Solubility of ionic solids

The solubility of ionic solids in water covers a wide range of values. Knowing the concentration of ions in aqueous solution is important in medicine and in chemical analysis. In this section, you will continue to study equilibrium. You will examine the solubility equilibria of ionic compounds in water. 

The dissolution of an ionic solid in a polar solvent, or the dissociation of an acid, are examples of processes for which AH° and AS0 are both relatively small and are often of equal importance in determining the sign and magnitude of AG°. This was seen to be the case for HF as discussed above, and the thermodynamic functions for the several steps into which its dissociation can be analysed are much greater in magnitude. As discussed further in Section 3.3, the same is true in the analysis of the solubilities of ionic solids. 

Our theme throughout this chapter is to manipulate equilibria to control the solubilities of ionic solids. In the first section we describe general features of the equilibria that govern dissolution and precipitation. In the remaining sections we explore quantitative aspects of these equilibria, including the effects of additional solutes, of acids and bases, and of ligands that can bind to metal ions to form complex ions. 

Solubility of Ionic Solids Our next topic is heterogeneous solution equilibria. A substance submerged in a liquid will generally begin to dissolve. The extremely low chemical potential p of this substance in the pure solvent rises rapidly—remember that p 00 for c 0—with increasing dissolution and therefore concentration. 

Chapter 7 introduced the solubility rules (Section 7.5, Table 7.2), which give us a qualitative description of the solubility of ionic solids. Molecular solids may also be soluble in water depending on whether the solid is polar. Table sugar (C12H22O11), for example, is polar and soluble in water. Nonpolar solids, such as lard and vegetable shortening, are usually insoluble in water. 

Unfortunately, we cannot determine from first principles the relative strengths of these two forces for a given solid. For this reason, among others, we cannot predict in advance the water solubilities of ionic solids, which cover an enormous range of possibilities. 

A wide variety of physical properties are important in the evaluation of ionic liquids (ILs) for potential use in industrial processes. These include pure component properties such as density, isothermal compressibility, volume expansivity, viscosity, heat capacity, and thermal conductivity. However, a wide variety of mixture properties are also important, the most vital of these being the phase behavior of ionic liquids with other compounds. Knowledge of the phase behavior of ionic liquids with gases, liquids, and solids is necessary to assess the feasibility of their use for reactions, separations, and materials processing. Even from the limited data currently available, it is clear that the cation, the substituents on the cation, and the anion can be chosen to enhance or suppress the solubility of ionic liquids in other compounds and the solubility of other compounds in the ionic liquids. For instance, an increase in allcyl chain length decreases the mutual solubility with water, but some anions ([BFJ , for example) can increase mutual solubility with water (compared to [PFg] , for instance) [1-3]. While many mixture properties and many types of phase behavior are important, we focus here on the solubility of gases in room temperature IFs. 

As we noted in Chapter 4, the solubility of ionic compounds in water varies tremendously from one solid to another. The extent to which solution occurs depends on a balance between two forces, both electrical in nature ... 

Young s equation is a plausible, widely used result, but experimental verification is often rendered difficult e.g., the two terms which involve the interface between the solid and the two other phases cannot be measured independently. Furthermore, many complications can arise with contact angle measurements ys values of ionic solids based on contact angle measurements are different from those estimated from solubility (Table 6.1) (cf. Table A.4.1). 

The Kelvin equation may also be applied to the equilibrium solubility of a solid in a liquid. In this case the ratio p/p0 in Equation (40) is replaced by the ratio a/a0, where a0 is the activity of dissolved solute in equilibrium with a flat surface, and a is the analogous quantity for a spherical surface. For an ionic compound having the general formula MmXn, the activity of a dilute solution is related to the molar solubility S as follows ... 

The increase in temperature increases adsorption of non-ionic surfactants on solid surfaces since the solubility of non-ionic surfactants in water decreases with increased temperature. On the other hand, increasing temperature decreases the adsorption of ionic surfactants on solid surfaces because the solubility of ionic surfactant increases with increased temperature. Furthermore, the presence of electrolytes increases the adsorption of ionic surfactants if the solid surface has the same charge as the surfactant head groups. 

As the lattice energies of a series of ionic solids increase, what might you expect to happen to the water solubilities ... 

The solubility of ionic substances in water varies greatly. For example, sodium chloride is quite soluble in water, whereas silver chloride (contains Ag+ and Cl- ions) is only very slightly soluble. The differences in the solubilities of ionic compounds in water typically depend on the relative affinities of the ions for each other (these forces hold the solid together) and the affinities of the ions for water molecules [which cause the solid to disperse (dissolve) in water]. Solubility is a complex issue that we will explore in much more detail in Chapter 17. However, the most important thing to remember at this point is that when an ionic solid does dissolve in water, the ions are dispersed and are assumed to move around independently. 

The high melting points of ionic solids indicate that a lot of energy must be supplied to separate the ions from one another. How is it possible that the ions can separate from one another when soluble ionic compounds are dissolved in water, often with essentially no temperature change ... 

Figure 9 shows that tryptophan can be solubilized with the addition of AOT, octanol, and water. At lower pressures, the surfactant partitions mostly into the liquid phase, and tryptophan is only sparingly soluble. At pressures above 140 bar, the liquid phase disappears as the AOT, octanol, and water form micelles in the fluid phase. The micelles cause the solubility of tryptophan to increase dramatically. The solubility becomes well above 0.1 wt.%, which is quite sufficient for practical applications. At pressures above 200 bar in the solid-fluid region, solubilities vary little with pressure, which is consistent with the relatively constant polarities shown in Figure 7 for similar values of Wq. This ability to adjust solubilities of ionic species at modest temperatures and pressures opens up the possibility of interesting new practical applications. 

In Lesson 6-2,1 mentioned that displacement reactions are a bit misleading at times. Sometimes, a product that is shown on the product side of the equation does not really appear in the physical chemical reaction. The reason for this has to do with the solubility of ionic substances in water. If a particular product is soluble, it will stay dissolved in the aqueous solution. If a product is insoluble, it will appear as a solid precipitate in the test vessel. It is important to know which products stay dissolved in the water, so we can make proper identification of the precipitates that do form as the result of the chemical reactions. Ionic equations are more realistic representations of these reactions that take place in aqueous solution. Ionic equations show the individual ions that exist in solution. When we take an ionic displacement reaction and remove the information that is misleading, we produce a net ionic equation. 

Figure 12.3 shows the temperature dependence of the solubility of some ionic compounds in water. In most but certainly not all cases, the solubility of a solid substance increases with temperature. However, there is no clear correlation between the sign of A/ZsoIj, and the variation of solubility with temperature. For example, the solution process of CaCl2 is exothermic, and that of NH4NO3 is endothermic. But the solubility of both compounds increases with increasing temperature. In general, the effect of temperature on solubility is best determined experimentally. 

Tabulated values of solubilities of ionic salts refer to the maximum amount of solid that will dissolve in a given mass of water to give a saturated solution. 

The negative of this value, 1050 kJ, is the lattice energy of LiF. The lattice energy ( attice) is the enthalpy change that occurs when 1 mol of ionic solid separates into gaseous ions. It indicates the strength of ionic interactions and influences melting point, hardness, solubility, and other properties. 

Fedorov et al. [73FED/SHM] measured the solubility of a solid with the reported composition of a trihydrate, Ni(103)2-3H20, in aqueous lithium perchlo-rate/nitrate mixtures at 298.15 K and ionic strengths from 0.5 to 4.0 M. Their solubility results near 25°C, on extrapolation to I =0, are similar to those from the other two studies. Based on the values obtained from experiments using solutions without nitrate, a value of the solubility product of Ni(I03)2 3H20, log,o K° = - (5.09 0.16) is calculated (see Appendix A). 

The available thermodynamic data are of two types stabihty constants, enthalpy and entropy of reaction for the formation of soluble complexes Th(S04) " " and solubihty data for various solid phases. The two sources are linked because the solubility of the solid phases depends on the chemical speciation, i.e., the sulphate complexes present in the aqueous phase. The analysis of the experimental stability constants has been made using the SIT model however, this method cannot be used to describe the often very high solubility of the solid sulphate phases. In order to describe these data the present review has selected a set of equilibrium constants for the formation of Th(S04) and Th(S04)2(aq) at zero ionic strength based on the SIT model and then used these as constants in a Gibbs energy minimisation code (NONLINT-SIT) for modelling experimental data to determine equilibrium constants for the formation of Th(S04)3 and the solubility products of different thorium sulphate solids phases. 

The solubility of NaCl in water and alcohols also shows an interesting trend and allows us to see the effects of solvent properties. The relevant data are shown in Table 5.3. As the size of the solvent molecules increases and the dielectric constant decreases, the solubihty of NaCl decreases. The size and character of the alkyl group becomes dominant over that of the polar OH group. Accordingly, the solubihty of ionic solids such as NaCl decreases with increasing size of the alkyl group. 

Assume that you have mixed two solutions, and a solid product (a precipitate) forms. How can you find out what the solid is What is its formula There are several possible approaches you can take to answering these questions. For example, we saw in Chapter 7 that we can usually predict the identity of a precipitate formed when two solutions are mixed in a reaction of this type if we know some facts about the solubilities of ionic compounds. 

There are a large number of factors that influence the solubility of atoms. When a solid material is soluble, it is able to be dissolved into a liquid (often, but not always, water). Lattice energy, ion size, and hydration energy play an important role in the solubility of ionic materials. 

Do amino acids exhibit any properties indicating that they behave as zwitterions If so, their behavior should be similar to that of ionic substances. (Section 8.2) Crystalline amino adds have relatively high melting points, usually above 200"C, which is characteristic of ionic solids. Amino adds are far more soluble in vrater than in nonpolar solvents. In addition, the dipole moments of amino acids are large, consistent with a laige separation of charge in the molecule. Thus, the ability of amino adds to ad simultaneously as adds and bases has important effeds on their properties. 

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