Acid-base equilibria buffered solutions

Sections 15.4 and 15.5 outline methods for calculating equilibria involving weak acids, bases, and buffer solutions. There we assume that the amount of hydronium ion (or hydroxide ion) resulting from the ionization of water can be neglected in comparison with that produced by the ionization of dissolved acids or bases. In this section, we replace that approximation by a treatment of acid-base equilibria that is exact, within the limits of the mass-action law. This approach leads to somewhat more complicated equations, but it serves several purposes. It has great practical importance in cases in which the previous approximations no longer hold, such as very weak acids or bases or very dilute solutions. It includes as special cases the various aspects of acid-base equilibrium considered earlier. Finally, it provides a foundation for treating amphoteric equilibrium later in this section. 

Scheme VIII has the form of Scheme II, so the relaxation time is given by Eq. (4-15)—appjirently. However, there is a difference between these two schemes in that L in Scheme VIII is also a participant in an acid-base equilibrium. The proton transfer is much more rapid than is the complex formation, so the acid-base system is considered to be at equilibrium throughout the complex formation. The experiment can be carried out by setting the total ligand concentration comparable to the total metal ion concentration, so that the solution is not buffered. As the base form L of the ligand undergoes coordination, the acid-base equilibrium shifts, thus changing the pH. This pH shift is detected by incorporating an acid-base indicator in the solution.

Buffers are used mainly to control the pH and the acid-base equilibrium of the solute in the mobile phase. They can also be used to influence the retention times of ionizable compounds. The buffer capacity should be maximum and should be uniform in the pH range of 2-8 commonly used in HPLC. The buffers should be soluble, stable, and compatible with the detector employed, e.g., citrates are known to react with certain HPLC hardware components. 

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. 

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... 

However, as in the cases of the other oxoanions considered above, the formed metaborate, BO2, shows appreciable acidic properties. After the pO drop at the equivalence point the oxo-anions continue to fix oxide ions, and the value of the ligand number rises to 5. This means that the second stage of the neutralization process takes place. Neutral orthoborate ion is the product resulting from such an interaction, and it is a strong Lux base. The buffer solutions metaborate/orthoborate exist, mainly, within a 1-2 pO range, and possess buffer numbers of the order of /3 -0.4 to - 0.6. The pK value describing the following equilibrium ... 

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. 

Buffering in aqueous solution is based on a weak acid-base equilibrium ... 

To find the pH of any 0.100 M solution of any acid, base, or 1 1 buffer, we simply put a straightedge at 45° on this plot passing through pH = pX of the acid-base as in Figure 4-1. For other concentrations and buffer ratios we shall have to plot the required material balance Rg curves. Six cases are shown in Figure 4-3 the three from Figure 4-2 and the three for 0.0100 M acid, base, and buffer. Three equilibrium condition ratio (H/X g) lines are shown... 

Buffers contain either a weak acid and its conjugate base or a weak base and its conjugate acid. Thus, a buffer solution contains both an acid species and a base species in equilibrium. To understand the action of a buffer, consider one that contains approximately equal molar amounts of a weak acid HA and its conjugate base A. When a strong acid is added to the buffer, it supplies hydronium ions that react with the base A . 

The values of [HA] and [A ] in this expression are the equilibrium concentrations of acid and base in the solution, not the concentrations added initially. However, a weak acid HA typically loses only a tiny fraction of its protons, and so [HA] is negligibly different from the concentration of the acid used to prepare the buffer, [HA]initia. Likewise, only a tiny fraction of the weakly basic anions A- accept protons, and so [A-] is negligibly different from the initial concentration of the base used to prepare the buffer. With the approximations A ] [base]initia and [HA] [acid]initia, we obtain the Henderson-Hasselbalch equation ... 

The pH is governed by the major solute species present in solution. As strong base is added to a solution of a weak acid, a salt of the conjugate base of the weak acid is formed. This salt affects the pH and needs to be taken into account, as in a buffer solution. Table 11.2 outlines the regions encountered during a titration and the primary equilibrium to consider in each region. 

The analysis carried out in Example reveals one of the key features of buffer solutions The equilibrium concentrations of both the weak acid and its conjugate base are essentially the same as their initial concentrations. 

The pH of a buffer solution depends on the weak acid equilibrium constant and the concentrations of the weak acid and its conjugate base. To show this, we begin by taking the logarithm of the acid equilibrium constant ... 

This equation is exact, but it can be simplified by applying one of the key features of buffer solutions. Any buffer solution contains both members of a conjugate acid-base pair as major species. In other words, both the weak acid and its conjugate base are present in relatively large amounts. As a result, the change to equilibrium, x, is small relative to each initial concentration, and the equilibrium concentrations are virtually the same as the initial leq linitial " = linitial... 

Use the seven-step strategy to calculate the pH of the buffer solution using the buffer equation. Then compare the amount of acid in the solution with the amount of added base. Buffer action is destroyed if the amount of added base is sufficient to react with all the acid.The buffering action of this solution is created by the weak acid H2 PO4 and its conjugate base HP04. The equilibrium constant for this... 

The comparison of I —> N and N —> I may also be explained by the buffered pH in the diffusion layer and leads to an interesting comparison between a process under kinetic control versus one under thermodynamic control. Because the bulk solution in process N —> I favors formation of the ionized species, a much larger quantity of drug could be dissolved in the N —> I solvent if the dissolution process were allowed to reach equilibrium. However, the dissolution rate will be controlled by the solubility in the diffusion layer accordingly, faster dissolution of the salt in the buffered diffusion layer (process I—>N) would be expected. In comparing N—>1 and N —> N, or I —> N and I —> I, the pH of the diffusion layer is identical in each set, and the differences in dissolution rate must be explained either by the size of the diffusion layer or by the concentration gradient of drug between the diffusion and the bulk solution. It is probably safe to assume that a diffusion layer at a different pH than that of the bulk solution is thinner than a diffusion layer at the same pH because of the acid-base interaction at the interface. In addition, when the bulk solution is at a different pH than that of the diffusion layer, the bulk solution will act as a sink and Cg can be eliminated from Eqs. (1), (3), and (4). Both a decrease in the h and Cg terms in Eqs. (1), (3), and (4) favor faster dissolution in processes N —> I and I —> N as opposed to N —> N and I —> I, respectively. 

Acid-base reactions of buffers act either to add or to remove hydrogen ions to or from the solution so as to maintain a nearly constant equilibrium concentration of H+. For example, carbon dioxide acts as a buffer when it dissolves in water to form carbonic acid, which dissociates to carbonate and bicarbonate ions ... 

New NH3/NH4+ buffer When 0.142 mol per liter of HC1 is added to the original buffer presented in (a), it reacts with the base component of the buffer, NH3, to form more of the acid component, NH4+ (the conjugate acid of NH3). Since HC1 is in the gaseous phase, there is no total volume change. A new buffer solution is created with a slightly more acidic pH. In this type of problem, always perform the acid-base limiting reactant problem first, then the equilibrium calculation. 

Each species within a buffer solution participates in an equilibrium reaction, as characterized by an equilibrium constant K. Adding an acid (or base) to a buffer solution causes the equilibrium to shift, thereby preventing the number of protons from changing, itself preventing changes in the pH. The change in the reaction s position of equilibrium is another manifestation of Le Chatelier s principle (see p. 166). 

Buffer and non-buffer solutions are prepared. The pHs of these solutions are determined before and after other substances—usually acids or bases—are added. (See the Equilibrium chapter.)... 

You learned about acids and bases in your previous chemistry course. In this chapter, you will extend your knowledge to learn how the structure of a compound determines whether it is an acid or a base. You will use the equilibrium constant of the reaction of an acid or base with water to determine whether the acid or base is strong or weak. You will apply your understanding of dissociation and pH to investigate buffer solutions solutions that resist changes in pH. Finally, you will examine acid-base titrations that involve combinations of strong and weak acids and bases. 

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