Ionic solids solubility

Write a balanced net ionic equation. Because H2C03 is a weak acid, it would tend to remain as carbonic acid, especially in the presence of a strong acid like HBr. KBr is an ionic solid, soluble in water therefore, it would exist as separate ions known as spectator ions, which are not written in the net ionic equation. 

When an element has more than one oxidation state the lower halides tend to be ionic whilst the higher ones are covalent—the anhydrous chlorides of lead are a good example, for whilst leadfll) chloride, PbCl2, is a white non-volatile solid, soluble in water without hydrolysis, leadflV) chloride, PbC, is a liquid at room temperature (p. 200) and is immediately hydrolysed. This change of bonding with oxidation state follows from the rules given on p.49... 

The force of attraction between H20 molecules and the ions, which tends to bring the solid into solution. If this factor predominates, the compound is very soluble in water, as is the case with NaCl, NaOH, and many other ionic solids. 

The force of attraction between oppositely charged ions, which tends to keep them in the solid state. If this is the major factor, the water solubility is very low. The fact that CaC03 and BaS04 are almost insoluble in water implies that interionic attractive forces predominate with these ionic solids. 

Frequently we find that the experimentally determined solubility of an ionic solid is larger than that predicted from /Qp. Consider, for example, PbCl2, where the solubility calculated from the relation... 

The effect illustrated in Example 16.6 is a general one. An ionic solid is less soluble in a solution containing a common ion than it is in water (Figure 16.3, p. 438). 

In every case, the alkali metal reacts to form a stable, ionic solid in which the alkali is present as an inert gas-like ion. The product is, in each case, a crystalline substance with high solubility in water. 

Ionic bond, 287, 288 dipole of, 288 in alkali metal halides, 95 vs. covalent, 287 Ionic character, 287 Ionic crystal, 81, 311 Ionic radius, 355 Ionic solids, 79, 81, 311 electrical conductivity, 80 properties of, 312 solubility in water, 79 stability of, 311... 

The great importance of the solubility product concept lies in its bearing upon precipitation from solution, which is, of course, one of the important operations of quantitative analysis. The solubility product is the ultimate value which is attained by the ionic concentration product when equilibrium has been established between the solid phase of a difficultly soluble salt and the solution. If the experimental conditions are such that the ionic concentration product is different from the solubility product, then the system will attempt to adjust itself in such a manner that the ionic and solubility products are equal in value. Thus if, for a given electrolyte, the product of the concentrations of the ions in solution is arbitrarily made to exceed the solubility product, as for example by the addition of a salt with a common ion, the adjustment of the system to equilibrium results in precipitation of the solid salt, provided supersaturation conditions are excluded. If the ionic concentration product is less than the solubility product or can arbitrarily be made so, as (for example) by complex salt formation or by the formation of weak electrolytes, then a further quantity of solute can pass into solution until the solubility product is attained, or, if this is not possible, until all the solute has dissolved. 

An electrolyte is a substance that, in solution, is present as ions. Ionic solids that are soluble in water are electrolytes because the ions become free to move when the solid dissolves (Fig. 1.2). Some electrolytes, however (such as acids), form... 

The equilibrium constant for the solubility equilibrium between an ionic solid and its dissolved ions is called the solubility product, Ksp, of the solute. For example, the solubility product for bismuth sulfide, Bi2S3, is defined as... 

Many substances that participate in aqueous reactions are soluble salts. These ionic solids dissolve in water to give solutions of cations and anions. For almost all salts, there is an upper limit to the amount that will dissolve in water. A salt solution is saturated when the amount dissolved has reached this upper limit of solubility. Any additional salt added to a saturated solution remains undissolved at the bottom of the vessel. When excess solid... 

One of the disadvantages of using a nonaqueous solvent is that in most cases ionic solids are less soluble than in water. There are exceptions to this. For example, silver chloride is insoluble in water, but it is soluble in liquid ammonia. As will be illustrated later, some reactions take place in opposite directions in a nonaqueous solvent and water. 

The solubility of solids in liquids is an important process for the analyst, who frequently uses dissolution as a primary step in an analysis or uses precipitation as a separation procedure. The dissolution of a solid in a liquid is favoured by the entropy change as explained by the principle of maximum disorder discussed earlier. However it is necessary to supply energy in order to break up the lattice and for ionic solids this may be several hundred kilojoules per mole. Even so many of these compounds are soluble in water. After break up of the lattice the solute species are dispersed within the solvent, requiring further energy and producing some weakening of the solvent-solvent interactions. 

For ionic solutions the strain energy seem to be relatively more important than for the metallic alloy systems [38-40] and the size difference between the two components being mixed dominates the energetics, although other factors are also of importance. In cases where the the covalency or ionicity of the components being mixed are largely different a limited solid solubility also must be expected, even... 

The importance of the size of the solute relative to that of the solvent mentioned above is evident also from experimental determinations of the extent of solid solubility in complex oxides and from theoretical evaluations of the enthalpy of solution of large ranges of solutes in a given solvent (e.g. a mineral). The enthalpy of solution for mono-, di- and trivalent trace elements in pyrope and similar systems shows an approximately parabolic variation with ionic radius [44], For the pure mineral, the calculated solution energies always show a minimum at a radius close to that of the host cation. 

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). 

Many ionic solids are also soluble in water. Magnesium fluoride, however, is an exception. It has a very low solubility in water. 

Sodium chloride and other soluble ionic solids dissolve in polar solvents such as water because of ion-dipole forces. An ion-dipole force is the force of attraction between an ion and a polar molecule (a dipole). For example, NaCl dissolves in water because the attractions between the Na and Cl ions and the water molecules provide enough energy to overcome the forces that bind the ions together. Figure 4.14 shows how ion-dipole forces dissolve any type of soluble ionic compound. 

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. 

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