negligibly small compared to IP and water, P i p / P w = a P = PlP + P W The partial pressures of IP and those of water were evaluated easily. The result are in Table A8. Excess chemical potentials of IP were then calculated as above in the IP-H2O-urea systems. 86 

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