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Chapter 15: Problem 3

Mixing together solutions of acetic acid and sodium hydroxide can make a buffered solution. Explain. How does the amount of each solution added change the effectiveness of the buffer?

Short Answer

Expert verified
Mixing acetic acid (CH3COOH) and sodium hydroxide (NaOH) forms a buffered solution containing acetic acid and its conjugate base, acetate ions (CH3COO-). A buffer solution resists changes in pH when small amounts of acid or base are added. The effectiveness of a buffer solution depends on the concentrations of the weak acid and its conjugate base. The Henderson-Hasselbalch equation illustrates this relationship: \( pH = pK_a + log_{10} \frac{[CH_3COO^-]}{[CH_3COOH]} \). By adjusting the amount of acetic acid and sodium hydroxide added, we can control the pH of the resulting buffered solution. If the concentrations of acetic acid and acetate ions are nearly equal, the buffer solution will have a maximum buffering capacity, making it more effective at resisting changes in pH.

Step by step solution

01

Understand the concept of a buffer solution

A buffer solution is a solution that resists changes in pH when small amounts of an acid or a base are added to it. It consists of a mixture of a weak acid and its conjugate base, or a weak base and its conjugate acid.
02

Understand the reaction between acetic acid and sodium hydroxide

When acetic acid (CH3COOH), a weak acid, is mixed with sodium hydroxide (NaOH), a strong base, a neutralization reaction occurs, producing sodium acetate (CH3COONa) and water (H2O). The balanced equation for this reaction is: \( CH_3COOH + NaOH \rightarrow CH_3COONa + H_2O \) Sodium acetate, formed during this reaction, dissociates in water to form acetate ions (CH3COO-) and sodium ions (Na+). The acetate ion is the conjugate base of acetic acid. The acidic component of the buffer system is acetic acid, and the basic component is the acetate ion.
03

Formation of the buffer solution

When acetic acid and sodium hydroxide are mixed, the resulting solution contains both acetic acid and its conjugate base, acetate ions: \( CH_3COOH \leftrightarrow H^+ + CH_3COO^- \) Since both acidic and basic components are present in the solution, it forms a buffer system that resists changes in pH when small amounts of acid or base are added.
04

Effect of the amount of each solution on buffer effectiveness

The effectiveness of a buffer solution depends on the concentrations of the weak acid and its conjugate base. The Henderson-Hasselbalch equation describes the relationship between the pH of a buffer solution and the concentrations of the acidic and basic components: \( pH = pK_a + log_{10} \frac{[CH_3COO^-]}{[CH_3COOH]} \) According to this equation, the pH of the buffer solution depends on the ratio of the concentrations of acetate ions (CH3COO-) and acetic acid (CH3COOH). By adjusting the amount of acetic acid and sodium hydroxide added to the solution, we can control the ratio of these two components and, consequently, the pH of the resulting buffered solution. If the concentrations of acetic acid and acetate ions are nearly equal, the buffer solution will be most effective at resisting changes in pH, as it will have a maximum buffering capacity. If the concentration of one component is much higher than the other, the buffering capacity will be decreased, making the buffer solution less effective at resisting changes in pH.

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Most popular questions from this chapter

Sketch the titration curve for the titration of a generic weak base B with a strong acid. The titration reaction is $$ \mathrm{B}+\mathrm{H}^{+} \rightleftharpoons \mathrm{BH}^{+} $$ On this curve, indicate the points that correspond to the following: a. the stoichiometric (equivalence) point b. the region with maximum buffering c. \(\mathrm{pH}=\mathrm{p} K_{\mathrm{a}}\)

Lactic acid is a common by-product of cellular respiration and is often said to cause the "burn" associated with strenuous activity. A \(25.0-\mathrm{mL}\) sample of \(0.100 M\) lactic acid \(\left(\mathrm{HC}_{3} \mathrm{H}_{5} \mathrm{O}_{3}, \mathrm{p} K_{\mathrm{a}}=\right.\) \(3.86\) ) is titrated with \(0.100 M\) NaOH solution. Calculate the \(\mathrm{pH}\) after the addition of \(0.0 \mathrm{~mL}, 4.0 \mathrm{~mL}, 8.0 \mathrm{~mL}, 12.5 \mathrm{~mL}, 20.0 \mathrm{~mL}\), \(24.0 \mathrm{~mL}, 24.5 \mathrm{~mL}, 24.9 \mathrm{~mL}, 25.0 \mathrm{~mL}, 25.1 \mathrm{~mL}, 26.0 \mathrm{~mL}, 28.0 \mathrm{~mL}\) and \(30.0 \mathrm{~mL}\) of the \(\mathrm{NaOH}\). Plot the results of your calculations as pH versus milliliters of \(\mathrm{NaOH}\) added.

Calculate the \(\mathrm{pH}\) at the halfway point and at the equivalence point for each of the following titrations. a. \(100.0 \mathrm{~mL}\) of \(0.10 \mathrm{M} \mathrm{HC}_{7} \mathrm{H}_{5} \mathrm{O}_{2}\left(K_{\mathrm{a}}=6.4 \times 10^{-5}\right)\) titrated by \(0.10 \mathrm{M} \mathrm{NaOH}\) b. \(100.0 \mathrm{~mL}\) of \(0.10 \mathrm{M} \mathrm{C}_{2} \mathrm{H}_{3} \mathrm{NH}_{2}\left(K_{\mathrm{b}}=5.6 \times 10^{-4}\right)\) titrated by \(0.20 \mathrm{M} \mathrm{HNO}_{3}\) c. \(100.0 \mathrm{~mL}\) of \(0.50 \mathrm{M} \mathrm{HCl}\) titrated by \(0.25 \mathrm{M} \mathrm{NaOH}\)

A friend asks the following: "Consider a buffered solution made up of the weak acid HA and its salt NaA. If a strong base like \(\mathrm{NaOH}\) is added, the HA reacts with the \(\mathrm{OH}^{-}\) to form \(\mathrm{A}^{-}\). Thus the amount of acid (HA) is decreased, and the amount of base \(\left(\mathrm{A}^{-}\right)\) is increased. Analogously, adding \(\mathrm{HCl}\) to the buffered solution forms more of the acid (HA) by reacting with the base \(\left(\mathrm{A}^{-}\right) .\) Thus how can we claim that a buffered solution resists changes in the \(\mathrm{pH}\) of the solution?' How would you explain buffering to this friend?

What are the major species in solution after \(\mathrm{NaHSO}_{4}\) is dissolved in water? What happens to the \(\mathrm{pH}\) of the solution as more \(\mathrm{NaHSO}_{4}\) is added? Why? Would the results vary if baking soda \(\left(\mathrm{NaHCO}_{3}\right)\) were used instead?
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