Water-ethanol mixture, mole fraction

The ionic conductivity of [C2CiIm][Bp4] was higher than of Li[C2QIm][BF4]. The high influence of water and ethanol on ionic conductivity was observed for Tallyl-3-methylimidazolium salt. The conductivity increases with an increase of temperature and the addition of water and ethanol. For example, the conductivity was changed from 3.82 mScm for pure IL at 293.15 K to 8.17 mScm for the mixture (IL + 0.3 water + 0.1 ethanol in mole fraction) [209]. 

The dependence of the dissociation rate on the mole fraction of the organic solvent is depicted in Figure 5. As seen in Figure 5, the rate of dissociation decreases exponentially with the mole fraction of the ethanol. This decrease in rate of dissociation cannot be attributed to the effect of the solvent on the dielectric constant of the solution. At Xethanoi = 0.2, the dielectric constant of the mixture is 66.2 (vs. 77.5 of water), but the rate of dissociation is slowed by an order of magnitude. The proton conductivity of the water—ethanol mixture decreases with the mole fraction of the solvent, but this decrease is not steep enough to account for the measured effect on the rate of dissociation (see Figure 5). This reasoning... 

The solvolysis of tran5-[Co(4-Etpy)4Cl2] in water-isopropanal at various temperatures has been studied in detail. The activation energy varied nonlinearly with the mole fraction of the cosolvent. The plot of log k vs. the reciprocal of the dielectric constant was also nonlinear. The influence of the solvent structure on the complex ion in the transition state dominates over that in the initial state. A similar study has also been carried out using water-ethanol mixtures. ... 

In Fig. 3.7, we show the entropy and the enthalpy of solvation of methane as a function of the mole fraction of ethanol. Note that the anomalously large and negative entropy and enthalpy of solvation are confined to the water-rich region of the water-ethanol mixtures. 

Figure 4.30 shows the variation of 8S a ) and 8H ai) as a function of the mole fraction of ethanol in water-ethanol mixtures. The most prominent feature of these two curves is the steep decrease in 8S and 8H in the water-rich region... 

Figure 6 26. Experimental data for water mole fraction in the permeate and the permeation rate. (Temperature, 35 C membrane, polyethylene grafted with polyacrylic acid system, water/ethanol mixtures.) (Reproduced from [240] with permission.)...

Fig. 1.1 Water + ethanol mixtures at 298.15 K. Comparison of excess molar heat capacities C- at atmospheric pressure. Dashed curve, calculated from heats of mixing at different temperatures (870TT1) points, direct experimental measurements o (80BEN1), (81GR01) xi is the ethanol mole fraction...

Figure 6. (A) E (NR) as a function of mole fraction of water in water-ethanol mixture, (B) Ej.(NR) as a function of mole fraction of ethanol in acetone- ethanol mixture.

The aquation kinetics of the chloropentaamminecobalt(III) ion in water-ethanol mixtures has been studied. The rate constants correlate well with the Grunwald-Winstein Y parameter and with the dielectric constant of the medium. The data supports a D mechanism for the reaction. The loss of chloride from the complexes cw-[Co(en)2(NH2CH2CH20H)Cl] and cw-[Co(en)2(NH2(CH2)3 0H)Cl] has been studied in aqueous ethyleneglycol at 40-65 °C in acidic media and at 20-35 °C in basic media.The rate constants decreased linearly with the increasing mole fraction of the cosolvent. The loss of chloride resulted in the formation of the chelated amino-alcohols as the main product. The observed solvent isotope effect (A h2oAd2o) = 112 at 50 °C, [HCIO4] =0.01 moldm for chloride release is lower than the value reported for the aquation of the cw-[Co(en)2(alkylamine)Cl] complexes (1.38-1.44). This result may indicate the lack of direct solvent intervention in the act of substitution at the cobalt(III) center, as expected for a true intramolecular reaction. 

Addition of co-solvents can also change the co-miscibility characteristics of ionic liquids. As an example, the hydrophobic [BMIM][PFg] salt can be completely dissolved in an aqueous ethanol mixture containing between 0.5 and 0.9 mole fraction of ethanol, whereas the ionic liquid itself is only partially miscible with pure water or pure ethanol [13]. The mixing of different salts can also result in systems with modified properties (e.g., conductivity, melting point). 

Consider a second example involving the separation of a mixture of ethanol and water that forms an azeotrope at around a mole fraction of ethanol of 0.88. It is proposed to use ethylene glycol as entrainer. An overall mass balance for the separation is shown in Figure 12.22a. As with the... 

Fio. 54. Sttbility constants for ion>pair formation with decyltrimethylammonium car boxylates at 25 C in various mole fraction ethanol-water mixtures. The carboxylic acids used are (O) decanoate, ( ) nonanoate, ( ) heptanoate, and (A) octanoate. The stability constants were determined conductimetrically. The data are taken from Oakenful and Fen wick (222). 

Figure 3.3-7 Ethanol/water/[BMIM][PF5] ternary phase diagram (a, left) and solute distribution in EtOH/water/IL mixtures (b, right) for [BMIMJiPEe] (O), [HMIMJiPFe] ( ), and [OMIMJiPFe] (V) as a function of initial mole fraction of ethanol in the aqueous phase, measured at 25 °C. From references [47, 48].

Into a water-cooled mixture of 1025 gm (25.0 moles) of acetonitrile, 1150 gm (25.0 moles) of absolute ethanol and 900 ml of chloroform is slowly bubbled 913 gm (25.0 moles) of anhydrous hydrogen chloride. The temperature rises to 35°-40°C and then is recooled to 20°C. After 48 hr at 20°-25°C, 5 liters of absolute ethanol is added and then the mixture is left for 2 days. The precipitated ammonium chloride is filtered, washed with ethanol, and the filtrate and washings added to 20 liters of 5 % sodium hydroxide solution. The product is extracted with chloroform, concentrated and fractionated under reduced pressure to afford 2430-3250 gm (60-80%), b.p. 70°-80°C (60 mm). 

Salt concentrations were obtained by the addition of known weights of dried salt to the solvent mixture. The compositions of the condensed equilibrium vapor samples and the previously prepared ethanol-water charges to the still were determined as previously outlined. The mole fractions of the salt, ethanol, and water charged to the Othmer still were thus accurately determined by mass balance calculations. 

The data in Tables I-XVI (see Appendix for all tables) show the isobaric vapor-liquid equilibrium results at the boiling point for potassium, ammonium, tetramethylammonium, tetraethylammonium, tetra-n-propylammonium, and tetra-n-butylammonium bromides in various ethanol-water mixtures at fixed liquid composition ratios. The temperature, t, is the boiling temperature for all solutions in these tables. In all cases, the ethanol-water composition was held constant between 0.20 and 0.35 mole fraction ethanol since it is in this range that the most dramatic salt effects on vapor-liquid equilibrium in this particular system should be observed. That is, previous data (12-15,38) have demonstrated that a maximum displacement of the vapor-liquid equilibrium curve by salts frequently occurs in this region. In the results presented here, it should be noted that Equation 1 has been modified to... 

In Table XVIII, there are several trends that can be noted in k if one proceeds through the R4NBr series. First of all, k tends to decrease as the size of the TAA cation increases and, in fact, tetra-n-butylammonium bromide shows a large salting-in effect. This trend is emphatically demonstrated by Figure 13, which shows the smoothed salt effects of the various salts studied in the ethanol-water system at x = 0.206. Secondly, it appears that there is a larger salting-out effect as the mole fraction of ethanol increases in the binary solvent mixture. 

Figure 13.23. Examples of vapor-liquid equilibria in presence of solvents, (a) Mixture of-octane and toluene in the presence of phenol, (b) Mixtures of chloroform and acetone in the presence of methylisobutylketone. The mole fraction of solvent is indicated, (c) Mixture of ethanol and water (a) without additive (b) with 10gCaCl2 in 100 mL of mix. (d) Mixture of acetone and methanol (a) in 2.3Af CaCl2 ip) salt-free, (e) Effect of solvent concentration on the activity coefficients and relative volatility of an equimolal mixture of acetone and water (Carlson and Stewart, in Weissbergers Technique of Organic Chemistry IV, Distillation, 1965). (f) Relative volatilities in the presence of acetonitrile. Compositions of hydrocarbons in liquid phase on solvent-free basis (1) 0.76 isopentane + 0.24 isoprene (2) 0.24 iC5 + 0.76 IP (3) 0.5 iC5 + 0.5 2-methylbutene-2 (4) 0.25-0.76 2MB2 + 0.75-0.24 IP [Ogorodnikov et al., Zh. Prikl. Kh. 34, 1096-1102 (1961)].

Figure 6.1. Concentration dependence of the Fick diffusivity for the binary mixtures (a) water (A)—acetone (B) at 45°C, (b) water (A)—ethanol (B) at 40°C in terms of water mole fraction (Tyn and Calus, 1975).

Fig. 11. Dependence of the exchange current of the Zn(II)/Zn(Hg) system on the mole fraction (jCj) of organic solvents in their mixtures with water. Curve 1, water-diox-ane curve 2, water-ethanol curve 3, water-acetonitrile curve 4, water-tetrahydrofuran curve 5, water-acetone.

Figure 1. Henry s constants of gases in binary solvent mixtures at 760 mmHg partial pressure ( , experimental A, calculated with the new eq 25 B, calculated with Krichevsky eq 1 in Figure la, C represents eq A2-3) (a) oxygen (2) in 1-propanol (1)—water (3) at 40 °C, (b) nitrogen (2) in 2-propanol (l)-water (3) at 40 °C, (c) oxygen (2) in ethanol (l)-water (3) at 40 °C, and (d) carbon dioxide (2) in 1-propanol (l)-water (3) at 40 °C. Xaic is the mole fraction of the alcohol in the gas-free mixture of solvents.

There are several experimental determinations of the solubility of naphthalene in ethanol-water mixtures at room temperature (Bennett and Canady, 1984 Morris, 1988 Dickhut et ah, 1989 LePree et ah, 1994). These data deviate appreciably from each other (Fig. 4). The analysis of the above data with Eq. (11) (Table 2) indicated that those regarding the solubility of naphthalene in ethanol-water mixtures at room temperature, obtained by various authors, were thermodynamically consistent in the dilute region however, the data of LePree et al. (1994), and Morris (1988) are thermodynamically inconsistent at high mole fractions of ethanol. Only the data for ethanol mole fractions... 

Fig. 4. The solubility of naphthalene (X2) in ethanol-water mixtures atroomtemperature (o)BennettandCanady, 1984 ( ) Morris, 1988 ( ) Dickhut et al., 1989 (x) LePree et al., 1994). xetoh is the mole fraction of ethanol in the solute-free mixed solvent.

Fig. 6. D values (A) calculated with Eq. (11) for the solubility of sulphamethoxypyridazine in ethanol-water mixture at room temperature. XEtoH is the mole fraction of ethanol in a solute-free ethanol—water mixture.

One should emphasize that the experimental data of different research groups exhibit large deviations from each other. This is illustrated in Figure 1 for the solubility of naphthalene in the ethanol + water mixture. Figure 1 shows that there are large differences between the experimental solubilities of naphthalene (2,8 and 9) in the ethanol + water mixture, which are particularly large for mole fractions of ethanol between 0.05 and 0.25. Therefore, the inaccuracy of the experimental data should be taken into account, since... 

FIGURE 2. Solubilities of naphthalene (S is the mole fraction of naphthalene) in the mixtures a) methanol + water and b) ethanol + water. The experimental data (0) were taken from Ref. (2). The solid lines represent the solubilities of naphthalene predicted using equation M4. The Wilson constants were taken from Gmeling s vapor-liquid equilibrium compilation (2 ). Thus, the only solubilities in pure water and cosolvents were used for prediction. 

---