Complex Ion and Precipitation Equilibria

Thus far, we have looked at the equilibrium established when acids, bases or both of these are added to water. 

In this chapter we will discuss the equilibrium of complex ion formation (Section 15.1), Solubiiity (Section 15.2) and precipitation (Section 15.3), discussed in Chapter 4, are reexamined with a focus on an equilibrium constant Ksp. We will also learn how to dissolve precipitates (Section 15.4) by making them react with strong acids or complexing agents forming complex ions, 

The equilibria involving complex ions and precipitates have applications in geology, medicine, and agriculture. In chemistry, you are more likely to meet up with these equilibria in the laboratory when you carry out experiments in qualitative analysis. 

Recall from Chapter 13 the Lewis model for the definition of an acid and a base. Many of the electron acceptors for Lewis acids are transition metal ions (Lewis acids). These metal ions combine with molecules (e.g., H2O, NH3) or other ions (e.g., CC, OH ) that 

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In this chapter, we describe the quantitative effects of acidity and complex-ation in precipitation equilibria and discuss precipitation titrations using silver nitrate and barium nitrate titrants with different kinds of indicators and their theory. You should review fundamental precipitation equilibria described in Chapter 10. Most ionic analytes, especially inorganic anions, e conveniently determined using ion chromatography (Chapter 21), but for high concentrations more precise determinations can be made by precipitation titration when applicable. 

The fact that the ion activities are measured rather than concentrations is ambivalent from the point of view of practical measurements. It is a great advantage for speciation purposes and for study of acid—base complexation and precipitation equilibria in solution. On the other hand, it is a drawback when the total content of an analyte in a sample is to be foimd, as great care must be taken to compensate or correct for interactions of the analyte with the sample matrix and the ambient atmosphere during calibration and measurement itself (cf. the problem of complexation of fluoride with iron(III) and alumi-num(III) ions in analyses of natural waters, or oxidation of sulfide by atmospheric oxygen). 

Equilibria Involving Compbx Ions 643 Formation of Complex Ions 643 Complex Ions and Solubility of Precipitates 645 CHAPTER REVIEW GUIDE 646 PROBLEMS 648... 

Complex formation, selective precipitation, and control of the pH of a solution all play important roles in the qualitative analysis of the ions present in aqueous solutions. There are many different schemes of analysis, but they follow the same general principles. Let s think through a simple procedure for the identification of five cations by following the steps that might be used in the laboratory. We shall see how each step makes use of solubility equilibria. 

Subject areas for the Series include solutions of electrolytes, liquid mixtures, chemical equilibria in solution, acid-base equilibria, vapour-liquid equilibria, liquid-liquid equilibria, solid-liquid equilibria, equilibria in analytical chemistry, dissolution of gases in liquids, dissolution and precipitation, solubility in cryogenic solvents, molten salt systems, solubility measurement techniques, solid solutions, reactions within the solid phase, ion transport reactions away from the interface (i.e. in homogeneous, bulk systems), liquid crystalline systems, solutions of macrocyclic compounds (including macrocyclic electrolytes), polymer systems, molecular dynamic simulations, structural chemistry of liquids and solutions, predictive techniques for properties of solutions, complex and multi-component solutions applications, of solution chemistry to materials and metallurgy (oxide solutions, alloys, mattes etc.), medical aspects of solubility, and environmental issues involving solution phenomena and homogeneous component phenomena. 

The ion concentrations that appear in the expression for the solubility product refer only to the simple ions in solution, and do not include the material in the precipitate because solids are not included in K expressions. Additional equilibria may exist between the simple ions and complexes in solution, as in the case of soluble complexes forming. Such equilibria would be governed by their own stability constants. 

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. 

Utilizes complex formation, selective precipitation, and pH control Uses solubility equilibria to remove and identify ions selectively... 

Consider just a few cases of aqueous equilibria. The magnificent formations i n limestone caves and the vast expanses of oceanic coral reefs result from subtle shifts in carbonate solubility equilibria. Carbonates also influence soil pH and prevent acidification of lakes by acid rain. Equilibria involving carbon dioxide and phosphates help organisms maintain cellular pH within narrow limits. Equilibria involving clays in soils control the availability of ionic nutrients for plants. The principles of ionic equilibrium also govern how water is softened, how substances are purified by precipitation of unwanted ions, and even how the weak acids in wine and vinegar influence the delicate taste of a fine French sauce. In this chapter, we explore three aqueous ionic equilibrium systems acid-base buffers, slightly soluble salts, and complex ions. 

In this chapter we will discuss the formation of solids from an aqueous solution and the resulting equilibria. We will also show how selective precipitation and the formation of complex ions can be used to do qualitative analysis. 

The equilibrium equations that normally have to be considered in the EKR modeling of a soil contaminated by heavy metals can be classified into one of the following categories complex formation reactions, precipitation of the metal hydroxides or of other species, ion exchange reactions, surface complexation reactions, etc. Anyway, the autoionization of water always has to be considered and the precipitation of carbonates, together with the carbonate-bicarbonate equilibrium, should normally also be considered. However, the above equations have only considered the species in aqueous phase, so if a species precipitates, a new master species has to be included in this equilibrium system, whose concentration would be the amount of the precipitated species per unit volume of water. This additional degree of freedom is constrained by the solubility product constant of the precipitate (KO, because the new solid phase is in equilibrium with the aqueous phase. If there exists Np precipitated species, the pure-phase equilibria can be represented with the following equation ... 

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