Question 1

Which of the following best describes the effect of adding a small amount of strong acid (HCl) to a buffer solution containing acetic acid and sodium acetate?

Accepted Answer

The pH will decrease slightly because the buffer will neutralize most of the added acid.

Question 2

A solution is prepared by mixing 0.10 mol of NH$_3$ and 0.10 mol of NH$_4$Cl in 1.0 L of water. Which equation should be used to calculate the pH of this buffer solution?

Accepted Answer

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

Question 3

Which of the following statements about polyprotic acids is correct?

Accepted Answer

Each successive proton is harder to remove, so $K_{a1} > K_{a2} > K_{a3}$.

Question 4

During a titration of a weak acid with a strong base, which region of the titration curve corresponds to the buffer region?

Accepted Answer

Between the initial addition of base and before the equivalence point, where both acid and its conjugate base are present.

Question 5

A student is asked to calculate the pH of a 0.01 M solution of HCl. Which is the correct approach?

Accepted Answer

Assume complete dissociation and use $\mathrm{pH} = -\log[\mathrm{H}^+]$.

Question 6

Which of the following best explains why the addition of NaCl to a saturated solution of AgCl causes AgCl to precipitate?

Accepted Answer

The common ion effect reduces the solubility of AgCl by increasing Cl$^-$ concentration, shifting equilibrium toward precipitation.

Question 7

If $K_{sp}$ for CaF$_2$ is $3.2 	imes 10^{-11}$, what is the molar solubility of CaF$_2$ in pure water?

Accepted Answer

$s = \sqrt{\dfrac{K_{sp}}{4}}$ where $K_{sp} = [Ca^{2+}][F^-]^2$, so $s = \sqrt[3]{\dfrac{K_{sp}}{4}}$

Question 8

Which of the following is true regarding the equivalence point in a titration of a weak acid with a strong base?

Accepted Answer

The pH at the equivalence point is greater than 7 due to the formation of a weak conjugate base.

Question 9

A buffer is prepared using 0.50 mol of acetic acid and 0.25 mol of sodium acetate in 1.0 L of solution. What is the pH if $pK_a$ of acetic acid is 4.76?

Accepted Answer

$\mathrm{pH} = 4.76 + \log\left(\dfrac{0.25}{0.50}\right)$

Question 10

Which of the following is the correct expression for the relationship between $K_a$, $K_b$, and $K_w$ for a conjugate acid-base pair?

Accepted Answer

$K_a 	imes K_b = K_w$

Question 11

A solution contains 0.10 M HF (weak acid, $K_a = 7.2 	imes 10^{-4}$). What is the pH of the solution?

Accepted Answer

Set up an ICE table and solve for $[H^+]$ using $K_a = \dfrac{[H^+][F^-]}{[HF]}$

Question 12

Which of the following best describes buffer capacity?

Accepted Answer

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

Question 13

Which of the following is true about the pH of a solution of a salt formed from a strong acid and a weak base (e.g., NH$_4$Cl)?

Accepted Answer

The solution will be acidic due to hydrolysis of the weak base's conjugate acid.

Question 14

Which of the following equations is used to calculate pOH from hydroxide ion concentration?

Accepted Answer

$\mathrm{pOH} = -\log[OH^-]$

Question 15

A saturated solution of BaSO$_4$ has $K_{sp} = 1.1 	imes 10^{-10}$. If [Ba$^{2+}$] = $1.0 	imes 10^{-5}$ M, what is [SO$_4^{2-}$] at equilibrium?

Accepted Answer

$[SO_4^{2-}] = \dfrac{K_{sp}}{[Ba^{2+}]}$

Question 16

Which of the following is the correct use of a RICE table in acid-base chemistry?

Accepted Answer

To determine equilibrium concentrations for weak acid or base dissociation.

Question 17

Which of the following best describes what happens when $Q_{sp} > K_{sp}$ for a slightly soluble compound?

Accepted Answer

Precipitation occurs because the solution is supersaturated.

Question 18

Which of the following is true about the effect of pH on the solubility of a salt containing a basic anion (e.g., CaCO$_3$)?

Accepted Answer

Decreasing pH increases solubility due to protonation of the anion.

Acid-Base Chemistry, Buffers, Acid-Base Titrations, and Equilibria of Slightly Soluble Compounds

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