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Chapter 15: Problem 13

A best buffer has about equal quantities of weak acid and conjugate base present as well as having a large concentration of each species present. Explain.

Short Answer

Expert verified
A best buffer has approximately equal quantities of weak acid and conjugate base, as well as large concentrations of each species, to maximize buffer capacity and effectively resist significant changes in pH when acid or base is added. This is explained by the Henderson-Hasselbalch equation, which shows that the pH of a buffer solution is optimal when the concentrations of weak acid and conjugate base are nearly equal, allowing successful neutralization of both added acid and base. Additionally, higher concentrations of these components enhance the buffer's resistance to pH changes, ultimately maximizing its capacity.

Step by step solution

01

Definition of a Buffer Solution

A buffer solution is a solution that can resist significant changes in pH when small amounts of acid or base are added. It consists of a weak acid and its conjugate base or a weak base and its conjugate acid.
02

The Henderson-Hasselbalch Equation

The Henderson-Hasselbalch equation relates the pH of a buffer solution to the concentrations of the weak acid and conjugate base in the solution. The equation is: \(pH = pKa + log\frac{[A^-]}{[HA]}\) Where: - pH is the acidity level of the solution - pKa is the negative logarithm of the acid dissociation constant (Ka) of the weak acid - [A^-] is the concentration of the conjugate base in the solution - [HA] is the concentration of the weak acid in the solution
03

Buffer Capacity and the Henderson-Hasselbalch Equation

Buffer capacity measures a buffer solution's ability to resist changes in pH when an acid or base is added. To maximize buffer capacity, the pH of the solution should be close to the pKa of the weak acid, so that the concentrations of the weak acid and conjugate base are approximately equal ([A^-] 鈮� [HA]). This ensures that the solution can effectively neutralize both added acid and base.
04

Large Concentrations of Weak Acid and Conjugate Base

The effectiveness of a buffer solution is also heavily dependent on the concentrations of its components. When a buffer solution has large concentrations of both weak acid and its conjugate base, it can absorb more acid or base before its pH changes significantly. Higher concentrations of these components make the buffer more resistant to pH changes, ultimately maximizing the buffer capacity.
05

Conclusion

In summary, a best buffer has approximately equal quantities of weak acid and conjugate base, as well as large concentrations of each species, because it maximizes buffer capacity and effectively resists significant changes in pH when acid or base is added. This can be explained by analyzing the Henderson-Hasselbalch equation and understanding the concept of buffer capacity.

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Most popular questions from this chapter

Repeat the procedure in Exercise \(67,\) but for the titration of 25.0 \(\mathrm{mL}\) of 0.100\(M\) propanoic acid \(\left(\mathrm{HC}_{3} \mathrm{H}_{3} \mathrm{O}_{2}, K_{\mathrm{a}}=1.3 \times 10^{-5}\right)\) with 0.100 \(\mathrm{M} \mathrm{NaOH}\) .

Consider the titration of 100.0 \(\mathrm{mL}\) of 0.100 \(\mathrm{M} \mathrm{HCN}\) by 0.100 \(\mathrm{M} \mathrm{KOH}\) at \(25^{\circ} \mathrm{C} .\left(K_{\mathrm{a}} \text { for } \mathrm{HCN}=6.2 \times 10^{-10} .\right)\) a. Calculate the pH after 0.0 \(\mathrm{mL}\) of KOH has been added. b. Calculate the pH after 50.0 \(\mathrm{mL}\) of KOH has been added. c. Calculate the pH after 75.0 \(\mathrm{mL}\) of KOH has been added. d. Calculate the pH at the equivalence point. e. Calculate the pH after 125 \(\mathrm{mL}\) of KOH has been added.

Th pH of blood is steady at a value of approximately 7.4 as a result of the following equilibrium reactions: $$ \mathrm{CO}_{2}(a q)+\mathrm{H}_{2} \mathrm{O}(l) \leftrightharpoons \mathrm{H}_{2} \mathrm{CO}_{3}(a q) \leftrightharpoons \mathrm{HCO}_{3}-(a q)+\mathrm{H}^{+}(a q) $$ The actual buffer system in blood is made up of \(\mathrm{H}_{2} \mathrm{CO}_{3}\) and \(\mathrm{HCO}_{3}\) - One way the body keeps the pH of blood at 7.4 is by regulating breathing. Under what blood ph conditions will the body increase breathing and under what blood pH conditions will the body decrease breathing? Explain.

Calculate the pH of a solution that is 0.60\(M\) HF and 1.00\(M \mathrm{KF}\)

Carbonate buffers are important in regulating the pH of blood at \(7.40 .\) If the carbonic acid concentration in a sample of blood is 0.0012 M, determine the bicarbonate ion concentration required to buffer the pH of blood at pH \(=7.40\) $$ \mathrm{H}_{2} \mathrm{CO}_{3}(a q) \rightleftharpoons \mathrm{HCO}_{3}^{-}(a q)+\mathrm{H}^{+}(a q) \quad K_{\mathrm{a}}=4.3 \times 10^{-7} $$
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