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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

What quantity (moles) of \(\mathrm{NaOH}\) must be added to \(1.0 \mathrm{~L}\) of \(2.0 \mathrm{M} \mathrm{HC}_{2} \mathrm{H}_{3} \mathrm{O}_{2}\) to produce a solution buffered at each \(\mathrm{pH}\) ? a. \(\mathrm{pH}=\mathrm{p} K_{\mathrm{a}}\) b. \(\mathrm{pH}=4.00\) c. \(\mathrm{pH}=5.00\)

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

One method for determining the purity of aspirin \(\left(\mathrm{C}_{9} \mathrm{H}_{8} \mathrm{O}_{4}\right)\) is to hydrolyze it with NaOH solution and then to titrate the remaining \(\mathrm{NaOH}\). The reaction of aspirin with \(\mathrm{NaOH}\) is as follows: \(\mathrm{C}_{?} \mathrm{H}_{3} \mathrm{O}_{4}(s)+2 \mathrm{OH}^{-}(a q)\) Aspirin $$ \begin{array}{c} \text { Bail } \mathrm{C}_{7} \mathrm{H}_{5} \mathrm{O}_{3}^{-}(a q)+\mathrm{C}_{2} \mathrm{H}_{3} \mathrm{O}_{2}^{-}(a q)+\mathrm{H}_{2} \mathrm{O}(l) \\ \text { Salicylate ion Acetate ion } \end{array} $$ A sample of aspirin with a mass of \(1.427 \mathrm{~g}\) was boiled in \(50.00 \mathrm{~mL}\) of \(0.500 M \mathrm{NaOH}\). After the solution was cooled, it took \(31.92 \mathrm{~mL}\) of \(0.289 \mathrm{M} \mathrm{HCl}\) to titrate the excess \(\mathrm{NaOH}\). Calculate the purity of the aspirin. What indicator should be used for this titration? Why?

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.

A certain indicator HIn has a \(\mathrm{p} K_{2}\) of \(3.00\) and a color change becomes visible when \(7.00 \%\) of the indicator has been converted to \(\mathrm{In}^{-}\). At what \(\mathrm{pH}\) is this color change visible?
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