Question 1

Which of the following best describes a buffer solution?

Accepted Answer

A solution containing a weak acid and its conjugate base.

Question 2

A 1.0 M HF solution has a pH of 1.6. What will happen to the pH if NaF is added to the solution? (Assume temperature is constant.)

Accepted Answer

The pH will increase above 1.6.

Question 3

What happens to the acid dissociation constant $$K_a$$ for HF if NaF is added to the solution at 25°C?

Accepted Answer

$$K_a$$ remains the same.

Question 4

Which of the following mixtures will result in a buffer solution?

Accepted Answer

10 mL of 1.0 M NH$$_3 + 10 mL of 1.0 M NH$_4Cl$$

Question 5

Which equation is used to calculate the pH of a buffer solution, and under what condition is it valid?

Accepted Answer

$$\mathrm{pH} = \mathrm{p}K_a + \log \left( \frac{[A^-]}{[HA]} \right)$$; valid only for buffer solutions where $$0.1 \u003c \frac{[A^-]}{[HA]} \u003c 10.$$

Question 6

A buffer solution is prepared by mixing 1.0 M HF and 0.5 M NaF. What is the pH of the solution? ($$K_a$$ for HF = $$7.2 	imes 10$$^{-4}$$)

Accepted Answer

$$\mathrm{pH} = 2.84$$

Question 7

What is the net ionic equation for the reaction that occurs when acetic acid ($$\mathrm{CH_3COOH}) is $$titrated with sodium hydroxide ($$\mathrm{NaOH}$$)?

Accepted Answer

$$\mathrm{CH_3COOH} + \mathrm{OH^-} \rightarrow \mathrm{CH_3COO^-} + \mathrm{H_2O}$$

Question 8

During the titration of 50 mL of 1.0 M acetic acid with 1.0 M NaOH, what is the pH at the half-equivalence point? ($$K_a$$ for acetic acid = $$1.8 	imes 10$$^{-5}$$)

Accepted Answer

$$\mathrm{pH} = 4.74$$

Question 9

At the equivalence point in the titration of acetic acid with NaOH, which of the following statements is true?

Accepted Answer

The pH is greater than 7.

Question 10

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 11

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

Accepted Answer

It maintains a nearly constant pH even when diluted with water.

Question 12

A solution is prepared by mixing 40.0 mL of 1.0 M HCl with 30.0 mL of 1.0 M Ba(OH)$$_2. $$What is the nature of the resulting solution?

Accepted Answer

Basic

Question 13

Which of the following statements about the solubility product constant ($$K_{sp}) is $$correct?

Accepted Answer

$$K_{sp} is $$constant at a given temperature.

Question 14

The solubility of Ag$$_2SO$_4 in $$pure water is 0.014 mol/L. What will happen to its solubility if 0.1 M AgNO$$_3 is $$added to the solution?

Accepted Answer

The solubility will decrease.

Question 15

Which of the following best describes the common ion effect?

Accepted Answer

The addition of an ion already present in equilibrium shifts the equilibrium to reduce solubility or ionization.

Question 16

Which of the following is the correct expression for the solubility product constant ($$K_{sp}) of $$CaCO$$_3$$?

Accepted Answer

$$K_{sp} = [Ca$$^{2+}$$] [CO_3$^{2-}]$$

Question 17

If $$Q \u003c K_{sp}$$ for a given salt in solution, what will happen?

Accepted Answer

More solid will dissolve until equilibrium is reached.

Question 18

Which of the following best explains why ocean water is buffered at a pH of about 8.1-8.3?

Accepted Answer

It contains weak acids and their conjugate bases, such as carbonate and bicarbonate ions.

Question 19

Which of the following is a correct application of Le Chatelier's Principle to buffer solutions?

Accepted Answer

Adding a common ion to a buffer will shift the equilibrium to reduce ionization of the weak acid or base.

Acid-Base Equilibria & Solution Equilibria: Buffers and pH Calculations

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