Weak-Base Equilibrium

For the nitration of the very weak base, acetophenone, there is reasonable agreement between observed and calculated activation parameters, and there is no doubt that nitration of the free base occurs at acidities below that of maximum rate. In this case the equilibrium concentration of free base is much greater than in the examples just discussed and there is no question of reaction upon encounter. ... 

The equilibrium constant for equation 6.13 is K. Since equation 6.13 is obtained by adding together reactions 6.11 and 6.12, may also be expressed as the product of Ka for CH3COOH and Kb for CH3COO-. Thus, for a weak acid, HA, and its conjugate weak base, A-,... 

Besides equilibrium constant equations, two other types of equations are used in the systematic approach to solving equilibrium problems. The first of these is a mass balance equation, which is simply a statement of the conservation of matter. In a solution of a monoprotic weak acid, for example, the combined concentrations of the conjugate weak acid, HA, and the conjugate weak base, A , must equal the weak acid s initial concentration, Cha- ... 

Equation 6.44 is written in terms of the concentrations of CH3COOH and CH3COO- at equilibrium. A more useful relationship relates the buffer s pH to the initial concentrations of weak acid and weak base. A general buffer equation can be derived by considering the following reactions for a weak acid, HA, and the salt of its conjugate weak base, NaA. 

Chemical Limitations to Beer s Law Chemical deviations from Beer s law can occur when the absorbing species is involved in an equilibrium reaction. Consider, as an example, an analysis for the weak acid, HA. To construct a Beer s law calibration curve, several standards containing known total concentrations of HA, Cmt, are prepared and the absorbance of each is measured at the same wavelength. Since HA is a weak acid, it exists in equilibrium with its conjugate weak base, A ... 

PK. — the negative logarithm of the equilibrium constant for acids or bases. This parameter is an indicator of the strength of an acid or base. Strong acids, such as H2SO4, and HCl, have low pK s (i.e., -1.0) while strong bases such as KOH and NaOH, have pK s close to 14.0. Weak acids and weak bases fall in the intermediate range. 

Suppose now that the pH is controlled by a weak base buffer, the equilibrium being written BH B + H, where B signifies a neutral base. The apparent dissociation constant is = (H )[B]/[BH ]. Following the earlier argument, we obtain... 

As an example of enolate-ion reactivity, aldehydes and ketones undergo base-promoted o halogenation. Even relatively weak bases such as hydroxide ion are effective for halogenation because it s not necessary to convert the ketone completely into its enolate ion. As soon as a small amount of enolate is generated, it reacts immediately with the halogen, removing it from the reaction and driving the equilibrium for further enolate ion formation. 

Equilibrium constants of weak bases can be measured in the laboratory by procedures very much like those used for weak acids. In practice, though, it is simpler to take advantage of a simple mathematical relationship between Kb for a weak base and Ka for its conjugate acid. This relationship can be derived by adding together the equations for the ionization of the weak acid HB and the reaction of the weak base B- with water ... 

In Chapter 13 we dealt with the equilibrium established when a single solute, either a weak acid or a weak base, is added to water. This chapter focuses on the equilibrium established when two different solutes are mixed in water solution. These solutes may be—... 

Consider now the salt of a strong acid and a weak base class (3). Here the initial high concentration of cations M + will be reduced by combination with the hydroxide ions of water to form the little-dissociated base MOH until the equilibrium ... 

Case 2. Salt of a strong acid and a weak base. The hydrolytic equilibrium is represented by ... 

All these species are in ceaseless dynamic equilibrium. Similarly, for a solution of a weak base, we visualize... 

The rest of this chapter is a variation on a theme introduced in Chapter 9 the use of equilibrium constants to calculate the equilibrium composition of solutions of acids, bases, and salts. We shall see how to predict the pH of solutions of weak acids and bases and how to calculate the extent of deprotonation of a weak acid and the extent of protonation of a weak base. We shall also see how to calculate the pH of a solution of a salt in which the cation or anion of the salt may itself be a weak acid or base. 

We calculate the pH of solutions of weak bases in the same way as we calculate the pH of solutions of weak acids—by using an equilibrium table. The protonation equilibrium is given in Eq. 9. To calculate the pH of the solution, we first calculate the concentration of OH ions at equilibrium, express that concentration as pOH, and then calculate the pH at 25°C from the relation pH + pOH = 14.00. For very weak or very dilute bases, the autoprotolysis of water must be taken into consideration. 

Proton transfer equilibrium is established as soon as a weak base is dissolved in water, and so we can calculate the hydroxide ion concentration from the initial concentration of the base and the value of its basicity constant. Because the hydroxide ions are in equilibrium with the hydronium ions, we can use the pOH and pKw to calculate the pH. 

To calculate the pH of a solution of a weak base, set up an equilibrium table to calculate pOH from the value ofKh and convert that pOH into pH by using pH + pOH = 14.00. 

To calculate the pH of a salt solution, we can use the equilibrium table procedure described in Toolboxes 10.1 and 10.2—an acidic cation is treated as a weak acid and a basic anion as a weak base. However, often we must first calculate the Ka or Kh for the acidic or basic ion. Examples 10.10 and 10.11 illustrate the procedure. 

For each of the following weak bases, write the proton transfer equilibrium equation and the expression for the equilibrium constant Kh. Identify the conjugate acid, write the appropriate proton transfer equation, and write the expression for the acidity constant Ka. (a) (CH3)2NH, dimethylamine ... 

Step 5 Use an equilibrium table to find the H.O concentration in a weak acid or the OH concentration in a weak base. Alternatively, if the concentrations of conjugate acid and base calculated in step 4 are both large relative to the concentration of hydronium ions, use them in the expression for /<, or the Henderson—Hasselbalch equation to determine the pH. In each case, if the pH is less than 6 or greater than 8, assume that the autoprotolysis of water does not significantly affect the pH. If necessary, convert between Ka and Kh by using Kw = KA X Kb. 

The major species in an aqueous solution determine which categories of equilibria are important for that solution. Each major species present in the solution must be examined in light of these general categories. Are any of the major species weak acids or weak bases Are there ions present that combine to form an insoluble salt Do any of the major species participate in more than one equilibrium Any chemical reaction can approach equilibrium from either direction. Consequently, there are six different t q)es of aqueous equilibria in which major species are reactants ... 

C16-0105. Write the equilibrium reaction and equilibrium constant expression for each of the following processes (a) Trimethylamine, (CH3)3 N, a weak base, is added to water, (b) Hydrofluoric acid, HF, a weak acid, is added to water, (c) Solid calcium sulfate, CaSOq, a sparingly soluble salt, is added to water. 

Most acids and bases are weak. A solution of a weak acid contains the acid and water as major species, and a solution of a weak base contains the base and water as major species. Proton-transfer equilibria determine the concentrations of hydronium ions and hydroxide ions in these solutions. To determine the concentrations at equilibrium, we must apply the general equilibrium strategy to these types of solutions. 

Ammonia is an example of a weak base. A weak base generates hydroxide ions by accepting protons from water but reaches equilibrium when only a fraction of its molecules have done so. The equilibrium constant for this type... 

The strong base is a soluble hydroxide that ionizes completely in water, so the concentration of OH matches the 0.25 M concentration of the base. For the weak base, in contrast, the equilibrium concentration of OH is substantially smaller than the 0.25 M concentration of the base. At any instant, only 0.8% of the ammonia molecules have accepted protons from water molecules, producing a much less basic solution in which OH is a minor species. The equilibrium concentration of unproton-ated ammonia is nearly equal to the Initial concentration. Figure 17-7 summarizes these differences. 

Besides water, the most common weak base is ammonia, NH3, whose proton transfer equilibrium with water appears in Section 16-. Many other weak bases are derivatives of ammonia called amines, hi these organic compounds, one, two, or three of the N—H bonds in ammonia have been replaced with N—C bonds. The nitrogen atom in an amine, like its counterpart in ammonia, has a lone pair of electrons that can form a bond to a proton. Water does not protonate an amine to an appreciable extent, so all amines are weak bases. Table 17-4 lists several examples of bases derived from ammonia. 

Salts that contain cations of weak bases are acidic. For example, the ammonium cation Is the conjugate acid of ammonia. When ammonium salts dissolve in water, NH4 ions transfer protons to H2 O molecules, generating H3 O and making the solution slightly acidic NH4" ((2 q) + H2 0(/) NH3(c2 q) + H3 O (a q) The equilibrium constant for this reaction can be calculated from Equation and for ammonia (Example ) ... 

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