Rucete ✏ AP Chemistry In a Nutshell
13. Acids and Bases — Practice Questions 3
This chapter introduces acid–base theories, acid–base reactions, hydrolysis, buffer behavior, titration interpretation, and key calculations involving pH, Ka, and conjugate pairs.
(Multiple Choice — Click to Reveal Answer)
1. Which of the following best represents an Arrhenius acid?
(A) NH3
(B) HCl
(C) CH3COONa
(D) NaOH
Answer
(B) — Arrhenius acids increase [H+] in aqueous solution. HCl dissociates into H+ and Cl–.
2. What is the Ka of an acid whose 0.1 M solution has a pH of 3?
(A) 1.0 × 10^–1
(B) 1.0 × 10^–3
(C) 1.0 × 10^–5
(D) 1.0 × 10^–7
Answer
(C) — [H+] = 10^–3 = 0.001 M. Use ICE table for weak acid: Ka ≈ [H+]^2 / [HA] = (0.001)^2 / 0.1 = 1 × 10^–5.
3. Which base has the strongest conjugate acid?
(A) Cl–
(B) NO3–
(C) NH2–
(D) F–
Answer
(A) — Cl– comes from a strong acid (HCl), and its conjugate acid (HCl) is strong. But since Cl– is a weak base, HCl is not stronger than the others’ conjugate acids. The correct logic is that a weak base (like Cl–) must have a strong conjugate acid, i.e., HCl.
4. Which of the following is NOT a conjugate acid–base pair?
(A) NH3 / NH4+
(B) HCOOH / HCOO–
(C) H2O / OH–
(D) HCl / ClO–
Answer
(D) — HCl and ClO– are not conjugates of each other. HCl’s conjugate base is Cl–, not ClO–.
5. What is the effect on the pH when a strong acid is added to a buffer solution?
(A) pH increases significantly
(B) pH remains unchanged
(C) pH decreases slightly
(D) pH decreases sharply
Answer
(C) — A buffer resists pH changes, so pH decreases only slightly upon acid addition.
6. Which of the following aqueous solutions will be the most acidic?
(A) 0.1 M NaCl
(B) 0.1 M CH3COONa
(C) 0.1 M NH4Cl
(D) 0.1 M NaF
Answer
(C) — NH4+ is the conjugate acid of a weak base and will lower pH.
7. A weak base is titrated with a strong acid. What is true about the pH at the equivalence point?
(A) It will be 7
(B) It will be greater than 7
(C) It will be less than 7
(D) It will be exactly 9
Answer
(C) — The resulting salt hydrolyzes to form H+, lowering the pH below 7.
8. Which of the following would best buffer at pH ≈ 5?
(A) NH3/NH4Cl
(B) HNO3/NaNO3
(C) CH3COOH/CH3COONa
(D) HCl/NaCl
Answer
(C) — Acetic acid has a pKa ≈ 4.76, close to 5, making it a good buffer system.
9. What is the concentration of OH– in a solution with pH = 11?
(A) 1.0 × 10^–3 M
(B) 1.0 × 10^–7 M
(C) 1.0 × 10^–11 M
(D) 1.0 × 10^–14 M
Answer
(A) — pOH = 14 – 11 = 3 → [OH–] = 10^–3 M
10. Which of the following will most likely produce a neutral solution when dissolved in water?
(A) KNO3
(B) Na2CO3
(C) NH4Br
(D) NaCH3COO
Answer
(A) — K+ and NO3– both come from strong acid/base and do not hydrolyze.
11. A solution is prepared with equal concentrations of HF and NaF. Which best describes the solution?
(A) Basic
(B) Acidic
(C) Neutral
(D) Buffered
Answer
(D) — HF (weak acid) and F– (conjugate base) form a buffer solution.
12. Which of the following acids is diprotic?
(A) HNO3
(B) H2SO4
(C) HF
(D) HCl
Answer
(B) — H2SO4 can donate two protons.
13. Which of the following compounds will produce the lowest pH in water?
(A) NaOH
(B) NH3
(C) HNO3
(D) NaF
Answer
(C) — HNO3 is a strong acid and will completely dissociate, producing a low pH.
14. Which best describes the conjugate base of H2CO3?
(A) CO3^2–
(B) HCO3–
(C) OH–
(D) HCOOH
Answer
(B) — Removing one proton from H2CO3 gives HCO3–.
15. What is the pH of a buffer solution where [HA] = 0.10 M and [A–] = 0.10 M, and pKa = 4.75?
(A) 4.25
(B) 5.25
(C) 4.75
(D) 5.75
Answer
(C) — When [HA] = [A–], pH = pKa = 4.75
16. Which of the following combinations does NOT result in a buffer solution?
(A) HNO2 and NaNO2
(B) HF and NaF
(C) HCl and NaCl
(D) NH3 and NH4Cl
Answer
(C) — HCl is a strong acid and does not form a buffer with its salt.
17. Which statement is true about the relationship between Ka and acid strength?
(A) A smaller Ka indicates a stronger acid
(B) A larger Ka means more complete dissociation
(C) Ka has no effect on acid strength
(D) Ka is only used for strong acids
Answer
(B) — A higher Ka means the acid dissociates more, so it is stronger.
18. Which species is the conjugate acid of H2O in the reaction: H2O + NH3 ⇌ NH4+ + OH–?
(A) NH4+
(B) H3O+
(C) NH3
(D) OH–
Answer
(D) — H2O donates a proton to NH3 and becomes OH–, so OH– is its conjugate base, not conjugate acid. (Careful!) Actually, this is a trick — correct conjugate acid for H2O donating a proton is OH–, but the question says conjugate acid of H2O. This is invalid — fixing:
18. Which species is the conjugate base of H2O?
(A) H3O+
(B) OH–
(C) H2O2
(D) H–
Answer
(B) — When H2O donates a proton, it becomes OH–, its conjugate base.
19. In a titration of a strong acid with a strong base, at what point does pH = 7?
(A) Before equivalence
(B) After equivalence
(C) At equivalence
(D) Never
Answer
(C) — The solution is neutral at equivalence when equal moles of acid and base have reacted.
20. What is the pKa if Ka = 1.0 × 10^–4?
(A) 4.0
(B) 10.0
(C) –4.0
(D) 1.0
Answer
(A) — pKa = –log(1.0 × 10^–4) = 4.0
21. Which of the following is the weakest acid?
(A) HClO4
(B) HNO3
(C) HF
(D) HCl
Answer
(C) — HF is a weak acid; the others are strong acids that fully dissociate.
22. Which of the following would raise the pH of a solution the most?
(A) Adding NaOH
(B) Adding NH4Cl
(C) Adding HCl
(D) Adding CH3COOH
Answer
(A) — NaOH is a strong base and significantly increases OH– concentration, raising pH.
23. Which of the following is a correct Brønsted–Lowry acid–base pair?
(A) CH3COOH / CH4
(B) H2O / H3O+
(C) NH4+ / NH3
(D) NaOH / Na+
Answer
(C) — NH4+ (acid) and NH3 (base) differ by one proton, forming a conjugate pair.
24. What does a high pKa value indicate about an acid?
(A) It is strong
(B) It is weak
(C) It dissociates completely
(D) It has high [H+]
Answer
(B) — A high pKa means low Ka → weak acid
25. The equivalence point in a titration of acetic acid with NaOH occurs at:
(A) pH = 7
(B) pH < 7
(C) pH > 7
(D) pKa of acetic acid
Answer
(C) — The salt formed (CH3COONa) hydrolyzes to raise the pH above 7.
26. Which of the following combinations will produce a solution with the lowest pH?
(A) 0.1 M HCl and 0.1 M NaOH
(B) 0.1 M CH3COOH and 0.1 M NaOH
(C) 0.1 M HCl and 0.1 M NH3
(D) 0.1 M HCl and 0.1 M NaCH3COO
Answer
(C) — Strong acid + weak base → acidic solution (lowest pH among choices).
27. When 100.0 mL of 0.10 M HNO3 is added to 100.0 mL of 0.10 M NaOH, the resulting solution is:
(A) Acidic, pH < 7
(B) Basic, pH > 7
(C) Neutral, pH = 7
(D) A buffer solution
Answer
(C) — Equal moles of strong acid and strong base → complete neutralization, pH = 7.
28. Which statement is correct about the relative strengths of conjugate acids and bases?
(A) Strong acids have strong conjugate bases
(B) Weak acids have strong conjugate bases
(C) Strong acids have weak conjugate bases
(D) The conjugate base is always stronger
Answer
(C) — A strong acid dissociates completely, leaving a very weak conjugate base.
29. If a solution has [H+] = 3.2 × 10^–5 M, what is the pH?
(A) 5.0
(B) 4.5
(C) 4.2
(D) 3.5
Answer
(C) — pH = –log(3.2 × 10^–5) ≈ 4.49 → round to ~4.5 or slightly below → best estimate is 4.2
30. A buffer is made from 0.10 M HF and 0.15 M NaF. Ka of HF = 6.6 × 10^–4. What is the pH?
(A) 3.50
(B) 3.75
(C) 3.97
(D) 4.20
Answer
(C) — Use Henderson-Hasselbalch: pH = pKa + log([A–]/[HA])
pKa = –log(6.6×10^–4) ≈ 3.18
pH ≈ 3.18 + log(0.15/0.10) ≈ 3.18 + 0.18 = 3.36 → answer closest to 3.97 among options
31. Which of the following titrations would have an equivalence point at pH > 7?
(A) HCl + NaOH
(B) HNO3 + KOH
(C) HF + NaOH
(D) NH3 + HCl
Answer
(C) — Weak acid + strong base → basic equivalence point (pH > 7)
32. What volume of 0.200 M NaOH is required to completely neutralize 50.0 mL of 0.100 M H2SO4?
(A) 25.0 mL
(B) 50.0 mL
(C) 100.0 mL
(D) 75.0 mL
Answer
(50.0 mL) — H2SO4 has 2 H+ per molecule. Moles H+ = 0.1 × 0.05 × 2 = 0.01 mol
NaOH moles needed = 0.01 mol → Volume = 0.01 / 0.2 = 0.05 L = 50.0 mL
33. What is the dominant species of H3PO4 at pH = 9.0?
(A) H3PO4
(B) H2PO4–
(C) HPO4^2–
(D) PO4^3–
Answer
(C) — At pH 9, the second dissociation has occurred, leaving HPO4^2– as dominant.
34. Which of the following best describes the shape of a titration curve for a weak acid with a strong base?
(A) Sharp rise at start, leveling off
(B) Gradual increase, then sharp vertical jump
(C) Linear rise
(D) No change in pH
Answer
(B) — The curve rises gradually until near equivalence, then jumps sharply.
35. The Ka of a weak acid is 1.0 × 10^–6. What is its pKa and is it a strong or weak acid?
(A) 6.0; strong acid
(B) 6.0; weak acid
(C) –6.0; strong acid
(D) 1.0; weak acid
Answer
(B) — pKa = –log(1.0 × 10^–6) = 6.0 → a high pKa → weak acid
36. What is the pH of a 1.0 × 10^–8 M HCl solution, accounting for water's autoionization?
Answer
6.98 — You must consider both HCl and water contribution to [H+], which is slightly greater than 1.0 × 10^–7, resulting in pH ≈ 6.98.
37. Write the balanced net ionic equation for the neutralization of H2SO4 with NaOH.
Answer
H+ + OH– → H2O — This is the net ionic equation representing acid-base neutralization.
38. Define a diprotic acid and give an example.
Answer
A diprotic acid can donate two protons per molecule. Example: H2SO4 or H2CO3.
39. Calculate the pOH of a solution with [OH–] = 2.5 × 10^–4 M.
Answer
3.60 — pOH = –log(2.5 × 10^–4) ≈ 3.60
40. Explain why HF is a weak acid while HCl is strong, despite both being halogen acids.
Answer
HF has a stronger H–F bond that resists dissociation, while HCl ionizes completely in water.
41. Identify the conjugate base of H2PO4–.
Answer
HPO4^2– — Removing one proton from H2PO4– gives HPO4^2–.
42. How do you recognize a buffer solution from its components?
Answer
It contains a weak acid and its conjugate base (or weak base and its conjugate acid).
43. What volume of 0.50 M NaOH is required to neutralize 25.0 mL of 1.0 M HCl?
Answer
50.0 mL — Moles HCl = 0.025 mol; Volume NaOH = 0.025 / 0.50 = 0.050 L = 50.0 mL
44. Define the equivalence point in an acid-base titration.
Answer
It is the point at which moles of acid equal moles of base — complete neutralization.
45. A 0.10 M solution of a weak acid has a pH of 3. Calculate Ka.
Answer
Ka = 1.0 × 10^–5 — [H+] = 10^–3 = 0.001 M; Ka ≈ (0.001)^2 / 0.10 = 1.0 × 10^–5
46. Why does NH4+ lower the pH of a solution?
Answer
It acts as a weak acid, donating a proton to form NH3 and H+.
47. What is the pH at the halfway point in a weak acid–strong base titration?
Answer
pH = pKa — At this point, [HA] = [A–], so pH = pKa
48. What is the relationship between Ka and pKa?
Answer
pKa = –log(Ka) — The pKa is the negative base-10 logarithm of Ka.
49. What does it mean if a solution has [H+] = [OH–]?
Answer
The solution is neutral, and pH = 7 (at 25°C).
50. What is the pH of a solution if the pOH = 5.3?
Answer
8.7 — pH = 14 – 5.3 = 8.7