Q: Which equation is the Henderson-Hasselbalch equation for a buffer solution?

$pH = pK_a + \log \left( \frac{[\text{base}]}{[\text{acid}]} \right)$

Question 1

Which of the following best describes the common ion effect in an aqueous solution?

Accepted Answer

The shift in equilibrium that occurs when an ion already present in the equilibrium is added to the solution.

Question 2

A solution is prepared by mixing 1.0 M HF and 1.0 M NaF. Given $K_a = 7.2 	imes 10^{-4}$ for HF, what is the main effect of adding NaF to the HF solution?

Accepted Answer

It increases the concentration of $F^-$, shifting the equilibrium to the left and increasing the pH.

Question 3

Which of the following mixtures will create a buffer solution?

Accepted Answer

0.10 M CH$_3$COOH and 0.10 M CH$_3$COONa

Question 4

A buffer is made by mixing 0.50 M acetic acid ($HC_2H_3O_2$) and 0.50 M sodium acetate ($NaC_2H_3O_2$). If 0.010 mol NaOH is added to 1.0 L of this buffer, what will happen to the pH?

Accepted Answer

The pH will increase only slightly.

Question 5

Which equation is the Henderson-Hasselbalch equation for a buffer solution?

Accepted Answer

$pH = pK_a + \log \left( \frac{[	ext{base}]}{[	ext{acid}]} ight)$

Question 6

A buffer is prepared with 0.75 M lactic acid ($HC_3H_5O_3$, $K_a = 1.4 	imes 10^{-4}$) and 0.25 M sodium lactate. What is the pH of the buffer? (Use the Henderson-Hasselbalch equation.)

Accepted Answer

pH = 3.85 + \log \left( \frac{0.25}{0.75} \right ) = 3.85 - 0.48 = 3.37

Question 7

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 buffer contains a weak acid and its conjugate base, which neutralize added $H^+$ or $OH^-$.

Question 8

A buffer is made from 0.10 M $NH_3$ and 0.10 M $NH_4Cl$. If the $K_b$ for $NH_3$ is $1.8 	imes 10^{-5}$, what is the pH of the buffer? ($K_w = 1.0 	imes 10^{-14}$)

Accepted Answer

$pOH = pK_b + \log \left( \frac{[NH_4^+]}{[NH_3]} \right ) = 4.74 + \log(1) = 4.74$, so $pH = 14 - 4.74 = 9.26$

Question 9

During the titration of a strong acid with a strong base, what is true at the equivalence point?

Accepted Answer

The pH is always 7.00 at 25°C.

Question 10

A 50.0 mL sample of 0.100 M HCN ($K_a = 4.9 	imes 10^{-10}$) is titrated with 0.100 M NaOH. What is the pH at the equivalence point?

Accepted Answer

pH = 11.00

Question 11

Which of the following statements about the effect of $K_a$ on a weak acid–strong base titration curve is correct?

Accepted Answer

A smaller $K_a$ results in a higher pH at the equivalence point.

Question 12

Which of the following best describes buffering capacity?

Accepted Answer

The amount of acid or base a buffer can absorb without a significant change in pH.

Question 13

A buffer is prepared with 0.10 mol $HC_2H_3O_2$ and 0.050 mol $NaC_2H_3O_2$ in 1.0 L. If 0.010 mol HCl is added, what is the new pH? ($K_a = 1.8 	imes 10^{-5}$)

Accepted Answer

$pH = 4.56$

Question 14

Which of the following is NOT a characteristic of a buffer solution?

Accepted Answer

It contains a strong acid and its conjugate base.

Question 15

If you want to prepare a buffer with pH = 9.00 using $NH_3$ ($K_b = 1.8 	imes 10^{-5}$) and $NH_4Cl$, what ratio of $[NH_4^+]/[NH_3]$ should you use? ($pK_a = 9.25$)

Accepted Answer

$\frac{[NH_4^+]}{[NH_3]} = 10^{9.25-9.00} = 1.78$

Question 16

Which of the following best explains why the pH of a buffer changes only slightly when a small amount of strong acid is added?

Accepted Answer

The strong acid is neutralized by the conjugate base present in the buffer.

Question 17

Which of the following is the correct expression for the acid dissociation constant $K_a$ for a weak acid $HA$?

Accepted Answer

$K_a = \frac{[H^+][A^-]}{[HA]}$

Chapter 17: Additional Aspects of Aqueous Equilibria – Buffers, Common Ion Effect, and Acid-Base Titrations

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