Aqueous solutions Solubility equilibria Solution

Colourless crystals m.p. I25°C, soluble in water and alcohol. In aqueous solution forms equilibrium with its lactones. Gluconic acid is made by the oxidation of glucose by halogens, by electrolysis, by various moulds or by bacteria of the Acetobacter groups. 

The thermodynamic solubility of a drug is the concentration of the compound that is dissolved in aqueous solution in equilibrium with the undissolved amount, when measured at 25°C after an appropriate time period. Aqueous solubility has long been recognized as a key molecular property in pharmaceutical science. Drug delivery, transport and distribution phenomena depend on solubility thus, it is of considerable value to possess information of the solubility value of a drug candidate, to be able to predict the solubility for unknown compounds and, finally, to be able to modify the structure of a compound in order to modulate its solubility value in an appropriate manner. 

In aqueous solution, the equilibrium between the cis- and franj-diaqua complexes lies almost completely toward the cis isomer881 (K 0.17, pH 3-4). The sparingly soluble potassium salt of the tram isomer may, however, be prepared by the slow evaporation of a saturated solution at room temperature,878,880 and the cis isomer by cooling a hot solution or by allowing potassium dichromate and oxalic acid to react in the presence of a minimal quantity of water.878,879,882 The tris complex was resolved by Werner in 1912,883 providing the first example of an anionic optically active coordination complex. 

It can precipitate as potassium hydrogen tartrate (KHT) or as calcium tartrate (CaT), the latter being practically insoluble in aqueous solutions. Their equilibrium solubility varies with temperature, pH, and alcohol content, while the presence of a few wine components, such as polysaccharides and mannoproteins, may hinder spontaneous nucleation even if the solution is supersaturated. From Figure 14 that shows the equilibrium tartaric acid-dissociated fractions versus pH and ethanol volumetric fraction (Berta, 1993 Usseglio-Tomasset and Bosia, 1978), it can be seen that in the typical pH range (3 4) of wines KHT is predominant. As temperature is reduced from 20 to 0°C, KHT solubility in water or in a 12% (v/v) hydro-alcoholic solution reduces from 5.11 to 2.45 kg/m3 or from 2.75 to 1.1 kg/m3, respectively (Berta, 1993). Each of these data also varies with pH and reaches a minimum at the pH value associated with the maximum concentration of the hydrogen tartrate anions. For the above-mentioned solutions, the solubility minimum shifts from pH 3.57 to pH 3.73 as the ethanol content increases from 0 to 12% (v/v) (Berta, 1993). 

In aqueous solution the equilibrium between the cis- and trans-[diaquabis-(oxalato)chromate(III)] ions lies almost completely on the side of the cis isomer pure crystals of the slightly soluble potassium salts of the trans isomer are formed by the slow evaporation of the solution at room temperature,2,4- 6 whereas the potassium salt of the cis isomer is obtained by cooling the hot solution or by allowing potassium dichromate and oxalic acid to react with only a minimal amount of water present,2,4-6 the method given here. 

Next we take a look at Fig. 14M, the simple potential/pH diagram representing the behayior of magnesium in aqueous solutions. For equilibrium between a solid and a soluble species, the concentration of the latter must be specified. In Fig. 14M(a) lines are shown for concentrations of Mg of 1 pM, 1 mM and 1 M. It is customary to simplify potential/pH diagrams by showing only the lines corresponding to a concentration of 1. tM of each soluble species. This convention is followed in all further potential/pH diagrams shown in this book. This is reasonable in view of the fact that a moderate rate of corrosion may... 

Although many of the substituted cyclopropenyl ions are marvels of stability, salts of the parent ion darken rather rapidly on heating and exposure to atmospheric moisture causes rapid decomposition. The hexachloroantimonate salt is stable for a long period at — 20° C and for a few hours at room temperature. As might be expected, the cyclopropenyl salts are soluble in polar solvents such as acetonitrile and dimethylformamide, but are insoluble in non-polar solvents. They, of course, react with protic solvents as described earlier. In aqueous solution the equilibrium of equation 24 is established. The pH required to... 

There has been some debate as to the identity of the aqueous arsenic species that are present in an aqueous solution in equilibrium with AS2S3 where that solution contains an excess of sulfide over that which would be present from dissolution of AS2S3 alone (see e.g. Krupp, 1990 Spycher and Reed, 1989). We had believed this controversy to have been settled recently by Helz et al. (1995) who utilized spectroscopy, molecular orbital calculations, and the solubility studies to arrive at the conclusion that the principal aqueous species in excess-sulfidic solutions are H2As3S6 (aq) and AsO(SH)2 (aq). 

This equation represents the solubility product of silver chloride. Solubility products are generally used to describe the solubility and equilibria of sparingly soluble salts in aqueous solutions. Solubility products of a number of substances are given in Table 1.3. It is important to remember that use of solubility product relations based on concentrations assumes that the solution is saturated, in equilibrium, and ideal (the activity coefficient is equal to one), and is therefore an approximation, except with very dilute solutions of one solute. 

We next consider the equilibrium or solubility of a vapor in aqueous solution. The equilibrium can be represented as ... 

Given the long gas-phase lifetime for formic acid (1 month), it is clear that heterogeneous processes (wet and dry deposition) also must be considered as sinks for this species, particularly in the lower troposphere. An important parameter in determining the rate of wet deposition (uptake of a gas into a hydrometer, followed by removal via precipitation) is the solubility of the gas in aqueous solution, an equilibrium process described by the Henry s law constant, Hx = [Jflaq/pz see also section I-B-4. Here, Hx is the Henry s law constant for species X, [X]aq is its equilibrium concentration in the aqueous phase, and px is its partial pressure of the gas. For the organic acids under consideration here, increased solubility can result from ionization in solution, e.g. ... 

From the standpoint of thermodynamics, the dissolving process is the estabHsh-ment of an equilibrium between the phase of the solute and its saturated aqueous solution. Aqueous solubility is almost exclusively dependent on the intermolecular forces that exist between the solute molecules and the water molecules. The solute-solute, solute-water, and water-water adhesive interactions determine the amount of compound dissolving in water. Additional solute-solute interactions are associated with the lattice energy in the crystalline state. 

These substances, having the formula CjHjNHCONH, and OC(NHCjH6)j respectively, are both formed when an aqueous solution of urea and aniline hydrochloride is heated. Their subsequent separation is based on the fact that diphenylurca is insoluble in boiling water, whereas monophenylurea is readily soluble. The formation of these compounds can be explained as follows. When urea is dissolved in water, a small proportion of it undergoes molecular rearrangement back to ammonium cyanate, an equilibrium thus being formed. 

In aqueous solution at 100° the change is reversible and equilibrium is reached when 95 per cent, of the ammonium cyanate has changed into urea. Urea is less soluble in water than is ammonium sulphate, hence if the solution is evaporated, urea commences to separate, the equilibrium is disturbed, more ammonium cyanate is converted into urea to maintain the equilibrium and evfflitually the change into urea becomes almost complete. The urea is isolated from the residue by extraction with boiling methyl or ethyl alcohol. The mechanism of the reaction which is generally accepted involves the dissociation of the ammonium cyanate into ammonia and cyanic acid, and the addition of ammonia to the latter ... 

CI2O is very soluble in water, a saturated solution at —9.4°C containing 143.6 g CI2O per 100 g H2O in fact the gas is the anhydride of hypochlorous acid, with which it is in equilibrium in aqueous solutions ... 

Thus, at equilibrium, aqueous solutions of mercury(I) salts will contain around 0.6% of mercury(II) and the rather finely balanced equilibrium is easily displaced. The presence of any reagent which reduces the activity (in effect the concentration) of Hg + more than that of Hg2 ", either by forming a less-soluble... 

The data given in Tables 1.9 and 1.10 have been based on the assumption that metal cations are the sole species formed, but at higher pH values oxides, hydrated oxides or hydroxides may be formed, and the relevant half reactions will be of the form shown in equations 2(a) and 2(b) (Table 1.7). In these circumstances the a + will be governed by the solubility product of the solid compound and the pH of the solution. At higher pH values the solid compound may become unstable with respect to metal anions (equations 3(a) and 3(b), Table 1.7), and metals like aluminium, zinc, tin and lead, which form amphoteric oxides, corrode in alkaline solutions. It is evident, therefore, that the equilibrium between a metal and an aqueous solution is far more complex than that illustrated in Tables 1.9 and 1.10. Nevertheless, as will be discussed subsequently, a similar thermodynamic approach is possible. 

The Number of Dipoles per Unit Volume. The Entropy Change Accompanying Proton Transfers. The Equilibrium between a Solid and Its Saturated Solution. Examples of Values of L and AF°. The Change of Solubility with Temperature. Uni-divalent and Other Solutes. Lithium Carbonate in Aqueous Solution. H2COj in Aqueous Solution. Comparison between HjCOj and Li2C03 in Aqueous Solution. Heats of Solution and the Conventional Free Energies and Entropies of Solution. 

A typical example is as follows. Benzoic acid, C6H5COOH, is a solid substance with only moderate solubility in water. The aqueous solutions conduct electric current and have the other properties of an acid listed in Section 11-2.1. We can describe this behavior with reaction (42) leading to the equilibrium relation (43) ... 

Equilibria that occur in aqueous solution are of particular interest, because water is the medium of life and a major influence on the geography of our planet. Many substances dissolve in water, and the solutes in an aqueous solution may participate in a number of different types of equilibria. Solubility itself is one important type of equilibrium, as we describe in Chapter 18. Acid-base reactions, considered in detail in Chapter 17, are another. To conclude this chapter, we describe how to determine which equilibria are most important in any particular aqueous solution. 

An aqueous solution of a soluble salt contains cations and anions. These ions often have acid-base properties. Anions that are conjugate bases of weak acids make a solution basic. For example, sodium fluoride dissolves in water to give Na, F, and H2 O as major species. The fluoride anion is the conjugate base of the weak acid HF. This anion establishes a proton transfer equilibrium with water ... 

The solubility product (. sp) describes the equilibrium of a salt dissolving in water. In the laboratory and in industry, solubility equilibria are often exploited in the opposite direction. Two solutions are mixed to form a new solution in which the solubility product of one substance is exceeded. This salt precipitates and can be collected by filtration. Example illustrates how precipitation techniques can be used to remove toxic heavy metals from aqueous solutions. 

By convention, [HA(s)] = [B(s)] = 1. Eqs. (6.1) represent the precipitation equilibria of the uncharged species, and are characterized by the intrinsic solubility equilibrium constant, Sq. The zero subscript denotes the zero charge of the precipitating species. In a saturated solution, the effective (total) solubility S, at a particular pH is defined as the sum of the concentrations of all the compound species dissolved in the aqueous solution ... 

Similarly, concepts of solvation must be employed in the measurement of equilibrium quantities to explain some anomalies, primarily the salting-out effect. Addition of an electrolyte to an aqueous solution of a non-electrolyte results in transfer of part of the water to the hydration sheath of the ion, decreasing the amount of free solvent, and the solubility of the nonelectrolyte decreases. This effect depends, however, on the electrolyte selected. In addition, the activity coefficient values (obtained, for example, by measuring the freezing point) can indicate the magnitude of hydration numbers. Exchange of the open structure of pure water for the more compact structure of the hydration sheath is the cause of lower compressibility of the electrolyte solution compared to pure water and of lower apparent volumes of the ions in solution in comparison with their effective volumes in the crystals. Again, this method yields the overall hydration number. 

Sanemasa, I., Araki, M., Deguchi, T., Nagai, H. (1981) Solubilities of benzene and alkylbenzenes in water. Methods for obtaining aqueous solutions saturated with vapors in equilibrium with organic liquids. Chem. Lett. 2, 255-258. 

The solubility of hydrated copper sulfate (CuS04 5H20) provides a simple example of how the solubility of a compound can be manipulated. CuS04 5H20 itself is very soluble in water, exhibiting an equilibrium solubility of 207 mg/ml at 20°C [44]. This high solubility is due to the dissociation of copper sulfate into its component ions upon dissolution into an aqueous solution ... 

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