Q: A 0.10 mol dm$$^{-3}$$ solution of HCl is prepared. What is the pH of this solution? (Assume complete dissociation.)

$$\text{pH} = -\log(0.10) = 1.00$$

Q: Which equation correctly represents the ionic product of water ($$K_w) at 25$^\circC$$?

$$K_w = [H^+][OH^-] = 1 \times 10$$^{-14}$$~\text{mol}^2~\text{dm}^{-6}$$

Q: Calculate the pH of a 0.010 mol dm$$^{-3}$$ NaOH solution. (Assume complete dissociation and $$K_w = 1 \times 10$$^{-14}$$ at 25$$^\circ$$C.)

$$[OH^-] = 0.010~\text{mol dm}^{-3}$$,~[H^+] = \frac{1 \times $$10^{-14}$$}{0.010} = 1 \times $$10^{-12}$$~\text{mol dm}^{-3}$$,~\text{pH} = 12.00$$

Q: A buffer solution is made by mixing 0.05 mol of ethanoic acid and 0.05 mol of sodium ethanoate in 1.0 dm$$^3 of $$water. Given $$K_a$$ for ethanoic acid is $$1.7 \times 10$$^{-5}$$ mol dm$$^{-3}$$, what is the pH of the buffer?

$$\text{pH} = -\log(1.7 \times 10$$^{-5}$$) = 4.77$$

Question 1

Which of the following best describes the relationship between the acid dissociation constant ($$K_a) $$and the strength of a weak acid?

Accepted Answer

A larger $$K_a$$ value indicates a stronger acid because more acid dissociates.

Question 2

A 0.10 mol dm$$^{-3}$$ solution of HCl is prepared. What is the pH of this solution? (Assume complete dissociation.)

Accepted Answer

$$	ext{pH} = -\log(0.10) = 1.00$$

Question 3

Which equation correctly represents the ionic product of water ($$K_w) at 25$^\circC$$?

Accepted Answer

$$K_w = [H^+][OH^-] = 1 	imes 10$$^{-14}$$~	ext{mol}^2~	ext{dm}^{-6}$$

Question 4

Calculate the pH of a 0.010 mol dm$$^{-3}$$ NaOH solution. (Assume complete dissociation and $$K_w = 1 	imes 10$$^{-14}$$ at 25$$^\circ$$C.)

Accepted Answer

$$[OH^-] = 0.010~	ext{mol dm}^{-3}$$,~[H^+] = \frac{1 	imes $$10^{-14}$$}{0.010} = 1 	imes $$10^{-12}$$~	ext{mol dm}^{-3}$$,~	ext{pH} = 12.00$$

Question 5

A buffer solution is made by mixing 0.05 mol of ethanoic acid and 0.05 mol of sodium ethanoate in 1.0 dm$$^3 of $$water. Given $$K_a$$ for ethanoic acid is $$1.7 	imes 10$$^{-5}$$ mol dm$$^{-3}$$, what is the pH of the buffer?

Accepted Answer

$$	ext{pH} = -\log(1.7 	imes 10$$^{-5}$$) = 4.77$$

Question 6

Which of the following best explains why the pH of a weak acid increases by less than 1 unit when diluted 10 times?

Accepted Answer

Because weak acids only partially dissociate, so the decrease in $$[H^+] is $$less than tenfold.

Question 7

A student adds a small amount of NaOH to an ethanoic acid buffer. What happens to the pH and why?

Accepted Answer

The pH increases slightly because $$OH^-$$ ions react with $$H^+$$, and the buffer equilibrium shifts to replace lost $$H^+.$$

Question 8

Which of the following equations correctly represents the dissociation of a weak acid, HA, in water?

Accepted Answer

$$HA_{(aq)} \rightleftharpoons H^+_{(aq)} + A^-_{(aq)}$$

Question 9

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

Accepted Answer

pH $$\approx 7$$

Question 10

Which indicator would be most suitable for a titration between a weak acid and a strong base, and why?

Accepted Answer

Phenolphthalein, because the pH at the equivalence point is above 7 and phenolphthalein changes color in the basic range.

Question 11

A buffer solution is prepared by mixing 0.200 mol dm$$^{-3}$$ ethanoic acid and 0.250 mol dm$$^{-3}$$ sodium ethanoate. After adding 0.005 mol NaOH to 500 cm$$^3 of $$this buffer, what is the new concentration of ethanoic acid?

Accepted Answer

Question 12

Which of the following best describes the role of a buffer in blood?

Accepted Answer

It maintains the pH of blood between 7.35 and 7.45 by neutralizing small amounts of added acid or base.

Question 13

A solution of a weak acid, HA, has a concentration of 0.10 mol dm$$^{-3}$$ and a $$K_a of$$ $$1.7 	imes 10$$^{-5}$$ mol dm$$^{-3}$$. What is the pH of the solution? (Assume $$[H^+] = \sqrt{K_a [HA]}$$)

Accepted Answer

$$[H^+] = \sqrt{1.7 	imes 10$$^{-5}$$ 	imes 0.10} = 1.3 	imes 10$$^{-3}$$~	ext{mol dm}^{-3}$$,~	ext{pH} = 2.89$

Question 14

Which of the following correctly describes the effect of temperature on the value of $$K_w$$ and the pH of pure water?

Accepted Answer

As temperature increases, $$K_w$$ increases and the pH of pure water decreases.

Question 15

During a titration of a weak acid with a strong base, what is the pH at the half-neutralization point, and why?

Accepted Answer

pH = p$$K_a$$, because $$[HA] = [A^-] at $$this point.

Question 16

Which of the following best explains why a buffer solution resists changes in pH when small amounts of acid are added?

Accepted Answer

The added $$H^+$$ ions react with the conjugate base present in the buffer, minimizing the increase in $$[H^+].$$

Question 17

A student is asked to calculate the pH of a buffer solution after adding a small amount of strong acid. Which equation should they use?

Accepted Answer

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

Question 18

Which of the following statements about the enthalpy change of neutralization is correct?

Accepted Answer

The enthalpy change of neutralization is always exothermic for reactions between strong acids and strong bases.

Acids, Bases, pH, and Buffer Solutions: Study Notes

Study Guide - Practice Questions

Study Guide - Flashcards