Acids and Bases—pH Calculation

The common-ion effect is an application of Le Chatelicr s principle to equilibrium systems of slightly soluble salts. A buffer is a solution that resists a change in pH if we add an acid or base. We can calculate the pH of a buffer using the Henderson-Hasselbalch equation. We use titrations to determine the concentration of an acid or base solution. We can represent solubility equilibria by the solubility product constant expression, Ksp. We can use the concepts associated with weak acids and bases to calculate the pH at any point during a titration. 

Example 3.7. Mixture of Acid and Base (Buffer) Calculate the pH ol a... 

Alternatively, we can use the Henderson-Hasselbalch equation with the initial concentrations of acid and base to calculate pH directly ... 

Chapter 14, Acids and Bases, discusses acids and bases and their strengths, conjugate acid-base pairs, the dissociation of weak acids and bases and water, pH and pOH, and buffers. Acid-base titration uses the neutralization reactions between acids and bases to calculate quantities of acid in a sample. Section 14.9, Acid-Base Properties of Salt Solutions, has been deleted. Combining Ideas from Chapters 11,12,13, and 14 follows as an interchapter problem set. 

Acids and bases are a big part of organic chemistry but the emphasis is much different from what you may be familiar with from your general chemistry course Most of the atten tion m general chemistry is given to numerical calculations pH percent loniza tion buffer problems and so on Some of this returns m organic chemistry but mostly we are concerned with the roles that acids and bases play as reactants products and catalysts m chemical reactions We 11 start by reviewing some general ideas about acids and bases... 

The fact that strong acids and bases are completely ionized in water makes it relatively easy to calculate the pH and pOH of their solutions (Example 13.3). 

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. 

Suppose we were asked to estimate the pH of 1.0 X 10 x m HCl(aq). If we used the techniques of Example 10.3 to calculate the pH from the concentration of the acid itself, we would find pH = 8.00. That value, though, is absurd, because it lies on the basic side of neutrality, whereas HC1 is an acid The error stems from there being two sources of hydronium ions, whereas we have considered only one. At very low acid concentrations, the supply of hydronium ions from the autoprotolysis of water is close to the supply provided by the very low concentration of HC1, and both supplies must be taken into account. The following two sections explain how to take autoprotolysis into account, first for strong acids and bases and then for weak ones. 

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 above examples assume that the strong base KOH is completely dissociated in solution and that the concentration of OH ions was thus equal to that of the KOH. This assumption is valid for dilute solutions of strong bases or acids but not for weak bases or acids. Since weak electrolytes dissociate only slightly in solution, we must use the dissociation constant to calculate the concentration of [H" ] (or [OH ]) produced by a given molarity of a weak acid (or base) before calculating total [H" ] (or total [OH ]) and subsequendy pH. 

In the case of dissociating or ionizing organic chemicals such as organic acids and bases, e.g., phenols, carboxylic acids and amines, it is desirable to calculate the concentrations of ionic and non-ionic species, and correct for this effect. A number of authors have discussed and reviewed the effect of pH and ionic strength on the distribution of these chemicals in the environment, including Westall et al. (1985), Schwarzenbach et al. (1988), Jafvert et al. (1990), Johnson and Westall (1990) and the text by Schwarzenbach, Gschwend and Imboden (1993). 

When an acid in solution is exactly neutralized with a base the resulting solution corresponds to a solution of the salt of the acid-base pair. This is a situation which frequently arises in analytical procedures and the calculation of the exact pH of such a solution may be of considerable importance. The neutralization point or end point in an acid-base titration is a particular example (Chapter 5). Salts may in all cases be regarded as strong electrolytes so that a salt AB derived from acid AH and base B will dissociate completely in solution. If the acid and base are strong, no further reaction is likely and the solution pH remains unaffected by the salt. However if either or both acid and base are weak a more complex situation will develop. It is convenient to consider three separate cases, (a) weak acid-strong base, (b) strong acid-weak base and (c) weak acid-weak base. 

Most pyrethroids undergo acid- and base-catalyzed hydrolysis to form the corresponding acid and alcohol (Fig. la), typically with U-shaped pH-rate profiles [8, 40]. The hydrolysis of pyrethroids in water basically obeys first-order kinetics with a half-life simply calculated from hydrolysis rate constant (A obs) as 0.693/kobs. Pyrethroids are generally stable under the acidic and neutral conditions at pH 4—7,... 

We can prepare a buffer of almost any pH provided we know the pAa of the acid and such values are easily calculated from the Ka values in Table 6.5 and in most books of physical chemistry and Equation (6.50). We first choose a weak acid whose pKa is relatively close to the buffer pH we want. We then need to measure out accurately the volume of acid and base solutions, as dictated by Equation (6.50). 

Our goal in this chapter is to help you understand the equilibrium systems involving acids and bases. If you don t recall the Arrhenius acid-base theory, refer to Chapter 4 on Aqueous Solutions. You will learn a couple of other acid-base theories, the concept of pH, and will apply those basic equilibrium techniques we covered in Chapter 14 to acid-base systems. In addition, you will need to be familiar with the log and 10 functions of your calculator. And, as usual, in order to do well you must Practice, Practice, Practice. 

In the next section, you will focus on the equilibrium of water. You will discover how the pH scale is related to the concentrations of the ions that form when water dissociates. As well, you will learn how to calculate the pH values of solutions of weak acids and bases. 

In an acid-base titration, you carefully measure the volumes of acid and base that react. Then, knowing the concentration of either the acid or the base, and the stoichiometric relationship between them, you calculate the concentration of the other reactant. The equivalence point in the titration occurs when just enough acid and base have been mixed for a complete reaction to occur, with no excess of either reactant. As you learned in Chapter 8, you can find the equivalence point from a graph that shows pH versus volume of one solution added to the other solution. To determine the equivalence point experimentally, you need to measure the pH. Because pH meters are expensive, and the glass electrodes are fragile, titrations are often performed using an acid-base indicator. 

Would suspect disproportionation to be second-order. Confirm by examining first three entries. Calculate 2nd order rate constant for each entry. Since acid and base forms, NiLH and NiL, are unreactive, then maximum rate close to pH 4 arises from reaction of NiLtf + with NiL + (k). Use equation analogous to (1.231) and confirm k = (3.4 0.5) X lO M- s. ... 

Speciation calculations can be performed for the weak acids and bases in a feshion similar to that presented earlier for Fe(III). The results of these calculations as a function of pH are shown in Figure 5.19. At the pH of seawater, the dominant species are carbonate, bicarbonate, ammonium, hydrogen phosphate, dihydrogen phosphate, and boric and silicic acid. In waters with low O2 concentrations, significant concentrations of HS can be present. 

The extent to which the pH of a solution is buffered against additions or removals of protons is measured by the solution s pH buffer capacity. This is defined as the amount of strong acid or base required to produce unit change in pH. The buffering depends on the transfer of protons between donors and acceptors, i.e. Bronsted acids and bases, which form conjugate acid-base pairs. The pH buffer capacity of a solution is calculated from the buffer capacities of the individual acid-base pairs present. 

These calculations are for the pH of weak acids and weak bases. It is well worth comparing the hgures we calculated above for strong acids and bases. Thus, a 0.1 M solution of the strong acid HCl had pH 1, and a 0.1 M solution of the strong base NaOH had pH 13. 

Calculation of pH values of aqueous solutions of strong and weak acids and bases... 

There are several ways of identifying whether a compound is an acid or a base, depending on what it does with protons and electrons pH and pOH calculations, along with the values of dissociation constants K and /CJ, can help chemists determine the properties of these acids and bases. 

This technique uses both direct and back titrations of weak acids and bases. Values of are obtained directly. In purely aqueous media, over the pH range 2-10, the titration of dilute (0.005 to 0.05 M) solutions of weak monovalent acids and bases with a glass electrode can lead to reliable thermodynamic pKs. Over this pH interval, the activity coefficients of the ionic species can be calculated by means of the Debye-Hiickel equation. Also, the activity coefficients of the neutral species remain essentially constant and... 

Tables 8.1 and 8.2 give calculated ct,a and (l-a,a) values, respectively, for various acids and bases in water at pH 7. Fig. 8.1 shows schematically the speciation of a given acid (or base) as a function of pH. Some example calculations are given in Illustrative Example 8.1. It should be reemphasized that the neutral and ionic forms of a given neutral acid (base) behave very differently in the environment. Depending on the process considered, either the neutral or ionic species may be the dominant factor in the compound s reactivity, even if the relative amount of that...

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