Rucete ✏ AP Chemistry In a Nutshell
13. Acids and Bases — Practice Questions 2
This chapter introduces acid–base theories, strength comparisons, buffer calculations, titration curve interpretation, and pH/pOH problem-solving.
(Multiple Choice — Click to Reveal Answer)
1. According to the Brønsted–Lowry definition, which of the following is a base?
(A) HCl
(B) H2O
(C) NH4+
(D) CH3COOH
Answer
(B) — Water can accept a proton, making it a Brønsted–Lowry base.
2. What is the conjugate acid of NH2–?
(A) NH3
(B) NH4+
(C) NH2+
(D) N2
Answer
(A) — Adding a proton to NH2– gives NH3, its conjugate acid.
3. Which of the following is the strongest base?
(A) Cl–
(B) NO3–
(C) NH2–
(D) H2O
Answer
(C) — NH2– is a very strong base due to its ability to readily accept protons.
4. Which compound acts as an acid in the following reaction: HCOOH + NH3 → HCOO– + NH4+?
(A) HCOOH
(B) NH3
(C) HCOO–
(D) NH4+
Answer
(A) — HCOOH donates a proton, acting as the acid.
5. What is the pOH of a 0.001 M NaOH solution?
(A) 3
(B) 7
(C) 11
(D) 14
Answer
(A) — pOH = –log[OH–] = –log(0.001) = 3
6. What is the name of the base that reacts with HCl to form NaCl and H2O?
(A) Na2CO3
(B) NaOH
(C) NH3
(D) CH3COONa
Answer
(B) — NaOH is a strong base that neutralizes HCl to form NaCl and water.
7. Which of the following is a Lewis acid?
(A) NH3
(B) H2O
(C) BF3
(D) OH–
Answer
(C) — BF3 can accept an electron pair, making it a Lewis acid.
8. What is the pKa of a weak acid with Ka = 4.0 × 10^–6?
(A) 3.40
(B) 5.40
(C) 6.40
(D) 7.40
Answer
(B) — pKa = –log(4.0 × 10^–6) ≈ 5.40
9. Which of the following salts will make a solution basic in water?
(A) NH4Cl
(B) NaNO3
(C) CH3COONa
(D) KBr
Answer
(C) — CH3COONa contains acetate, a weak base, making the solution basic.
10. Which of the following acids has the weakest conjugate base?
(A) HClO4
(B) H2CO3
(C) HF
(D) CH3COOH
Answer
(A) — Strong acids have very weak conjugate bases; HClO4 is very strong.
11. Which of the following substances can act as both an acid and a base?
(A) HCl
(B) NH4+
(C) H2O
(D) NaCl
Answer
(C) — Water is amphiprotic and can act as both a proton donor and acceptor.
12. What happens to pH when a small amount of strong base is added to a buffer?
(A) pH drops sharply
(B) pH stays the same
(C) pH increases slightly
(D) pH decreases slightly
Answer
(C) — Buffers resist pH change, but a small rise may occur.
13. The Ka of formic acid is 1.8 × 10^–4. What is the pH of a 0.10 M solution?
(A) 2.37
(B) 3.25
(C) 4.74
(D) 5.00
Answer
(B) — Use ICE table and Ka = x²/(0.10) → x ≈ 5.5 × 10^–3 → pH = –log(x) ≈ 3.25
14. What is the conjugate base of HSO4–?
(A) SO4^2–
(B) H2SO4
(C) HSO3–
(D) HSO4^2–
Answer
(A) — The conjugate base of HSO4– is SO4^2– (after losing a proton).
15. Which is true about strong acids?
(A) They have large Ka values and small pKa values
(B) They have low conductivity
(C) They have high pH
(D) They form strong conjugate bases
Answer
(A) — Strong acids dissociate completely and have very large Ka (and low pKa).
16. What is the hydroxide ion concentration in a solution with pH = 9.00?
(A) 1.0 × 10^–5 M
(B) 1.0 × 10^–9 M
(C) 1.0 × 10^–3 M
(D) 1.0 × 10^–10 M
Answer
(A) — pOH = 14 – 9 = 5 → [OH–] = 10^–5 M
17. What is the best explanation for why HF is a weak acid while HCl is strong?
(A) HF is larger in size
(B) HF has stronger H–F bond
(C) HCl is less polar
(D) HCl cannot dissociate
Answer
(B) — The strong H–F bond prevents full dissociation, making HF weak.
18. A solution has [H+] = 1.0 × 10^–8 M. What is its pH?
(A) 6.0
(B) 8.0
(C) 7.0
(D) 5.0
Answer
(B) — pH = –log[H+] = –log(1.0 × 10^–8) = 8.0
19. Which of the following is a weak acid?
(A) HNO3
(B) HCl
(C) H2SO4
(D) H2CO3
Answer
(D) — Carbonic acid (H2CO3) only partially dissociates, making it weak.
20. Which statement is true for the autoionization of water?
(A) H2O → H+ + OH–
(B) H2O + H2O ⇌ H3O+ + OH–
(C) H2O ⇌ H2 + O2
(D) H2O + OH– ⇌ H3O+
Answer
(B) — Water autoionizes to form hydronium and hydroxide ions.
21. What is the pH of a solution with [OH–] = 2.0 × 10^–4 M?
(A) 10.3
(B) 3.7
(C) 7.0
(D) 4.7
Answer
(A) — pOH = –log(2.0 × 10^–4) ≈ 3.7 → pH = 14 – 3.7 = 10.3
22. What is the role of water in the following reaction? NH3 + H2O ⇌ NH4+ + OH–
(A) Strong acid
(B) Weak base
(C) Acid
(D) Spectator ion
Answer
(C) — Water donates a proton to NH3, acting as the acid.
23. If a solution has a pH = 5 and another has pH = 3, how much more acidic is the second?
(A) 5 times
(B) 10 times
(C) 20 times
(D) 100 times
Answer
(D) — Each pH unit represents a tenfold change → 10^2 = 100 times more acidic.
24. Which species acts as the conjugate acid of OH– in aqueous solution?
(A) O2–
(B) H2O
(C) H3O+
(D) H2O2
Answer
(B) — The conjugate acid of OH– is H2O, which forms when OH– accepts a proton.
25. Which is more acidic: a 0.1 M HCl or a 0.1 M CH3COOH solution?
(A) CH3COOH
(B) HCl
(C) Both are equal
(D) Cannot be determined
Answer
(B) — HCl is a strong acid and fully dissociates, giving a lower pH.
26. Which solution has the highest [H+]?
(A) pH = 6.0
(B) pH = 4.0
(C) pH = 7.0
(D) pH = 5.0
Answer
(B) — Lower pH means higher [H+]; pH 4 has the highest [H+].
27. Which of the following would most likely result in a pH above 7 when dissolved in water?
(A) KBr
(B) NH4Cl
(C) NaF
(D) HNO3
Answer
(C) — F– (from NaF) hydrolyzes to form OH–, raising the pH.
28. A weak acid and its conjugate base are mixed in equal concentrations. What is the resulting pH relative to pKa?
(A) pH < pKa
(B) pH > pKa
(C) pH = pKa
(D) pH cannot be determined
Answer
(C) — When [HA] = [A–], pH = pKa (Henderson-Hasselbalch equation).
29. Which salt will hydrolyze in water to produce a basic solution?
(A) KCl
(B) NH4NO3
(C) Na2CO3
(D) Ca(NO3)2
Answer
(C) — CO3^2– hydrolyzes to form OH–, making the solution basic.
30. The pH at the halfway point of a titration of a weak acid with a strong base is:
(A) equal to 7
(B) equal to the pKa of the acid
(C) dependent on the concentration
(D) always acidic
Answer
(B) — At halfway to equivalence, [HA] = [A–] → pH = pKa
31. A strong base is titrated with a strong acid. At equivalence, what is the pH of the solution?
(A) Below 7
(B) 7
(C) Above 7
(D) Depends on the base
Answer
(B) — Strong acid + strong base → neutral solution at pH = 7
32. What happens if equal volumes of 0.10 M HCl and 0.10 M NH3 are mixed?
(A) The pH becomes strongly basic
(B) A buffer forms
(C) The pH becomes strongly acidic
(D) A salt and water form without buffering
Answer
(B) — NH4+ and NH3 form a conjugate pair → buffer system
33. What volume of 0.10 M NaOH is needed to neutralize 25.0 mL of 0.10 M HCl?
(A) 12.5 mL
(B) 25.0 mL
(C) 50.0 mL
(D) 10.0 mL
Answer
(B) — Equal molar concentrations and volume → 25.0 mL required
34. What is the main species present in a 0.1 M solution of H2SO4 after the first dissociation?
(A) SO4^2–
(B) HSO4–
(C) H2SO4
(D) OH–
Answer
(B) — The first dissociation is complete → HSO4– is the dominant species before full second dissociation
35. Which of the following is best for preparing a buffer solution with pH ≈ 9?
(A) CH3COOH and CH3COONa
(B) H2CO3 and NaHCO3
(C) NH3 and NH4Cl
(D) HCl and NaCl
Answer
(C) — NH3/NH4+ has a pKa around 9.25, ideal for buffering near pH 9
36. What is the pH of a 1.0 × 10^–7 M HCl solution, considering water’s autoionization?
Answer
6.98 — Must account for both HCl and H2O contribution: total [H+] slightly above 1.0 × 10^–7 M, so pH ≈ 6.98
37. Write the balanced equation for the neutralization of sulfuric acid (H2SO4) with sodium hydroxide (NaOH).
Answer
H2SO4 + 2NaOH → Na2SO4 + 2H2O — H2SO4 provides two protons; NaOH provides two hydroxides.
38. What is the primary difference between a strong acid and a weak acid at the molecular level?
Answer
Strong acids dissociate completely in solution, while weak acids only partially dissociate.
39. Define a monoprotic acid and give one example.
Answer
A monoprotic acid can donate only one proton per molecule, e.g., HCl.
40. Explain why the pH of rainwater is naturally slightly acidic.
Answer
CO2 dissolves in water forming carbonic acid (H2CO3), which lowers the pH slightly.
41. Calculate the [H+] in a solution with pH = 2.70.
Answer
2.0 × 10^–3 M — [H+] = 10^–2.70 ≈ 2.0 × 10^–3 M
42. What is the role of NH4+ in a NH4Cl solution regarding acidity?
Answer
NH4+ acts as a weak acid, releasing H+ and lowering the pH.
43. Identify the conjugate base of H2CO3.
Answer
HCO3– — It forms when carbonic acid donates one proton.
44. Describe the effect of temperature increase on the Kw of water.
Answer
Kw increases with temperature, resulting in higher [H+] and [OH–].
45. What is the pOH of a 1.0 × 10^–9 M HCl solution?
Answer
5.0 — [H+] from HCl ≈ 1.0 × 10^–9 M, [OH–] = 10^–5 M → pOH = 5.0
46. Why is HF considered a weak acid despite being a halogen acid like HCl?
Answer
HF has a very strong H–F bond that resists dissociation, unlike HCl.
47. Predict the pH of a solution made by mixing equal volumes of 0.1 M NaOH and 0.1 M HNO3.
Answer
7.0 — Strong acid and base completely neutralize each other → neutral solution.
48. How does dilution affect the strength of a weak acid?
Answer
Dilution doesn’t change the intrinsic strength (Ka), but lowers the concentration and increases dissociation.
49. Write the net ionic equation for the neutralization of CH3COOH with NaOH.
Answer
CH3COOH + OH– → CH3COO– + H2O — Net ionic form of weak acid + strong base.
50. What is the typical pH range of an effective buffer system?
Answer
pKa ± 1 — Buffers are effective within one pH unit of their acid’s pKa value.