CHAPTER Acid-Base Equilibria

C18-0040. List all the types of calculations described in Chapter 18 in which acid-base equilibrium expressions play a role. 

Our goal in this chapter is to help you continue learning about acid-base equilibrium systems and, in particular, buffers and titrations. If you are a little unsure about equilibria and especially weak acid-base equilibria, review Chapters 14 and 15. You will also learn to apply the basic concepts of equilibria to solubility and complex ions. Two things to remember (1) The basic concepts of equilibria apply to all the various types of equilibria, and (2) Practice, Practice, Practice. 

Recall from Chapter 4 (text Section 4.6) that an acid-base equilibrium favors formation of the weaker acid and base. Also remember that the weaker acid forms the stronger conjugate base, and vice versa. 

Having a conceptual understanding of the effect is a good starting point, but we still need to be able to understand the quantitative relationships between the different components in the equilibrium mixture. In this section, we will see how to deal with the common-ion effect in acid-base equilibrium problems. You will find that these problems are very similar to the weak acid problems earlier in the chapter. 

The alkoxide doesn t have to be made first, though, because alcohols dissolved in basic solution are at least partly deprotonated to give alkoxide anions. How much alkoxide is present depends on the pH of the solution and therefore the pKa of the base (Chapter 8), but even a tiny amount is acceptable because once this has added it will be replaced by more alkoxide in acid-base equilibrium with the alcohol. In this example, allyl alcohol adds to pent-2-enal, catalysed by sodium-hydroxide in the presence of a buffer. 

In this chapter we have encountered many different situations involving aqueous solutions of acids and bases, and in the next chapter we will encounter still more. In solving for the equilibrium concentrations in these aqueous solutions, you may be tempted to create a pigeonhole for each possible situation and to memorize the procedures necessary to deal with each particular situation. This approach is just not practical and usually leads to frustration Too many pigeonholes are required, because there seems to be an infinite number of cases. But you can handle any case successfully by taking a systematic, patient, and thoughtful approach. When analyzing an acid-base equilibrium problem, do not ask yourself how a memorized solution can be used to solve the problem. Instead, ask yourself this question What are the major species in the solution, and how does each behave chemically ... 

In the course of acute or chronic liver disease, the biochemical functions of the liver may be compromised indefinitely the outcome is decompensated liver insufficiency. (s. pp 277, 381) (s. tab. 20.4) The stage of decompensation is synonymous with the onset of life-threatening complications. These mainly take the form of hepatic encephalopathy with transition to hepatic coma (see chapter 15), oedema and ascites with imbalance of the electrolytes and the acid-base equilibrium (see chapter 16) through to the hepatorenal syndrome (see chapter... 

Many real-world applications of chemistry and biochemistry involve fairly complex sets of reactions occurring in sequence and/or in parallel. Each of these individual reactions is governed by its own equilibrium constant. How do we describe the overall progress of the entire coupled set of reactions We write all the involved equilibrium expressions and treat them as a set of simultaneous algebraic equations, because the concentrations of various chemical species appear in several expressions in the set. Examination of relative values of equilibrium constants shows that some reactions dominate the overall coupled set of reactions, and this chemical insight enables mathematical simplifications in the simultaneous equations. We study coupled equilibria in considerable detail in Chapter 15 on acid-base equilibrium. Here, we provide a brief introduction to this topic in the context of an important biochemical reaction. 

In this text, we use the symbol H30 in those chapters that deal with acid/base equilibria and acid/base equilibrium calculations. In the remaining chapters, we simplify to the more convenient H", with the understanding that this symbol represents the hydrated proton. 

The radiolysis of water leads to the formation of the three radical species OH (hydroxyl radical), H (hydrogen atom) and e q" (hydrated electron), within the nanosecond time scale [3]. In aerated medium (that is in the presence of dioxygen, Oj concentration dissolved in water being 2 x 10 mol I ), the free radicals H and e are replaced by HO2 and Oj radical species, respectively, which are related by an acid/base equilibrium (pK (HO27O2 ) = 4.8).The radiolytic yields (G-values expressed in moles per Joule) of each radical species are well known 2.8 x 10 mol J and 3.4 x 10 mol J , respectively for OH and O2 free radicals at pH = 7 (Chapter 1) [3]. 

The concepts and equations of acid-base dissociation have referred chiefly to aqueous solutions. Recently, interest in the behavior of acids and bases in solvents other than water has increased considerably. The classical definition of an acid and a base, which is satisfactory for water solutions, is too limited for other solvents. Because of the great importance of the general question of the acid-base equilibrium, the clear and fruitful views of Bronsted are exhaustively considered in a special (fourth) chapter. ... 

Before studying this chapter, review Section 5.4 on limiting quantities problems and, before Section 10.2, Acid-Base Equilibrium, review net ionic equations from the textbook. 

The acidity or basicity of a solution is frequently an important factor in chemical reactions. The use of buffers of a given pH to maintain the solution pH at a desired level is very important. In addition, fundamental acid-base equihbria are important in understanding acid-base titrations and the effects of acids on chemical species and reactions, for example, the effects of complexation or precipitation. In Chapter 6, we described the fundamental concept of equilibrium constants. In this chapter, we consider in more detail various acid-base equilibrium calculations, including weak acids and bases, hydrolysis, of salts of weak acids and bases, buffers, polyprotic acids and their salts, and physiological buffers. Acid-base theories and the basic pH concept are reviewed first. 

The importance of water as a medium for inorganic reactions stems not only from the fact that it is far more readily available than any other solvent, but also because of the abundance of accurate physicochemical data for aqueous solutions compared with the relative scarcity of such data for solutions in non-aqueous solvents. This chapter is concerned mainly with equilibria and we begin by reviewing calculations involving acid-base equilibrium constants. 

Chapter 15 Applications of Aqueous Equilibria" was a long chapter, which dealt with difficult material. In order to make this chapter more manageable for students, we have split this chapter in two "Chapter 15 Acid-Base Equilibrium" and "Chapter 16 Solubility and Complex Ion Equilibrium."... 

The last step in complexity that we will explore in detail for the construction of diagrams to illustrate redox equilibria involves the addition of heterogeneous equilibria to redox and acid-base equilibrium diagrams. We will illustrate this system with a ps-pH diagram for iron species in aqueous solution containing no anions other than hydroxide, We will expand on this diagram later in this chapter during the discussion of iron chemistry. 

Solutions Manual (0-13-147882-6) The Solutions Manual, prepared by Jan W. Simek of California Polytechnic State University, contains complete solutions to all the problems. The Solutions Manual also gives helpful hints on how to approach each kind of problem. This supplement is a useful aid for any student, and it is particularly valuable for students who feel they understand the material but need more help with problem solving. Appendix 1 of the Solutions Manual summarizes the lUPAC system of nomenclature. Appendix 2 reviews and demonstrates how acidity varies with structure in organic molecules, and how one can predict the direction of an acid-base equilibrium. Brief answers to many of the in-chapter problems are given at the back of this book. These answers are sufficient for a student on the right track, but they are of limited use to one who is having difficulty working the problems. 

Large hydrogen isotope effects were found experimentally very soon after the discovery of deuterium in 1932, and there is now an extensive literature on the subject, recently supplemented by work with tritium. Much of this relates to the kinetics of reactions involving proton transfer, but there is also a large amount of information on acid-base equilibria. A recent review estimated that about 300 papers on isotope effects are now published each year, and a considerable proportion of these relate to the isotopes of hydrogen. The present chapter will be devoted to the effect of isotope substitution on acid-base equilibrium constants, while Chapter 12 will deal with kinetic isotope effects in proton transfer reactions. 

Acid-base chemistry is quite important in biological systems, and acid-base reactions drive many common processes. As for this chapter, one useful example of the acid-base equilibrium process is metabolic acidosis and metabolic... 

E. J. King, Acid-Base Equilibrium, Macmillan, New York, 1965, Chapter 9. 

In this chapter, we treat a variety of acid-base equilibrium situations to give practice in using pH calculations and concepts in a broad span of science. 

In aqueous solutions, two ions have dominant roles. These ions, the hydronium ion, H30 (or hydrogen ion, H" ), and the hydroxide ion, (OH ), are available in any aqueous solution as a result of the self-ionization of water, a reaction of water with itself, which we will describe in the next section. This will also give us some background to acid-base equilibrium calculations, which we will discuss in Chapter 17. 

In the chapter Acid-Base Titration and pH, you learned that the selfionization of water is an equilibrium reaction. 

Hypochlorite is the most inexpensive and the most effective disinfectant known. Its activity is mainly due to the presence of hypochlorous acid (see Chapter 16). The strongest biocidal activity occurs at pH 6—8, where hypochlorous acid predominates because of the acid-base equilibrium, but products need a considerably higher pH in order to ensure hypochlorite stability. 

As pointed out in Chapter 4, an acid-base indicator is useful in determining the equivalence point of an acid-base titration. This is the point at which reaction is complete equivalent quantities of acid and base have reacted. If the indicator is chosen properly, the point at which it changes color (its end point) coincides with the equivalence point To understand how and why an indicator changes color, we need to understand the equilibrium principle involved. 

The discussion of acid-base titrations in Chapter 4 focused on stoichiometry. Here, the emphasis is on the equilibrium principles that apply to the acid-base reactions involved. It is convenient to distinguish between titrations involving—... 

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