The two Equations are both forms of the Henderson-Hasselbalch equationA rearranged version of the equilibrium constant expression that provides a direct way to calculate the pH of a buffer pK solution: pH = + log([base]/[acid])., named after the two early-20th-century chemists who first a noticed that this rearranged version of the equilibrium constant expression provides an easy way to calculate the pH of a buffer solution. In general, the validity of the Henderson-Hasselbalch equation may be limited to solutions whose concentrations are at least 100 times greater than their Ka values. There are three special cases where the Henderson-Hasselbalch equation is easily interpreted without the need for calculations: [base] = [acid]. these conditions, [base]/[acid] = 1 in the Equation above. Because log 1 = 0, pH = pKa, regardless of the actual concentrations of the acid and base. Recall that this corresponds to the midpoint in the titration of a weak acid or a weak base. [base]/[acid] = 10. In the Equation above, because log 10 = 1, pH = pKa + 1. [base]/[acid] = 100. In the Equation above, because log 100 = 2, pH = pKa + 2. Each time we increase the [base]/[acid] ratio by 10, the pH of the solution increases by 1 pH unit. Conversely, if the [base]/[acid] ratio is 0.1, then pH = pKa 1. Each additional factor-of-10 decrease − in the [base]/[acid] ratio causes the pH to decrease by 1 pH unit. Note the Pattern If [base] = [acid] for a buffer, then pH = pKa. Changing this ratio by a factor of 10 either way changes the pH by ±1 unit. The Henderson-Hasselbalch equation can also be used to calculate the pH of a buffer solution adding a given amount of strong acid or strong base. Note the Pattern Only the amounts (in or millimoles) of the acidic and basic components of the buffer are needed to use the Henderson-Hasselbalch equation, not their concentrations.
The two Equations are both forms of the Henderson-Hasselbalch equationA rearranged version of the equilibrium constant expression that provides a direct way to calculate the pH of a buffer solution: pH = pK a + log([base]/[acid])., named after the two early-20th-century chemists who first noticed that this rearranged version of the equilibrium constant expression provides an easy way to calculate the pH of a buffer solution. In general, the validity of the Henderson-Hasselbalch equation may be limited to solutions whose concentrations are at least 100 times greater than their Ka values. There are three special cases where the Henderson-Hasselbalch equation is easily interpreted without the need for calculations: [base] = [acid]. these conditions, [base]/[acid] = 1 in the Equation above. Because log 1 = 0, pH = pKa, regardless of the actual concentrations of the acid and base. Recall that this corresponds to the midpoint in the titration of a weak acid or a weak base. [base]/[acid] = 10. In the Equation above, because log 10 = 1, pH = pKa + 1. [base]/[acid] = 100. In the Equation above, because log 100 = 2, pH = pKa + 2. Each time we increase the [base]/[acid] ratio by 10, the pH of the solution increases by 1 pH unit. Conversely, if the [base]/[acid] ratio is 0.1, then pH = pKa − 1. Each additional factor-of-10 decrease in the [base]/[acid] ratio causes the pH to decrease by 1 pH unit. Note the Pattern If [base] = [acid] for a buffer, then pH = pKa. Changing this ratio by a factor of 10 either way changes the pH by ±1 unit. The Henderson-Hasselbalch equation can also be used to calculate the pH of a buffer solution adding a given amount of strong acid or strong base. Note the Pattern Only the amounts (in or millimoles) of the acidic and basic components of the buffer are needed to use the Henderson-Hasselbalch equation, not their concentrations.
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