Q: You are asked to prepare a buffer with pH 7.00 using acetic acid ($$\mathrm{p}K_a = 4.74) $$and sodium acetate. If the acetic acid concentration is 0.010 M, what must the acetate concentration be? (Use the Henderson-Hasselbalch equation.)

$$\mathrm{pH} = 4.74 + \log \left( \frac{[\text{acetate}]}{0.010} \right) = 7.00$$, so $$\log \left( \frac{[\text{acetate}]}{0.010} \right) = 2.26$$, $$[\text{acetate}] = 0.010 \times 10$$^{2.26}$$ = 0.182 M$$

Question 1

Which of the following best describes the function of a buffer solution?

Accepted Answer

A buffer solution maintains a constant pH by neutralizing added acids or bases through the presence of a weak acid and its conjugate base (or weak base and its conjugate acid).

Question 2

A buffer is prepared by mixing 0.30 mol of acetic acid ($$\mathrm{CH_3COOH}$$, $$K_a = 1.8 	imes 10$$^{-5}$$) and 0.30 mol of sodium acetate in enough water to make 1.0 L of solution. What is the pH of this buffer? (Use the Henderson-Hasselbalch equation: $$\mathrm{pH} = \mathrm{p}K_a + \log \left( \frac{[	ext{A}^-]}{[	ext{HA}]} ight)$$)

Accepted Answer

$$\mathrm{pH} = 4.74 + \log \left( \frac{0.30}{0.30} \right) = 4.74$$

Question 3

Which of the following best explains the common ion effect as it relates to buffer solutions?

Accepted Answer

The common ion effect refers to the shift in equilibrium that occurs when a solution already contains one of the ions involved in the equilibrium, thereby suppressing the ionization of a weak acid or base.

Question 4

A buffer is made by mixing 0.12 M lactic acid ($$K_a = 1.4 	imes 10$$^{-4}$$) and 0.10 M sodium lactate. What is the pH of the buffer? (Use $$\mathrm{p}K_a = -\log K_a$$ and the Henderson-Hasselbalch equation.)

Accepted Answer

$$\mathrm{p}K_a = 3.85$$, $$\mathrm{pH} = 3.85 + \log \left( \frac{0.10}{0.12} \right) = 3.72$$

Question 5

Which of the following statements about buffer capacity is correct?

Accepted Answer

Buffer capacity is greatest when the concentrations of the weak acid and its conjugate base are both high and approximately equal.

Question 6

A buffer solution is prepared by adding ammonium chloride to a 0.500 M ammonia solution. If the desired pH is 9.20 and $$K_b$$ for ammonia is $$1.80 	imes 10$$^{-5}$$, what is the required ratio of $$[\mathrm{$$NH_4$$^+}]/[\mathrm{$$NH_3$$}]$$? (Use $$\mathrm{pOH} = \mathrm{p}K_b + \log \left( \frac{[\mathrm{$$NH_4$$^+}]}{[\mathrm{$$NH_3$$}]} ight)$$ and $$\mathrm{pH} + \mathrm{pOH} = 14$$.)

Accepted Answer

$$\mathrm{p}K_b = 4.74$$, $$\mathrm{pOH} = 14 - 9.20 = 4.80$$, $$4.80 = 4.74 + \log \left( \frac{[\mathrm{NH_4$^+}]}{0.500} \right)$$, so $$\log \left( \frac{[\mathrm{NH_4$^+}]}{0.500} \right) = 0.06$$, $$\frac{[\mathrm{NH_4$^+}]}{0.500} = 10$$^{0.06}$$ = 1.15$$, $$[\mathrm{NH_4$^+}]/[\mathrm{NH_3$}] = 1.15$$

Question 7

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 of the buffer?

Accepted Answer

The pH remains essentially unchanged because the ratio $$\frac{[	ext{A}^-]}{[	ext{HA}]}$$ does not change.

Question 8

Which of the following combinations would NOT result in a buffer solution?

Accepted Answer

Mixing equal volumes of 0.10 M hydrochloric acid and 0.10 M sodium chloride.

Question 9

A buffer is most effective when:

Accepted Answer

The concentrations of the weak acid and its conjugate base are equal, and the pH is close to the $$\mathrm{p}K_a of $$the acid.

Question 10

You are asked to prepare a buffer with pH 7.00 using acetic acid ($$\mathrm{p}K_a = 4.74) $$and sodium acetate. If the acetic acid concentration is 0.010 M, what must the acetate concentration be? (Use the Henderson-Hasselbalch equation.)

Accepted Answer

$$\mathrm{pH} = 4.74 + \log \left( \frac{[	ext{acetate}]}{0.010} ight) = 7.00$$, so $$\log \left( \frac{[	ext{acetate}]}{0.010} ight) = 2.26$$, $$[	ext{acetate}] = 0.010 	imes 10$$^{2.26}$$ = 0.182 M$$

Question 11

Which of the following best describes what happens when a small amount of strong acid is added to a buffer composed of ammonia ($$\mathrm{NH_3$}) $$and ammonium ion ($$\mathrm{NH_4$^+}$$)?

Accepted Answer

The added acid reacts with $$\mathrm{NH_3$} to $$form $$\mathrm{NH_4$^+}$$, minimizing the change in pH.

Question 12

A hospital needs to prepare a buffer solution to maintain blood pH at 7.4. Which buffer system would be most appropriate, considering the effective pH range of a buffer?

Accepted Answer

A buffer with a $$\mathrm{p}K_a$$ close to 7.4, such as the $$\mathrm{H_2PO_4^-}/\mathrm{$$HPO_4$$^{2-}}$$ system ($$\mathrm{p}K_a = 7.2).$$

Question 13

Which of the following statements about the Henderson-Hasselbalch equation is FALSE?

Accepted Answer

It can be used for strong acid/strong base solutions.

Question 14

Calculate the number of grams of ammonium chloride ($$\mathrm{NH_4Cl}$$, molar mass = 53.5 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}$$, $$\mathrm{p}K_b = 4.74$$)

Accepted Answer

First, $$\mathrm{pOH} = 14 - 9.20 = 4.80$$. $$4.80 = 4.74 + \log \left( \frac{[\mathrm{NH_4$^+}]}{0.500} ight)$$, so $$\log \left( \frac{[\mathrm{NH_4$^+}]}{0.500} ight) = 0.06$$, $$[\mathrm{NH_4$^+}] = 0.500 	imes 10$$^{0.06}$$ = 0.575 M. $$Moles needed: $$0.575 	imes 2.00 = 1.15$$ mol. Mass: $$1.15 	imes 53.5 = 61.5 g.$$

Question 15

Which of the following best describes the effect of adding a small amount of strong base to an acetic acid/acetate buffer?

Accepted Answer

The added base reacts with acetic acid to form acetate, minimizing the change in pH.

Question 16

Which of the following is a limitation of buffer solutions?

Accepted Answer

Buffers have a limited capacity and can only neutralize small amounts of added acid or base before the pH changes significantly.

Buffer Solutions: Principles, Calculations, and Applications

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