Q: Which equation is used to calculate the pH of a buffer solution when the concentrations of the acid and its conjugate base are known?

$$pH = pK_a + \log \frac{[\text{base}]}{[\text{acid}]}$$

Q: Given $$K_{sp}$$ for $$PbCl_2 is$$ $$1.17 \times 10$$^{-5}$$, what is the molar solubility of $$PbCl_2$$ in pure water?

$$s = 1.50 \times 10$$^{-2}$$ M$$^{1/3}$$

Q: A buffer is most effective when:

The ratio $$[\text{base}]:[\text{acid}] is $$close to 1.

Question 1

Which of the following best describes a buffer solution?

Accepted Answer

A solution that contains a weak acid and its conjugate base, and resists changes in pH when H+ or OH− are added.

Question 2

A buffer is prepared by mixing 0.100 mol of acetic acid ($$\mathrm{HC_2H_3O_2}) $$and 0.100 mol of sodium acetate ($$\mathrm{NaC_2H_3O_2}) in 1.0 L of $$water. If $$K_a = 1.8 	imes 10$$^{-5}$$, what is the pH of the buffer?

Accepted Answer

$$pH = 4.74$$

Question 3

Which equation is used to calculate the pH of a buffer solution when the concentrations of the acid and its conjugate base are known?

Accepted Answer

$$pH = pK_a + \log \frac{[	ext{base}]}{[	ext{acid}]}$$

Question 4

A buffer solution is made by mixing 0.50 M ammonia ($$NH_3) $$and 0.20 M ammonium chloride ($$NH_4Cl). If$$ $$pK_b$$ for ammonia is 4.75, what is the pH of the solution? ($$pK_a + pK_b = 14)$$

Accepted Answer

$$pH = 9.95$$

Question 5

Which of the following statements about buffer capacity is correct?

Accepted Answer

Buffer capacity increases with increasing absolute concentrations of both acid and base.

Question 6

A 50.0 mL sample of 0.200 M NaOH is titrated with 0.200 M HNO$$_3. $$What is the pH after adding 30.0 mL of HNO$$_3$$?

Accepted Answer

$$pH = 12.7$$

Question 7

During the titration of a weak acid with a strong base, why is the pH at the equivalence point greater than 7?

Accepted Answer

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

Question 8

Which indicator would be most appropriate for detecting the equivalence point in the titration of a strong acid with a strong base?

Accepted Answer

Bromothymol blue ($$pH$$ transition range 6.0–7.6)

Question 9

Which of the following best explains why the solubility of $$CaF_2$$ decreases in a solution containing $$NaF$$?

Accepted Answer

The common ion effect reduces the solubility of $$CaF_2$$ due to increased $$F^-$$ concentration.

Question 10

Given $$K_{sp}$$ for $$PbCl_2 is$$ $$1.17 	imes 10$$^{-5}$$, what is the molar solubility of $$PbCl_2$$ in pure water?

Accepted Answer

$$s = 1.50 	imes 10$$^{-2}$$ M$$^{1/3}$$

Question 11

A solution contains $$0.0150 M$$ $$Pb(NO_3$)_2$$ and $$0.00350 M$$ $$NaBr. $$Will $$PbBr_2$$ ($$K_{sp} = 4.67 	imes 10$$^{-6}$$) precipitate?

Accepted Answer

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

Question 12

Which of the following statements about complex ion formation is correct?

Accepted Answer

Complex ion formation can increase the solubility of metal salts by removing metal ions from solution.

Question 13

Which of the following hydroxides is amphoteric and can dissolve in both acidic and basic solutions?

Accepted Answer

$$Al(OH)_3$$

Question 14

A buffer is most effective when:

Accepted Answer

The ratio $$[	ext{base}]:[	ext{acid}] is $$close to 1.

Question 15

Which of the following best describes the relationship between $$K_a$$ and $$K_b$$ for a conjugate acid-base pair?

Accepted Answer

$$K_a 	imes K_b = K_w$$

Question 16

Which of the following is true about the titration curve of a polyprotic acid?

Accepted Answer

It has as many equivalence points as there are ionizable protons.

Question 17

Which of the following compounds will be more soluble in acidic solution than in neutral solution?

Accepted Answer

$$BaF_2$$

Question 18

If $$Q \u003e K_{sp}$$ for a salt in solution, what will happen?

Accepted Answer

Precipitation will occur until $$Q = K_{sp}.$$

Question 19

Which of the following best explains why a supersaturated solution will precipitate when a seed crystal is added?

Accepted Answer

The seed crystal provides a surface for excess solute to crystallize, reducing $$Q to$$ $$K_{sp}.$$

Aqueous Ionic Equilibrium: Buffers, Titrations, and Solubility

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