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Chapter 15: Problem 2

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?

Short Answer

Expert verified
A buffered solution resists changes in pH because it contains both a weak acid (HA) and its conjugate base (A-), which can react with either H+ or OH- ions in solution. When small amounts of acids or bases are added, the concentrations of HA and A- are only slightly affected and the pH remains relatively constant, as seen through the Henderson-Hasselbalch equation: pH = pKa + log\(\frac{[A-]}{[HA]}\). This ability of a buffered solution to maintain pH within a narrow range is crucial for many biological processes and chemical reactions.

Step by step solution

01

Understanding the Components of a Buffered Solution

A buffered solution is composed of a weak acid (HA) and its conjugate base (A-), usually in the form of a salt such as NaA. In a buffered solution, both the weak acid and its conjugate base are present in appreciable amounts. The weak acid can react with added bases by donating a proton, while the conjugate base can react with added acids by accepting a proton. This ability for both components to participate in acid-base reactions helps to maintain the pH of the solution within a narrow range.
02

Role of Weak Acid and Conjugate Base When Strong Base is Added

If a strong base such as NaOH is added to the buffered solution, the OH- from the strong base will react with the weak acid (HA) in the solution, forming A- and water: HA + OH- → A- + H2O As a result, the concentration of HA in the solution decreases, and the concentration of A- increases.
03

Role of Weak Acid and Conjugate Base When Strong Acid is Added

If a strong acid such as HCl is added to the buffered solution, the H+ from the strong acid will react with the conjugate base (A-) in the solution, forming HA: A- + H+ → HA As a result, the concentration of A- in the solution decreases, and the concentration of HA increases.
04

Relating pH Changes to the Concentration of Weak Acid and Conjugate Base

To understand how the pH changes in the buffered solution, we can use the Henderson-Hasselbalch equation, which relates the pH of the solution to the concentrations of the weak acid and its conjugate base: pH = pKa + log\(\frac{[A-]}{[HA]}\) Since the concentration of HA and A- doesn't change drastically when small amounts of strong acids or bases are added to the buffered solution, the pH remains relatively constant. However, if a significant amount of acid or base is added, the buffer's capacity will be exceeded, and the pH will change to a greater extent.
05

Explaining Buffering to the Friend

A buffered solution resists changes in pH when acids or bases are added because it contains both a weak acid (HA) and its conjugate base (A-), which are capable of reacting with either H+ or OH- ions in solution. When small amounts of acids or bases are added to the buffered solution, the pH doesn't change significantly because the concentrations of HA and A- are only slightly affected, as observed through the Henderson-Hasselbalch equation. This demonstrates the ability of a buffered solution to maintain pH within a narrow range, which is essential for many biological processes and chemical reactions.

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Most popular questions from this chapter

What volume of \(0.0100 M\) NaOH must be added to \(1.00 \mathrm{~L}\) of \(0.0500 \mathrm{M} \mathrm{HOCl}\) to achieve a \(\mathrm{pH}\) of \(8.00 ?\)

A \(0.210-\mathrm{g}\) sample of an acid (molar mass \(=192 \mathrm{~g} / \mathrm{mol}\) ) is titrated with \(30.5 \mathrm{~mL}\) of \(0.108 \mathrm{M} \mathrm{NaOH}\) to a phenolphthalein end point. Is the acid monoprotic, diprotic, or triprotic?

Calculate the \(\mathrm{pH}\) of each of the following buffered solutions. a. \(0.10 M\) acetic acid \(/ 0.25 M\) sodium acetate b. \(0.25 M\) acetic acid/0.10 \(M\) sodium acetate c. \(0.080 M\) acetic acid \(/ 0.20 M\) sodium acetate d. \(0.20 M\) acetic acid/0.080 \(M\) sodium acetate

One method for determining the purity of aspirin \(\left(\mathrm{C}_{9} \mathrm{H}_{8} \mathrm{O}_{4}\right)\) is to hydrolyze it with \(\mathrm{NaOH}\) solution and then to titrate the remaining \(\mathrm{NaOH}\). The reaction of aspirin with \(\mathrm{NaOH}\) is as follows: A sample of aspirin with a mass of \(1.427 \mathrm{~g}\) was boiled in \(50.00 \mathrm{~mL}\) of \(0.500 \mathrm{M} \mathrm{NaOH}\). After the solution was cooled, it took \(31.92 \mathrm{~mL}\) of \(0.289 M \mathrm{HCl}\) to titrate the excess \(\mathrm{NaOH}\). Calculate the purity of the aspirin. What indicator should be used for this titration? Why?

Malonic acid \(\left(\mathrm{HO}_{2} \mathrm{CCH}_{2} \mathrm{CO}_{2} \mathrm{H}\right)\) is a diprotic acid. In the titration of malonic acid with \(\mathrm{NaOH}\), stoichiometric points occur at \(\mathrm{pH}=3.9\) and \(8.8\). A \(25.00-\mathrm{mL}\) sample of malonic acid of unknown concentration is titrated with \(0.0984 \mathrm{M} \mathrm{NaOH}\), requiring \(31.50 \mathrm{~mL}\) of the \(\mathrm{NaOH}\) solution to reach the phenolphthalein end point. Calculate the concentration of the initial malonic acid solution. (See Exercise 113.)
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