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Chapter 15: Problem 59

What characteristic properties do buffered solutions possess?

Short Answer

Expert verified
Buffered solutions possess characteristic properties such as the presence of a weak acid and its conjugate base or a weak base and its conjugate acid, limited buffer capacity, and working within a specific pH range. They can maintain a relatively constant pH despite the addition of small amounts of acid or base by neutralizing these added compounds. The Henderson-Hasselbalch equation can be used to calculate the pH of a buffered solution. Buffered solutions have various applications in biological, chemical, and industrial processes that require maintaining stable pH levels.

Step by step solution

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1. Key Components of Buffered Solutions

A buffered solution is made up of a weak acid and its conjugate base, or a weak base and its conjugate acid. The weak acid or base partially ionizes in the solution, while its conjugate counterpart neutralizes small amounts of added acids or bases.
02

2. Buffer Capacity

Buffered solutions have limited a buffer capacity, which is the ability of a buffer to resist changes in pH. The buffer capacity depends on the concentration of the weak acid/base and its conjugate in the solution. Buffer solutions can only withstand a certain amount of added acid or base before the pH begins to change significantly.
03

3. pH Range

Buffered solutions can maintain a relatively constant pH within a specific pH range. This pH range is determined by the pKa of the weak acid or pKb of the weak base. The Henderson-Hasselbalch equation can help us calculate the pH of a buffered solution: For weak acids: \[pH = pKa + \log{\frac{[\text{conjugate base}]}{[\text{weak acid}]}}\] For weak bases: \[pOH = pKb + \log{\frac{[\text{conjugate acid}]}{[\text{weak base}]}}\] Where pKa and pKb are the negative logarithms of the acid and base dissociation constants, respectively.
04

4. Resistance to pH Change

Buffered solutions minimize pH changes by neutralizing added acids or bases. When a small amount of acid is added to a buffer, the conjugate base present in the buffer neutralizes the added H+ ions. In contrast, when a small amount of base is added, the weak acid present in the buffer donates H+ ions to neutralize the added OH- ions. This process helps maintain the pH of the buffered solution within its preferred range.
05

5. Common Applications

Buffered solutions are essential in many biological and chemical systems where maintaining a stable pH is critical. Some common applications include maintaining pH levels in blood, aquatic habitats, and in various industrial and laboratory processes.

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Most popular questions from this chapter

For each pair of concentrations, tell which represents the more basic solution. a. \(\left[\mathrm{H}^{+}\right]=0.000013 \mathrm{M}\) or \(\left[\mathrm{OH}^{-}\right]=0.0000032 \mathrm{M}\) b. \(\left[\mathrm{H}^{+}\right]=1.03 \times 10^{-6} \mathrm{M}\) or \(\left[\mathrm{OH}^{-}\right]=1.54 \times 10^{-8} \mathrm{M}\) c. \(\left[\mathrm{OH}^{-}\right]=4.02 \times 10^{-7} \mathrm{M}\) or \(\left[\mathrm{OH}^{-}\right]=0.0000001 \mathrm{M}\)

Anions containing hydrogen (for example, \(\mathrm{HCO}_{3}^{-}\) and \(\mathrm{H}_{2} \mathrm{PO}_{4}^{2-}\) ) show amphoteric behavior when reacting with other acids or bases. Write equations illustrating the amphoterism of these anions.

For a hydrogen ion concentration of \(2.33 \times 10^{-6} \mathrm{M}\) how many decimal places should we give when expressing the pH of the solution?

Calculate the hydrogen ion concentration, in moles per liter, for solutions with each of the following \(\mathrm{pH}\) or \(\mathrm{pOH}\) values. a. \(\mathrm{pOH}=0.90\) b. \(\mathrm{pH}=0.90\) c. \(\mathrm{pOH}=10.3\) d. \(\mathrm{pH}=5.33\)

Which of the following acids are classified as strong acids? a. \(\mathrm{HNO}_{3}\) b. \(\mathrm{CH}_{3} \mathrm{COOH}\left(\mathrm{HC}_{2} \mathrm{H}_{3} \mathrm{O}_{2}\right)\) c.\( HCl\) d. \(HF\) e. \(\mathrm{HClO}_{4}\)
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