Ionic Salts - Solubility Rules

Although it is possible to deduce the solubility of common inorganic 

1 All nitrates (NO j) and most ethanoates, (acetates, CH3COO ) are soluble. 

2 All common chlorides are soluble except for silver chloride (AgCl) mercury(ll) chloride (HgCl2) and lead(II) chloride (PbCli), although the latter is soluble in hot water. 

3 All common sulfates are soluble, except for barium sulfate (BaS04) and lead sulfate (PbS04). [Calcium sulfate (CaS04) and silver sulfate (Ag2S04) are only slightly soluble.] 

4 All common salts of the metals of Group 1 in the periodic table - sodium, potassium (Na, K) - and all common ammonium salts (NH4) are soluble. 

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Water. It should come as no surprise that ordinary water can be an excellent solvent for many samples. Due to its extremely polar nature, water will dissolve most substances of likewise polar or ionic nature. Obviously, then, when samples are composed solely of ionic salts or polar substances, water would be an excellent choice. An example might be the analysis of a commercial iodized table salt for sodium iodide content. A list of solubility rules for ionic compounds in water can be found in Table 2.1. 

Use the Ionic Compounds activity (eChapter 4.4) to determine the formula of each of the insoluble iron(III) salts. Then, using your knowledge of the solubility rules, write a molecular, ionic, and net ionic equation for an aqueous reaction that would produce each salt. 

In aqueous solutions of ionic compounds, the ions act independently of each other. Soluble ionic compounds are written as their separate ions. We must be familiar with the solubility rules presented in Chapter 8 and recognize that the following types of compounds are strong electrolytes strong acids in solution, soluble metallic hydroxides, and salts. (Salts, which can be formed as the products of reactions of acids with bases, include all ionic compounds except strong acids and bases and metalhc oxides and hydroxides.) Compounds must be both ionic and soluble to be written in the form of their separate ions. (Section 9.1)... 

Notice that in Table 8.1 and Figure 8.3 the term salt is used to mean ionic compound. Many chemists use the terms salt and ionic compound interchangeably. In Example 8.1, we will illustrate how to use the solubility rules to predict the products of reactions among ions. 

The solubility of ionic salts can be predicted by noting the relative sizes of the ions in the material. If there is a large difference in size of the ions the substance is more likely to be soluble. Lithium is the smallest ion among the alkali metals. Based on the rule noted above, which lithium halogen salt is the most soluble and which is the least soluble Explain your reasoning. 

In Section 3.3, we introduced the term solubility to measure how much of a solute a particular solvent can dissolve. This idea is frequently translated into the solubility rules for ionic substances, which we introduced in Section 3.3 (page 91). Those rules classify both silver bromide and silver cyanide as insoluble. But in Example Problem 12.3, we implied that there was a measurable difference in solubility between those two compounds. So we need to refine our notion of what it means for a compound to be insoluble. We will be more accurate if we refer to salts such as silver bromide and silver cyanide as sparingly soluble. Given enough time and a constantly refreshing solvent (in other words, not equilibrium conditions), even the most insoluble salt will dissolve. The dissolution of mountains by rainfall represents such a process, albeit one that requires hundreds of thousands of years. So how can we specify just how soluble an insoluble compound is ... 

To help understand the solubility rules for ionic compounds, think about the magnitudes of the charges on the cation and anion. The higher the charges are, the more strongly the anion and cation attract one another and the less soluble the salt is likely to be. For example, almost all salts composed of 1 + cations and 1 - anions are soluble (Ag salts are a notable exception). Almost all salts composed of 2 +/2 - and 3 +/3 - ion pairs are insoluble (or only slightly soluble). 

We demonstrated precipitation reactions in the context of the solubility rules of chemistry. The example we used was the reaction of silver ion with chloride ion in water. You may recall you learned that the ionic solid silver nitrate dissolves in water, but if a solution of sodium chloride (common table salt) is added to a silver nitrate solution, a fluffy white solid compound, silver chloride, forms. We usually write the equation for that reaction as... 

Physical and ionic adsorption may be either monolayer or multilayer (12). Capillary stmctures in which the diameters of the capillaries are small, ie, one to two molecular diameters, exhibit a marked hysteresis effect on desorption. Sorbed surfactant solutes do not necessarily cover ah. of a sohd iaterface and their presence does not preclude adsorption of solvent molecules. The strength of surfactant sorption generally foUows the order cationic > anionic > nonionic. Surfaces to which this rule apphes include metals, glass, plastics, textiles (13), paper, and many minerals. The pH is an important modifying factor in the adsorption of all ionic surfactants but especially for amphoteric surfactants which are least soluble at their isoelectric point. The speed and degree of adsorption are increased by the presence of dissolved inorganic salts in surfactant solutions (14). 

The alkali metals react with many other elements directly to make binary solids. The alkali halides are often regarded as the most typical ionic solids. Their lattice energies agree closely with calculations although their structures do not all conform to the simple radius ratio rules, as all have the rock salt (NaCl) structure at normal temperature and pressure, except CsCl, CsBr and Csl, which have the eight-coordinate CsCl structure. The alkali halides are all moderately soluble in water, LiF being the least so. 

General Rules for Solubility of Ionic Compounds (Salts) in Water at 25 °C... 

One limitation of the HLB concept is its failure to account for variations in system conditions from that at which the HLB is measured (e.g., temperature, electrolyte concentration). For example, increasing temperature decreases the water solubility of a nonionic surfactant, ultimately causing phase separation above the cloud point, an effect not captured in a temperature-independent HLB value. When both water and oil are present, the temperature at which a surfactant transitions from being water soluble to oil soluble is known as the phase inversion temperature (PIT). Below the PIT, nonionic surfactants are water soluble, while above the PIT. they are oil soluble. Thus, from Bancroft s rule, a nonionic surfactant will form an 0/W emulsion below its PIT and a W/0 emulsion above its PIT. Likewise, increasing salt concentrations reduces the water solubility of ionic surfactant systems. At elevated salt concentrations, ionic surfactants will eventually partition into the oil phase. This is illustrated in Fig. 13. which shows aqueous micelles at lower salt concentrations and oil-phase inverse micelles at higher salt concentrations. Increasing the system temperature will likewise cause this same transition for nonionic surfactant systems. 

Solution Properties. Zwitterionic polymers show interesting aqueous solution behavior. As a general rule, they are insoluble in pure water due to the formation on intra- and interchain ion contacts resulting in an ionically cross-linked network-type structure. Polyampholytes and polybetaines which are not soluble become soluble upon the addition of low molecular weight electrolytes, such as NaCl (Fig. 51). This dissolution process can best be understood in terms of the low molecular weight electrolyte penetrating the ionically cross-linked network whereupon the ions screen the net attractive interactions between the polymer chains and hence promote solubility. The addition of the salt also results in an-tipolylelectrolyte behavior, ie chain expansion upon the addition of the salt. 

With forensic analyses of toxicological and drug evidence, water solubility of the analytes is of paramount concern for sample preparation and extraction. Solubility is also important toxicologically, since it plays a role in determining how, where, and how quickly a drug is absorbed. The like-dissolves-like rule stiU applies but is broadened to include acid-base character and solubility of salts. Thus, target analytes may be ionic compounds, molecular compounds, or in the case of many drugs, salts like cocaine hydrochloride (cocaine HCl). 

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