Q: Which of the following acid-conjugate base combinations would result in an ideal buffer solution with a pH of 5.7?

Nitrous acid ($$HNO_2) $$and sodium nitrite ($$K_a = 4.6 \times 10$$^{-4}$$)

Question 1

Which of the following statements best describes the purpose of the Henderson-Hasselbalch equation in buffer chemistry?

Accepted Answer

It allows calculation of the pH of a buffer solution using the concentrations of a weak acid and its conjugate base.

Question 2

Given a buffer solution containing 2.0 M nitrous acid ($$HNO_2) $$and 1.48 M lithium nitrite ($$LiNO_2$$), with $$K_a = 4.6 	imes 10$$^{-4}$$, what is the pH of the solution? (Use pK_a$$ = -\log K_a$$)

Accepted Answer

$$3.21$$

Question 3

A buffer is most effective when:

Accepted Answer

The concentrations of the weak acid and its conjugate base are approximately equal.

Question 4

What is the correct form of the Henderson-Hasselbalch equation for a buffer solution?

Accepted Answer

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

Question 5

A buffer is prepared by mixing 500 mL of 1.4 M $$C_6H_5NH_2$$ and 220 mL of 1.2 M $$C_6H_5NH_3^+. If$$ $$K_b$$ for $$C_6H_5NH_2 is$$ $$3.8 	imes 10$$^{-10}$$, what is the pH of the buffer? (Assume pK_b$$ = -\log K_b$$ and pK_a + pK_b = 14$)

Accepted Answer

$$5.23$$

Question 6

If a buffer solution has a pH of 4.7 and is made from boric acid ($$H_3BO_3) $$with $$K_a = 5.4 	imes 10$$^{-10}$$, what is the required ratio of $$[CB]/[WA]$$ (conjugate base to weak acid) in the buffer?

Accepted Answer

$$91$$

Question 7

Calculate the mass of NaO2H (sodium hydroxide) that must be added to 1.0 L of 0.35 M HBr to make a buffer with a pH of 8.5. $$K_a of $$hypobromous acid is $$1.8 	imes 10$$^{-9}$$. (Molar mass of NaO2H = 56 g/mol)

Accepted Answer

$$10.5 g$$

Question 8

Which of the following acid-conjugate base combinations would result in an ideal buffer solution with a pH of 5.7?

Accepted Answer

Nitrous acid ($$HNO_2) $$and sodium nitrite ($$K_a = 4.6 	imes 10$$^{-4}$$)

Question 9

The buffering range of a solution containing lactic acid ($$K_a = 1.4 	imes 10$$^{-4}$$) and sodium lactate is best described as:

Accepted Answer

$$2.85 to$$ $$4.85$$

Question 10

A buffer solution is prepared by mixing equal concentrations of a weak acid and its conjugate base. What is the relationship between the pH and the $$pK_a of $$the acid in this solution?

Accepted Answer

The pH is equal to the $$pK_a.$$

Question 11

If the $$[A^-]/[HA]$$ ratio in a buffer is increased by a factor of 10, how does the pH change?

Accepted Answer

The pH increases by 1 unit.

Question 12

Which of the following best explains why a buffer resists changes in pH upon addition of small amounts of acid or base?

Accepted Answer

The weak acid neutralizes added base and the conjugate base neutralizes added acid.

Question 13

A buffer solution is made by mixing 0.5 mol of acetic acid ($$CH_3COOH) $$and 0.5 mol of sodium acetate ($$CH_3COONa) in 1.0 L of $$water. If $$K_a$$ for acetic acid is $$1.8 	imes 10$$^{-5}$$, what is the pH of the buffer?

Accepted Answer

$$4.74$$

Question 14

Which of the following would NOT be a suitable buffer system for maintaining a pH near 9?

Accepted Answer

Acetic acid and sodium acetate ($$pK_a \approx 4.74)$$

Question 15

A buffer solution is prepared with $$[HA] = 0.10 M $$and $$[A^-] = 0.20 M. If$$ $$pK_a = 4.75$$, what is the pH of the solution?

Accepted Answer

$$5.05$$

Question 16

If you want to prepare a buffer with a pH of 6.0 using a weak acid with $$pK_a = 5.0$$, what should the $$[A^-]/[HA]$$ ratio be?

Accepted Answer

$$10$$

Question 17

Which of the following real-life scenarios best illustrates the importance of buffer solutions?

Accepted Answer

Maintaining blood pH within a narrow range in the human body.

Henderson-Hasselbalch Equation and Buffer Solutions

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