Q: Which equation is used to calculate the pH of a buffer solution made from a weak acid and its conjugate base?

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

Q: At the half-equivalence point during the titration of a weak acid with a strong base, what is true about the pH and concentrations?

pH = pK$$_a$$ and $$[\text{acid}] = [\text{salt}]$$

Q: Which of the following is the correct formula for buffer capacity?

$$\mathrm{Buffer\ Capacity = \frac{\text{number of moles of OH^-\ or\ H_3O^+\ added}}{(\text{pH change})(\text{volume of buffer in L})}}$$

Question 1

Which of the following best describes the function of a buffer solution?

Accepted Answer

A buffer solution resists changes in pH when small amounts of acids or bases are added.

Question 2

What is the main chemical reaction that occurs when OH$$^-$$ ions are added to an ethanoic acid/sodium ethanoate buffer?

Accepted Answer

$$\mathrm{OH^- (aq) + H^+ (aq) \rightarrow H_2O (l)}$$

Question 3

Which mixture would form a buffer solution with a pH below 7?

Accepted Answer

NaOH (aq) and excess CH$$_3$$COOH (aq)

Question 4

In the blood, which equilibrium helps maintain pH between 7.35 and 7.45?

Accepted Answer

$$\mathrm{CO_2$ (g) + H_2O (l) \rightleftharpoons H^+ (aq) + $$HCO_3$$^- (aq)}$$

Question 5

What happens to the blood pH if the concentration of H$$^+$$ ions increases significantly?

Accepted Answer

The equilibrium shifts to the left, reducing H$$^+$$ concentration and stabilizing pH.

Question 6

Which equation is used to calculate the pH of a buffer solution made from a weak acid and its conjugate base?

Accepted Answer

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

Question 7

A buffer solution is prepared by mixing equal volumes of 1.00 mol dm$$^{-3}$$ ethanoic acid and 1.00 mol dm$$^{-3}$$ sodium ethanoate at 298 K. What is the concentration of CH$$_3$$COOH and CH$$_3$$COO$$^- in $$the final solution?

Accepted Answer

CH$$_3$$COOH = 0.500 mol dm$$^{-3}$$, CH$$_3$$COO$$^- = 0.500 $$mol dm$$^{-3}$$

Question 8

At the half-equivalence point during the titration of a weak acid with a strong base, what is true about the pH and concentrations?

Accepted Answer

pH = pK$$_a$$ and $$[	ext{acid}] = [	ext{salt}]$$

Question 9

A buffer solution is made by mixing 50 cm$$^3 of 0.200 $$mol dm$$^{-3} NH$_3(aq) $$with 50 cm$$^3 of NH$_4Cl(aq). If $$the pH of the solution is 9.55 and $$K_b(NH$_3$$) = $$5.62 	imes 10$$^{-10}$$ mol dm$$^{-3}$$, what is the concentration of NH_4$Cl(aq) used?

Accepted Answer

$$0.112$$ mol dm$$^{-3}$$

Question 10

Which factor most directly increases the buffer capacity of food?

Accepted Answer

Higher protein content

Question 11

Calculate the pH of a buffer solution containing 0.305 mol dm$$^{-3} of $$ethanoic acid and 0.520 mol dm$$^{-3}$$ sodium ethanoate. The $$K_a of $$ethanoic acid is $$1.74 	imes 10$$^{-5}$$ mol dm$$^{-3}$$ at 298 K.

Accepted Answer

pH = 5.00

Question 12

Why is solution A a buffer solution but solution B is not, given the following preparations?
Solution A: 50 cm$$^3 of 0.100 $$mol dm$$^{-3} CH$_3$$COOH(aq) + 25 cm$$^3 of 0.100 $$mol dm$$^{-3}$$ NaOH(aq)
Solution B: 25 cm$$^3 of 0.200 $$mol dm$$^{-3} CH$_3$$COOH(aq) + 50 cm$$^3 of 0.100 $$mol dm$$^{-3}$$ NaOH(aq)

Accepted Answer

Solution A contains both weak acid and its conjugate base in significant amounts, while solution B has excess base.

Question 13

Which of the following is the correct formula for buffer capacity?

Accepted Answer

$$\mathrm{Buffer\ Capacity = \frac{	ext{number of moles of OH^-\ or\ H_3O^+\ added}}{(	ext{pH change})(	ext{volume of buffer in L})}}$$

Question 14

During aerobic respiration, which product contributes to the acid-base equilibrium in blood?

Accepted Answer

CO$$_2$$

Question 15

If a buffer solution is made with a weak base and its conjugate acid, what will be the expected pH range?

Accepted Answer

pH \u003e 7

Question 16

At the equivalence point in the titration of a weak acid with a strong base, what is true about the pH?

Accepted Answer

pH \u003e 7

Question 17

Which of the following best explains why ethanoate ions in a buffer solution prevent the pH from decreasing when H$$^+$$ ions are added?

Accepted Answer

Ethanoate ions react with added H$$^+$$ ions to form ethanoic acid, reducing free H$$^+$$ concentration.

Acid-Base Equilibria and Buffer Solutions

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