Acids, Bases, and Buffers

Beyond simple pH and pOH lie titrations and buffers. Titrations allow you to determine the concentration of acids and bases. Buffers maintain the pH of a solution by reacting to changes and neutralizing them. 

Suppose we need to make up a buffer of a particular pH. For instance, we might be culturing bacteria and need to maintain a precise pH. To choose the most appropriate buffer system, we need to know the value of the pH at which a given buffer stabilizes a solution. Because buffer solutions are merely mixed solutions, we can calculate their pH in the same way we did in Example 11.1. There is no need to distinguish between acid and base buffers in the calculation, because the principles are the same for both. 

AQUEOUS SOLUTIONS AND ACID-BASE CHEMISTRY Weak Acids and Bases—Buffers... 

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

Fig. 5.6 Anodic polarization curves for iron dissolution (solid curves) and for total current density of iron plus oxygen evolution (dashed curves) after 1 h at steady state in deaerated 0.15 M Na3P04 solution. Indicated pH obtained by use of acid and base buffers and additions of H2S04 or NaOH. Redrawn from Ref 5...

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. 

Living systems depend on buffers for pH control most of these are complex buffers that contain a mixture of various acids and bases. Buffers can be made from weak acids/bases and the salts of other weak acids/bases. For example, we could make an acidic buffer from acetic acid and sodium bicarbonate (NaHC03, a salt of the weak acid carbonic acid). Or, we could make a basic buffer from NH3 and NH2(CH3)2Br (the salt of the weak base di-methylamine, HN(CH3)2). But it is more difficult to calculate pH s or other concentration values from these mixed buffer systems. We have limited our discussion and detailed examples to buffers that contain a weak acid/base and the salt of icfiX. same acid/base. 

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 role that acid and base catalysts play can be quantitatively studied by kinetic techniques. It is possible to recognize several distinct types of catalysis by acids and bases. The term specie acid catalysis is used when the reaction rate is dependent on the equilibrium for protonation of the reactant. This type of catalysis is independent of the concentration and specific structure of the various proton donors present in solution. Specific acid catalysis is governed by the hydrogen-ion concentration (pH) of the solution. For example, for a series of reactions in an aqueous buffer system, flie rate of flie reaction would be a fimetion of the pH, but not of the concentration or identity of the acidic and basic components of the buffer. The kinetic expression for any such reaction will include a term for hydrogen-ion concentration, [H+]. The term general acid catalysis is used when the nature and concentration of proton donors present in solution affect the reaction rate. The kinetic expression for such a reaction will include a term for each of the potential proton donors that acts as a catalyst. The terms specific base catalysis and general base catalysis apply in the same way to base-catalyzed reactions. 

Scheme (b) includes reactions formerly described by a variety of names, such as dissociation, neutralisation, hydrolysis and buffer action (see below). One acid-base pair may involve the solvent (in water H30+ —H2OorH20 — OH ), showing that ions such as HsO+ and OH- are in principle only particular examples of an extended class of acids and bases though, of course, they do occupy a particularly important place in practice. It follows that the properties of an acid or base may be greatly influenced by the nature of the solvent employed. 

This shows that the pM value of the solution is fixed by the value of K and the ratio of complex-ion concentration to that of the free ligand. If more of M is added to the solution, more complex will be formed and the value of pM will not change appreciably. Likewise, if M is removed from the solution by some reaction, some of the complex will dissociate to restore the value of pM. This recalls the behaviour of buffer solutions encountered with acids and bases (Section 2.20), and by analogy, the complex-ligand system may be termed a metal ion buffer. 

Ionic strength adjuster buffer 565, 570 Ionisation constants of indicators, 262, (T) 265 of acids and bases, (T) 832, 833, 834 see also Dissociation constants Ionisation suppressant 793 Iron(II), D. of by cerium(IV) ion, (cm) 546 by cerium(IV) sulphate, (ti) 382 by potassium dichromate, (ti) 376 by potassium permanganate, (ti) 368 see also under Iron... 

Examination of the effect of pH on the rates of protodeboronation of the 2,6-dimethoxy compound at 90 °C in malonic acid-sodium malonate buffer solutions of ionic strength 0.14 gave the data in Table 199. A plot of these data revealed the curve shown in Fig. 3 (one of the points was misplotted on the original) and the linear portions of the plot were attributed to acid and base catalysis as shown on Fig. 3, and since the rates in the region of pH 4-5 are higher than would be... 

A catalytic system may contain active components other than H30+, H2O, and OH-. Weak acids and bases may also be efficient catalysts. These include, of course, both components of the buffer. Their contributions are in addition to the three terms seen before. If they are designated as BH+ and B, the rate constant is... 

Buffers are often prepared with equal molar concentrations of the conjugate acid and base. In these equimolar solutions,... 

This very simple result makes it easy to make an initial choice of a buffer we just select an acid that has a pfC, close to the pH that we require and prepare an equimolar solution with its conjugate base. When we prepare a buffer for pH > 7 we have to remember that the acid is supplied by the salt, that the conjugate base is the base itself, and that the pKa is that of the conjugate acid of the base (and hence related to the pKh of the base by pR l + pKh = p/conjugate acid and base have unequal concentrations—such as those considered in Examples 11.1 and 11.2—are buffers, but they may be less effective than those in which the concentrations are nearly equal (see Section 11.3). Table 11.1 lists some typical buffer systems. 

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 of a buffer solution is close to the pKa of the weak acid component when the acid and base have similar concentrations. 

Buffer capacity also depends on the relative concentrations of weak acid and base. Broadly speaking, a buffer is found experimentally to have a high capacity for acid when the amount of base present is at least 10% of the amount of acid. Otherwise, the base is used up quickly as strong acid is added. Similarly, a buffer has a high capacity for base when the amount of acid present is at least 10% of the amount of base, because otherwise the acid is used up quickly as strong base is added. 

This chapter describes several Important applications of aqueous equilibria. We begin with a discussion of buffer chemistry, followed by a description of acid and base titration reactions. Then we change our focus to examine the solubility equilibria of inorganic salts. The chapter concludes with a discussion of the equilibria of complex Ions. 

Human blood contains a variety of acids and bases that maintain the pH very close to 7.4 at all times. Close control of blood pH Is critical because death results if the pH of human blood drops below 7.0 orrises above 7.8. This narrow pH range corresponds to only a fivefold change in the concentration of hydronium ions. Chemical equilibria work in the blood to hold the pH within this narrow window. Close control of pH is achieved by a buffer solution, so called because it protects, or buffers, the solution against pH variations. 

The sum of the moles of the products is equal to the number of initial moles of acetate, and the final amounts of acid and base are comparable. Thus, the calculations support the conclusion that the final mixture is a buffer solution. 

The first four steps of the seven-step strategy are identical to the ones in Example. In this example, addition of a strong acid or base modifies the concentrations that go into the buffer equation. We need to determine the new concentrations (Step 5) and then apply the buffer equation (Step 6). In dealing with changes in amounts of acid and base, it is often convenient to work with moles rather than molarities. The units cancel in the concentration term of the buffer equation, so the ratio of concentrations can be... 

Ottiger, C. Wunderli-Allenspach, H., Partition behavior of acids and bases in a phosphatidylcholine liposome-buffer equilibrium dialysis system, Eur. J. Pharm. Sci. 5, 223-231 (1997). 

Log k appears to correlate with log P for standards between log P —0.5 to 5.0. One limitation of this method is that solutes must be electrically neutral at the pH of the buffer solution because electrophoretic mobility of the charged solute leads to migration times outside the range of Tm and TEof- Basic samples are therefore run at pH 10, and acidic samples at pH 3, thus ensuring that most weak acids and bases will be in their neutral form. This method has been used in a preclinical discovery environment with a throughput of 100 samples per week [24]. 

The linear buffer is created by mixing several weak acids and bases in varying concentrations such that the titration curve is linear in terms of pH, and the ionic... 

Buffers minimize changes in pH when strong acids and bases are added. 

A buffer s job is to keep the pH of a solution from changing very much when either strong acids or strong bases are added to it. To handle the addition of both acids and bases, a buffer contains an acid (to react with added base) and a base (to react with added acid). Take a glass of water (beaker sounds too chemical, and blood sounds too gruesome)... 

In the titration curves shown in Fig. 23-5, you start with the fully protonated form of the amino acid. Notice that at pH s that are not near the pKa of any functional group, the pH changes more when base is added. Also notice that there are multiple buffer regions (where the pH doesn t change rapidly when base is added) when there are multiple acid and base groups present. If the pAVs of two groups are close to each... 

The C02-bicarbonate buffer is a little different from buffers using the usual kind of acids and bases, but it is extremely important in maintaining the acid-base balance of the blood. The acid form of the bicarbonate buffer is actually a gas dissolved in water. Dissolved C02 is turned into an acid by hydration to give H2C03. Hydrated C02 is then much like a carboxylic acid. It gives up a proton to a base and makes bicarbonate, HCO 3. 

Unless standards are prepared in buffered media, positive or negative deviations may result from measurements made at 348 nm or 372 nm respectively Alternatively, measurements can be made at the isosbesticpoint, i.e. where the absorbance curves of each form intersect, and where absorbance is not a function of equilibrium concentrations but only of the overall concentration. Solutions of weak acids and bases should also be measured at their isosbestic points for the same reason. 

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). 

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