Weak Bases and Their Equilibrium Constants

Like the weak acids, a large number of solutes act as weak bases. It is convenient to classify weak bases into two groups, molecules and anions. 

Molecules. As pointed out in Chapter 4, there are many molecular weak bases, including the organic compounds known as amines. The simplest weak base is ammonia, whose reversible Bronsted-Lowry reaction with water is represented by the equation 

This reaction proceeds to only a very small extent. In 0.10 M NH3, the concentrations of NH4+ and OH- are only about 0.0013 M. 

Anions. An anion derived from a weak add is itself a weak base. A typical example is the fluoride ion, F-, which is the conjugate base of the weak acid HF. The reaction of the F- ion with water is 

The OH- ions formed make the solution basic. A 0.10 M solution of sodium fluoride, NaF, has a pH of about 8.1. 

Write an equation to explain why each of the following produces a basic water solution. 

React each basic anion with a water molecule. 

The weak base picks up the proton (H+) and increases its charge by one unit to create its conjugate acid. 

The presence of OH as a product is the reason these anions in water are considered to be basic. 

We showed in Section 13.4 how Ka of a weak acid can be used to calculate [H ] in a solution of that acid. In a very similar way, Kb can be used to find [OH ] in a solution of a weak base. 

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Weak acids and bases exist in equilibrium with their ions, and their equilibrium constants (Ka) are small. The position of equilibrium is measured by the equilibrium constant [equation (4-10)] and Kt. [equation (4-12)]. The dis-... 

Equation 2-7 is more commonly called the Henderson-Ha.sselbalch equation and is the basis for most calculations involving weak acids and bases. It is used to calculate the pH of solutions of weak acids, weak ba.ses. and buffers consisting of weak acids and their conjugate bases or weak bases and their conjugate acids. Because the pK is a modified equilibrium constant, it corrects for the fact that weak acids do itnl completely react with water. 

The carbon-metal bonds of organolithium and organomagnesium compounds have appreciable carbamomc character Carbanions rank among the strongest bases that we 11 see m this text Their conjugate acids are hydrocarbons—very weak acids indeed The equilibrium constants for ionization of hydrocarbons are much smaller than the s for water and alcohols thus hydrocarbons have much larger pA s... 

Equilibrium Constants for Weak Acids and Their Conjugate Bases... 

To predict the pH of mixtures of weak acids or bases and their salts quantitatively, we set up an equilibrium table, as described in Toolbox 10.1. Then we use the acidity or basicity constant to calculate the concentration of hydronium ions present in the solution. The only difference is that now the conjugate acid and base are both present initially, so the first line of the table must have their initial concentrations. For instance, in the mixed acetic acid/sodium acetate solution, both acetic acid and its conjugate base, acetate ions, are present initially. In the ammonia/ammo-nium chloride solution, both the base (ammonia) and its conjugate acid (the ammonium ions) are present initially. 

As usual, [H20] is omitted from the equilibrium constant expression. Table 15.4 lists some typical weak bases and gives their Kb values. (The term base-protonation constant might be a more descriptive name for Kb, but the term base-dissociation constant is still widely used.)... 

Buffers help maintain a relatively constant hydrogen ion concentration. The most common buffers consist of weak acids and their conjugate bases. A buffered solution can resist pH changes because an equilibrium between the buffer s components... 

As the titration begins, mostly HAc is present, plus some H and Ac in amounts that can be calculated (see the Example on page 45). Addition of a solution of NaOH allows hydroxide ions to neutralize any H present. Note that reaction (2) as written is strongly favored its apparent equilibrium constant is greater than lO As H is neutralized, more HAc dissociates to H and Ac. As further NaOH is added, the pH gradually increases as Ac accumulates at the expense of diminishing HAc and the neutralization of H. At the point where half of the HAc has been neutralized, that is, where 0.5 equivalent of OH has been added, the concentrations of HAc and Ac are equal and pH = pV, for HAc. Thus, we have an experimental method for determining the pV, values of weak electrolytes. These p V, values lie at the midpoint of their respective titration curves. After all of the acid has been neutralized (that is, when one equivalent of base has been added), the pH rises exponentially. 

Section 19.1 discusses the Brpnsted theory of acids and bases, which extends the concepts of add and base beyond aqueous solutions and also explains the acidic or basic nature of solutions of most salts. Dissociation constants, the equilibrium constants for the reactions of weak acids or bases with water, are introduced in Section 19.2. The concept of the ionization of covalent compounds is extended to water itself in Section 19.3, which also covers pH, a scale of acidity and basicity. Section 19.4 describes buffer solutions, which resist change in their acidity or basicity even when some strong acid or base is added. Both the preparation and the action of buffer solutions are explained. Section 19.5 discusses the equilibria of acids containing more than one ionizable hydrogen atom per molecule. 

Many organic compounds can act as weak Bronsted-Lowry acids or bases. Their reactions involve the transfer of H+ ions, or protons (Section 10-4). Like similar reactions of inorganic compounds, these acid-base reactions of organic acids and bases are usually fast and reversible. Consequently, we can discuss the acidic or basic properties of organic compounds in terms of equilibrium constants (Section 18-4). 

The polar O-H bond of alcohols makes them weak acids. By the Bronsted-Lowry definition, acids are hydrogen ion donors and bases are hydrogen ion acceptors in chemical reactions. Strong acids are 100% ionized in water and weak acids are only partially ionized. Weak acids establish an equilibrium in water between their ionized and unionized forms. This equilibrium and the strength of an acid is described by the acidity constant, Ka. Ka is defined as the concentrations of the ionized forms of the acids (H30+ and A-) divided by the un-ionized form... 

In 1923, Broensted was the first to develop an acid-base concept that was no longer related to substances, but rather to the function of particles. Acids are proton donors and are capable, with suitable reaction partners, to donate protons to base particles or proton acceptors, i.e. protolysis or proton transfer reaction. For example, HC1 molecules, as acid particles, transfer protons when colliding with H20 molecules (see Fig. 7.3). In this sense, the proton donors of pure sulfuric acid are H2S04 molecules, the acid particles of the sulfuric acid solution are the hydronium ions (or also the hydrogen sulfate ions in concentrated solutions). In weak acids, the protolysis equilibrium is to be considered, equilibria and their constants are well defined. 

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. 

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