Acid-Base Equilibria, Buffers, and Salt Solutions Study Guide

Q5. Determine the pH of a 0.22 M NaF solution at 25°C. The Ka of HF is 3.5 × 10-5. (Notice this is a salt)

Background

Topic: Salt Solutions and Hydrolysis

This question tests your understanding of how salts derived from weak acids and strong bases affect pH. NaF is the salt of Na+ (from NaOH, a strong base) and F- (from HF, a weak acid). The F- ion hydrolyzes in water, making the solution basic.

Key Terms and Formulas

• Hydrolysis: Reaction of an ion with water to produce either H+ or OH-.

• Base ionization constant ($K_b$): $K_b = \frac{K_w}{K_a}$, where $K_w = 1.0 \times 10^{-14}$ at 25°C.

• For F-: $F^- + H_2O \rightleftharpoons HF + OH^-$

• pH calculation: $pH = 14 - pOH$

Step-by-Step Guidance

1. Calculate $K_b$ for F- using $K_b = \frac{K_w}{K_a}$, where $K_a$ is for HF.

2. Set up the equilibrium expression for F- hydrolysis: $K_b = \frac{[HF][OH^-]}{[F^-]}$.

3. Let $x$ be the concentration of OH- produced. Write the equilibrium concentrations: $[F^-] = 0.22 - x$, $[OH^-] = x$, $[HF] = x$.

4. Plug these values into the $K_b$ expression and solve for $x$ (but do not finish the calculation).

Try solving on your own before revealing the answer!

Q6. Calculate the pH of a buffer that is 0.225 M HC2H3O2 and 0.162 M NaC2H3O2. The Ka for HC2H3O2 is 1.8 × 10-5.

Background

Topic: Buffer Solutions and Henderson-Hasselbalch Equation

This question tests your ability to calculate the pH of a buffer solution using the concentrations of a weak acid and its conjugate base.

Key Terms and Formulas

• Buffer: A solution that resists changes in pH when small amounts of acid or base are added.

• Henderson-Hasselbalch Equation: $pH = pK_a + \log\left(\frac{[A^-]}{[HA]}\right)$

• $pK_a = -\log(K_a)$

• $[HA]$ = concentration of weak acid (HC2H3O2), $[A^-]$ = concentration of conjugate base (NaC2H3O2)

Step-by-Step Guidance

1. Calculate $pK_a$ using $pK_a = -\log(K_a)$, where $K_a = 1.8 \times 10^{-5}$.

2. Identify $[HA] = 0.225$ M and $[A^-] = 0.162$ M.

3. Plug these values into the Henderson-Hasselbalch equation: $pH = pK_a + \log\left(\frac{0.162}{0.225}\right)$.

4. Calculate the logarithmic term, but do not finish the calculation.

Try solving on your own before revealing the answer!

Q7. Consider a buffer composed of an acid HA and its conjugate base A-. Which solution would make the most effective buffer?

Background

Topic: Buffer Capacity and Effectiveness

This question tests your understanding of what makes a buffer most effective. The effectiveness depends on the concentrations and the ratio of acid to conjugate base.

Key Terms and Formulas

• Buffer capacity: The amount of acid or base a buffer can neutralize before the pH changes significantly.

• Effective buffer: Most effective when $[HA] \approx [A^-]$ and both concentrations are high.

Step-by-Step Guidance

1. Compare the ratios of $[HA]$ to $[A^-]$ in each option. The most effective buffer has a ratio close to 1:1.

2. Consider the absolute concentrations: higher concentrations provide greater buffer capacity.

3. Identify which option has both a 1:1 ratio and the highest concentrations.

Try solving on your own before revealing the answer!

Final Answers

Q5: 10.20

Q6: 4.60

Q7: B) 0.500 M HA; 0.500 M A-

For Q5, the solution is basic due to F- hydrolysis. For Q6, the Henderson-Hasselbalch equation gives a pH close to the pKa. For Q7, the most effective buffer is the one with equal and highest concentrations of acid and conjugate base.