Weak acids Henderson-Hasselbalch equation

Multiprotic weak acids can be used to prepare buffers at as many different pH s as there are acidic protons. For example, a diprotic weak acid can be used to prepare buffers at two pH s and a triprotic weak acid can be used to prepare three different buffers. The Henderson-Hasselbalch equation applies in each case. Thus, buffers of malonic acid (pKai = 2.85 and = 5.70) can be prepared for which... 

Any solution containing comparable amounts of a weak acid, HA, and its conjugate weak base, A-, is a buffer. As we learned in Chapter 6, we can calculate the pH of a buffer using the Henderson-Hasselbalch equation. 

This relationship is known as the Henderson-Hasselbalch equation. Thus, the pH of a solution can be calculated, provided and the concentrations of the weak acid HA and its conjugate base A are known. Note particularly that when [HA] = [A ], pH = pAl,. For example, if equal volumes of 0.1 MHAc and 0.1 M sodium acetate are mixed, then... 

The Henderson-Hasselbalch equation provides a general solution to the quantitative treatment of acid-base equilibria in biological systems. Table 2.4 gives the acid dissociation constants and values for some weak electrolytes of biochemical interest. 

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 ... 

A note on good practice Keep in mind the approximations required for the use of the Henderson-Hasselbalch equation (that the concentrations of both the weak acid and its conjugate base are much greater than the hydronium ion concentration). Because the equation uses molar concentration instead of activities, it also ignores the interactions between ions. 

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 Henderson-Hasselbalch Equation Describes the Behavior of Weak Acids Buffers... 

The Henderson-Hasselbalch equation is derived below. A weak acid, HA, ionizes as follows ... 

The case of a buffer consisting of a weak base and its acidic form (for example, NH3 and NH4) is treated in an analogous way. Equations of the type of (1.4.26) are sometimes called the Henderson-Hasselbalch equations. 

Words that can be used as topics in essays 5% rale buffer common ion effect equilibrium expression equivalence point Henderson-Hasselbalch equation heterogeneous equilibria homogeneous equilibria indicator ion product, P Ka Kb Kc Keq KP Ksp Kw law of mass action Le Chatelier s principle limiting reactant method of successive approximation net ionic equation percent dissociation pH P Ka P Kb pOH reaction quotient, Q reciprocal rule rule of multiple equilibria solubility spectator ions strong acid strong base van t Hoff equation weak acid weak base... 

Consider how a weak electrolyte is distributed across the gastric mucosa between plasma (pH 7.4) and gastric fluid (pH 1.0). In each compartment, the Henderson-Hasselbalch equation gives the ratio of acid-base concentrations. The negative logarithm of the acid dissociation constant is designated here by the symbol pAa rather than the more precisely correct pK1. 

Buffers are solutions that resist a change in pH when we add an acid or base. A buffer contains both a weak acid (HA) and its conjugate base (A-). The acid part will neutralize any base added and the base part of the buffer will neutralize any acid added to the solution. We may calculate the hydronium ion concentration of a buffer by rearranging the Ka expression to yield the Henderson-Hasselbalch equation, which we can use to calculate the pH of a buffer ... 

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. 

Thus, for a weak acid with a given Ka (or pKJ and a given ratio of conjugate base concentration to acid concentration, the pH may be calculated. Or, given the desired pH and IQ (pKa), the ratio of salt concentration to acid concentration can be calculated and the buffer subsequently prepared. Equations (5.26) to (5.30) are each a form of the Henderson-Hasselbalch equation for dealing with buffer solutions. 

The second pair of equations relate pH to pATa for weak acids and weak bases, though they are actually variants of the Henderson-Hasselbalch equation ... 

One should always be careful when using the Henderson-Hasselbalch equation pH = pK + log ([A ]/[HA]). The value of [HA] is not always moles of the weak acid per unit volume, cha, that had been added to the moles of the conjugate base per unit volume, Ca-. There are five species for this acid/base equilibria HA, A, H+, OH , and Na+ (or, whatever cation portion of the salt of the conjugate base is present). These species are all interrelated by five expressions ... 

B. The Henderson-Hasselbalch equation is derived from the rearrangement of the equilibrium equation for dissociation of a weak acid. 

Figure I-I. Weak acids act as buffers in a pH range near their pK s. According to the Henderson-Hasselbalch equation, when the ratio of conjugate base to conjugate acid, [A ]/[HA] is plotted versus pH, a titration curve is generated that indicates a region of good buffering at pH = pK I pH unit.

The answer is D. Weak acids like lactic acid never completely dissociate in solution and are thus defined by the property that at least some of the protonated (undissociated acid) form and the unprotonated (conjugate base) form of the acid are present at all concentrations and pH conditions. The indicated of 5.2 is consistent with the idea that the lactate anion retains a strong affinity for protons, a hallmark of a weak acid. The lactate anion is highly water-soluble. All weak acids obey the Henderson-Hasselbalch equation. 

IONIZATION OF WEAK ACIDS AND WEAK BASES THE HENDERSON-HASSELBALCH EQUATION... 

Note that the protonated form of a weak acid is the neutral, more lipid-soluble form, whereas the unprotonated form of a weak base is the neutral form. The law of mass action requires that these reactions move to the left in an acid environment (low pH, excess protons available) and to the right in an alkaline environment. The Henderson-Hasselbalch equation relates the ratio of protonated to unprotonated weak acid or weak base to the molecule s pKa and the pH of the medium as follows ... 

The addition of H+ to this solution favours the back reaction while the addition of base favours the forward reaction. The weak acid/salt pair thus acts to minimize ApH. An analogous situation exists for buffers consisting of a weak base and its salt. The pH of a buffer can be calculated from the concentration of its components by the Henderson-Hasselbalch equation... 

Hasselbalch equation, which is important for understanding buffer action and acid-base balance in the blood and tissues of vertebrates. This equation is simply a useful way of restating the expression for the dissociation constant of an acid. For the dissociation of a weak acid HA into H+ and A-, the Henderson-Hasselbalch equation can be derived as follows ... 

Titration curve of alanine By applying the Henderson-Hasselbalch equation to each dissociable acidic group, it is possi ble to calculate the complete titration curve of a weak acid. Figure 1.11 shows the change in pH that occurs during the addition of base to the fully protonated form of alanine (I) to produce the completely deprotonated form (III). Note the following ... 

The Henderson-Hasselbalch equation can be used to calculate the quantitative relationship between the concentration of a weak acid and its conjugate base. 

The Henderson-Hasselbalch equation tells us the pH of a solution, provided we know the ratio of the concentrations of conjugate acid and base, as well as p/f, for the acid. If a solution is prepared from the weak base B and its conjugate acid, the analogous equation is... 

The pH is calculated from the Henderson-Hasselbalch equation for the weak acid, BH+,... 

The most effective buffering system contains equal concentrations of the acid, HA, and its conjugate base, A-. According to the Henderson-Hasselbalch equation (2.6), when [A-] is equal to [HA], pH equals pKa. Therefore, the pKa of a weak acid-base system represents the center of the buffering region. The effective range of a buffer system is generally two pH units, centered at the pKa value (Equation 2.9). 

The Henderson-Hasselbalch equation says that the pH of a buffer solution has a value close to the pKa of the weak acid, differing only by the amount log [base]/[add]. When [base]/[acid] = 1, then log [base]/[acid] = 0, and the pH equals the pKa. 

The real importance of the Henderson-Hasselbalch equation, particularly in biochemistry, is that it tells us how the pH affects the percent dissociation of a weak acid. Suppose, for example, that you have a solution containing the amino acid glycine, one of the molecules from which proteins are made, and that the pH of the solution is 2.00 pH units greater than the pKa of glycine ... 

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