Question 1

Which of the following best explains why glycolic acid is toxic when ingested in large amounts?

Accepted Answer

It overwhelms the blood's buffer system, causing acidosis and reducing oxygen transport.

Question 2

A buffer solution is prepared by mixing 0.100 M acetic acid ($$\mathrm{HC_2H_3O_2}) $$and 0.100 M sodium acetate ($$\mathrm{NaC_2H_3O_2}). $$Which equation correctly represents the acid dissociation equilibrium for this buffer?

Accepted Answer

$$\mathrm{HC_2H_3O_2} (aq) + H_2O (l) \rightleftharpoons C_2H_3O_2^- (aq) + H_3O^+ (aq)$$

Question 3

Which of the following best describes the common ion effect?

Accepted Answer

Adding a salt containing the conjugate base of a weak acid suppresses the acid's ionization, raising the pH.

Question 4

Calculate the pH of a buffer solution containing 0.14 M HF ($$pK_a = 3.15) $$and 0.071 M KF using the Henderson-Hasselbalch equation.

Accepted Answer

$$\mathrm{pH = 3.15 + \log\left(\frac{0.071}{0.14}\right) = 2.86}$$

Question 5

A buffer is most effective when:

Accepted Answer

The concentrations of acid and base are equal and large.

Question 6

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

Accepted Answer

It is valid when the 'x is small' approximation holds, typically when $$[	ext{acid}]$$ and $$[	ext{base}]$$ are much greater than $$K_a.$$

Question 7

A buffer solution contains 0.100 mol $$\mathrm{HC_2H_3O_2}$$ and 0.100 mol $$\mathrm{NaC_2H_3O_2} in 1.00 L. If 0.010 $$mol NaOH is added, what is the new concentration of $$\mathrm{HC_2H_3O_2}$$?

Accepted Answer

0.090 M

Question 8

Which of the following best describes the effect of adding a strong acid to a buffer solution?

Accepted Answer

The added acid reacts with the conjugate base, forming more weak acid and minimizing pH change.

Question 9

Which of the following equations correctly relates $$pK_a$$ and $$pK_b$$ for a conjugate acid-base pair?

Accepted Answer

$$pK_a + pK_b = 14$$

Question 10

A titration curve for a strong acid with a strong base shows a sharp rise in pH at the equivalence point. What is the pH at the equivalence point if both solutions are 0.100 M?

Accepted Answer

pH = 7.00

Question 11

Which of the following best explains why the pH at the equivalence point of a weak acid-strong base titration is greater than 7?

Accepted Answer

The conjugate base of the weak acid hydrolyzes to produce OH$$^-$$ ions.

Question 12

Which of the following is the correct expression for the solubility product constant ($$K_{sp}) $$for $$\mathrm{$$PbCl_2$$}$$?

Accepted Answer

$$K_{sp} = [Pb$$^{2+}$$][Cl^-]^2$$

Question 13

If the reaction quotient $$Q$$ for a salt is less than $$K_{sp}$$, what will happen?

Accepted Answer

The solution is unsaturated and more salt can dissolve.

Question 14

Which of the following will increase the solubility of $$\mathrm{AgCl} in $$water?

Accepted Answer

Adding NH$$_3 to $$the solution, which forms a complex ion with Ag$$^+.$$

Question 15

A solution contains 0.0175 M Ca(NO$$_3$$)$$_2$$ and 0.0175 M NaOH. Will Ca(OH)$$_2$$ precipitate? ($$K_{sp} of Ca(OH$$)$$_2$$ = $$4.68 	imes 10$$^{-6}$$)

Accepted Answer

Yes, because $$Q \u003e K_{sp}.$$

Question 16

Which of the following best describes an amphoteric hydroxide?

Accepted Answer

A hydroxide that is more soluble in both acidic and basic solutions due to its ability to act as an acid or base.

Question 17

Which of the following is a correct statement about indicators in acid-base titrations?

Accepted Answer

An indicator should change color at a pH close to the equivalence point of the titration.

Question 18

Which of the following best describes selective precipitation in qualitative analysis?

Accepted Answer

Adding a reagent that forms an insoluble salt with one cation but not others, allowing separation.

Aqueous Ionic Equilibrium: Buffers, Titrations, and Solubility

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