Q: Which equation is most commonly used to calculate the pH of a buffer solution when the concentrations of the acid and its conjugate base are known?

$\mathrm{pH} = \mathrm{p}K_a + \log \left( \frac{[\text{conjugate base}]}{[\text{acid}]} \right)$

Question 1

Which of the following best describes the mechanism by which a buffer solution resists changes in pH when a small amount of strong acid is added?

Accepted Answer

The conjugate base in the buffer reacts with the added acid, converting it to the weak acid and minimizing pH change.

Question 2

A buffer is prepared by mixing 0.30 mol of acetic acid ($\mathrm{CH_3COOH}$, $K_a = 1.8 	imes 10^{-5}$) and 0.30 mol of sodium acetate in enough water to make 1.0 L of solution. What is the pH of this buffer?

Accepted Answer

pH = $4.74$

Question 3

Which equation is most commonly used to calculate the pH of a buffer solution when the concentrations of the acid and its conjugate base are known?

Accepted Answer

$\mathrm{pH} = \mathrm{p}K_a + \log \left( \frac{[	ext{conjugate base}]}{[	ext{acid}]} ight)$

Question 4

A buffer is made by mixing 0.12 M lactic acid ($K_a = 1.4 	imes 10^{-4}$) and 0.10 M sodium lactate. What is the pH of the solution?

Accepted Answer

pH = $3.85$

Question 5

Which of the following statements about buffer capacity is correct?

Accepted Answer

Buffer capacity is highest when the concentrations of acid and conjugate base are both high and nearly equal.

Question 6

Suppose you dilute a buffer composed of 0.25 M acetic acid and 0.25 M acetate by a factor of 2. What happens to the pH?

Accepted Answer

The pH remains nearly unchanged.

Question 7

Calculate the number of grams of ammonium chloride ($\mathrm{NH_4Cl}$, molar mass $53.5\ \mathrm{g/mol}$) that must be added to 2.00 L of a 0.500 M ammonia solution to obtain a buffer of pH = 9.20. ($K_b$ for ammonia $= 1.80 	imes 10^{-5}$).

Accepted Answer

$68.2\ \mathrm{g}$

Question 8

Which of the following is NOT a method for preparing a buffer solution?

Accepted Answer

Mixing a strong acid with a strong base in equal amounts.

Question 9

A buffer solution is most effective at resisting changes in pH when:

Accepted Answer

The pH is equal to the pKa of the acid.

Question 10

If a small amount of strong base is added to a buffer composed of $\mathrm{NH_3}$ and $\mathrm{NH_4^+}$, what happens?

Accepted Answer

$\mathrm{NH_4^+}$ reacts with the base to form $\mathrm{NH_3}$ and water, minimizing pH change.

Question 11

Which of the following best explains why a buffer composed of acetic acid and acetate cannot effectively maintain a pH of 7.00?

Accepted Answer

The buffer's pKa is too far from the desired pH, so the ratio of base to acid would need to be extremely high.

Question 12

What is the effective pH range for a buffer based on a weak acid with $\mathrm{p}K_a = 4.74$?

Accepted Answer

$3.74$ to $5.74$

Question 13

A buffer solution is prepared by adding a salt containing the conjugate base of a weak acid. Which of the following is a correct example?

Accepted Answer

Adding sodium acetate to acetic acid.

Question 14

Which of the following is a real-world application of buffer solutions?

Accepted Answer

Maintaining blood pH in humans.

Question 15

If the concentration of the conjugate base in a buffer is increased relative to the acid, what happens to the pH?

Accepted Answer

The pH increases.

Question 16

Which of the following is a correct statement about the common ion effect in buffer solutions?

Accepted Answer

The common ion effect suppresses the ionization of the weak acid, stabilizing pH.

Question 17

A buffer is prepared using 0.010 M acetic acid (pKa = 4.74) and acetate. What concentration of acetate is needed to achieve a pH of 7.00?

Accepted Answer

$1.8\ \mathrm{M}$

Question 18

Which of the following scenarios would cause a buffer to lose its effectiveness?

Accepted Answer

Adding a large amount of strong acid or base that exceeds the buffer's capacity.

Buffer Solutions: Principles, Calculations, and Applications

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