Rucete ✏ AP Chemistry In a Nutshell
13. Acids and Bases — Practice Questions
This chapter introduces acid–base theories, pH and pOH, buffer solutions, titration curves, and hydrolysis reactions.
(Multiple Choice — Click to Reveal Answer)
1. According to Arrhenius theory, what defines a base?
(A) Donates H+ ions
(B) Accepts OH– ions
(C) Produces OH– ions in solution
(D) Neutralizes acids completely
Answer
(C) — Arrhenius bases are defined as substances that increase OH– concentration in water.
2. Which of the following is a Brønsted-Lowry base?
(A) HCl
(B) NH3
(C) H2O
(D) CH3COOH
Answer
(B) — NH3 can accept a proton, making it a Brønsted-Lowry base.
3. What is the conjugate base of H2SO4 after losing one proton?
(A) HSO4–
(B) SO4^2–
(C) H2SO3
(D) H3SO4+
Answer
(A) — The conjugate base after losing one proton is HSO4–.
4. Which of the following is a strong base?
(A) NH3
(B) CH3NH2
(C) Ca(OH)2
(D) HCOOH
Answer
(C) — Ca(OH)2 is a strong base that dissociates fully in water.
5. Which ion is amphiprotic?
(A) Cl–
(B) NH4+
(C) H2PO4–
(D) Na+
Answer
(C) — H2PO4– can both donate and accept a proton.
6. What is the pH of a neutral solution at 25 °C?
(A) 6.34
(B) 7.00
(C) 6.77
(D) 5.00
Answer
(B) — At 25 °C, the pH of a neutral solution is exactly 7.00.
7. Which pair constitutes a buffer solution?
(A) HCl and NaCl
(B) HF and NaF
(C) NaOH and HCl
(D) NH4Cl and NaNO3
Answer
(B) — HF and NaF are a conjugate acid–base pair and form a buffer.
8. Which acid is considered strong?
(A) CH3COOH
(B) HF
(C) HNO3
(D) H2CO3
Answer
(C) — HNO3 is one of the six strong acids that dissociate completely.
9. The pOH of a 0.01 M NaOH solution is:
(A) 2
(B) 10
(C) 12
(D) 14
Answer
(A) — [OH–] = 0.01 M → pOH = –log(0.01) = 2
10. What happens when an acid reacts with a metal hydroxide?
(A) Only salt forms
(B) Only water forms
(C) Salt and water form
(D) No reaction occurs
Answer
(C) — Neutralization between an acid and base produces salt and water.
11. What is the name of H2SO3?
(A) Sulfuric acid
(B) Sulfurous acid
(C) Sulfous acid
(D) Sulfite acid
Answer
(B) — H2SO3 is sulfurous acid, derived from the sulfite ion (SO3^2–).
12. What is the pH of a 1.0 × 10^–3 M HCl solution?
(A) 3
(B) 11
(C) 7
(D) 1
Answer
(A) — pH = –log[H+] = –log(1.0 × 10^–3) = 3
13. Which base is weak?
(A) NaOH
(B) KOH
(C) NH3
(D) Ba(OH)2
Answer
(C) — NH3 is a weak base; it does not fully dissociate in water.
14. What species results from the complete dissociation of HNO3 in water?
(A) HNO2 and NO3–
(B) H+ and NO3–
(C) NO2– and H+
(D) HNO3 remains undissociated
Answer
(B) — HNO3 is a strong acid and dissociates completely into H+ and NO3–.
15. What is the pH of a solution with [OH–] = 1.0 × 10^–4 M?
(A) 4
(B) 10
(C) 7
(D) 6
Answer
(B) — pOH = 4 → pH = 14 – 4 = 10
16. What does the term “polyprotic acid” mean?
(A) An acid that forms multiple salts
(B) An acid with multiple hydrogen atoms
(C) An acid that can donate more than one proton
(D) An acid that cannot ionize
Answer
(C) — Polyprotic acids donate more than one proton (e.g., H3PO4, H2SO4).
17. Which pair is a conjugate acid–base pair?
(A) HCl and Cl–
(B) NaOH and Na+
(C) HNO3 and NO2–
(D) NH4+ and HCl
Answer
(A) — HCl and Cl– differ by one H+, making them a conjugate pair.
18. Which compound is a basic anhydride?
(A) SO2
(B) CO2
(C) CaO
(D) H2SO4
Answer
(C) — CaO is a basic anhydride; it forms Ca(OH)2 with water.
19. Which solution is likely to be acidic?
(A) NaF
(B) NH4Cl
(C) NaNO3
(D) Na2CO3
Answer
(B) — NH4+ is the conjugate acid of NH3, making the solution acidic.
20. The Ka of an acid is 1.8 × 10^–5. What is the pKa?
(A) 5.74
(B) 3.45
(C) 4.74
(D) 2.11
Answer
(C) — pKa = –log(1.8 × 10^–5) ≈ 4.74
21. What is the conjugate acid of NH3?
(A) NH2–
(B) NH4+
(C) NO3–
(D) N2
Answer
(B) — NH3 + H+ → NH4+
22. Which of the following has the lowest pH?
(A) 0.1 M HCl
(B) 0.1 M NaOH
(C) 0.1 M CH3COOH
(D) 0.1 M NaCl
Answer
(A) — HCl is a strong acid and will have the lowest pH.
23. Which acid is weakest among the following?
(A) HClO4
(B) HClO
(C) H2SO4
(D) HI
Answer
(B) — HClO is a weak acid; the others are strong acids.
24. Which of the following pairs will not result in a buffer solution?
(A) HNO2 and NaNO2
(B) NH3 and NH4Cl
(C) HCl and NaCl
(D) HF and NaF
Answer
(C) — HCl is a strong acid and cannot form a buffer with its salt.
25. Which is a property of a buffer solution?
(A) It changes pH rapidly when base is added
(B) It contains only a strong acid
(C) It resists changes in pH
(D) It is always basic
Answer
(C) — Buffer solutions resist changes in pH upon addition of small amounts of acid or base.
26. What is the pH of a buffer that contains 0.250 M acetic acid (Ka = 1.8 × 10^–5) and 0.150 M sodium acetate?
(A) 4.43
(B) 4.57
(C) 4.74
(D) 5.00
Answer
(B) — Using the Henderson-Hasselbalch equation: pH = 4.74 + log(0.150/0.250) = 4.57
27. Which species cannot exist together in significant amounts in an aqueous solution?
(A) Na+ and OH–
(B) NH4+ and NH3
(C) Cl– and HCl
(D) CO3^2– and H2CO3
Answer
(D) — These differ by two protons and are not a conjugate acid–base pair; they do not coexist in buffer systems.
28. Which of the following will produce a solution with the highest pH?
(A) 0.10 M NH4Cl
(B) 0.10 M NaCl
(C) 0.10 M NaHCO3
(D) 0.10 M HNO3
Answer
(C) — NaHCO3 is basic due to the hydrolysis of HCO3–, producing OH– ions.
29. Which of the following is the best combination to prepare a buffer at pH ≈ 10?
(A) CH3COOH / CH3COONa
(B) HNO2 / NaNO2
(C) NH3 / NH4Cl
(D) HCl / NaCl
Answer
(C) — NH3 has a pKa around 9.25 and is ideal for preparing a buffer near pH 10.
30. What is the pH at the equivalence point when 0.100 M HF is titrated with 0.100 M NaOH?
(A) 7.00
(B) < 7
(C) > 7
(D) Cannot be determined
Answer
(C) — The conjugate base F– formed at equivalence hydrolyzes to produce OH–, so pH > 7.
31. A weak base has Kb = 1.8 × 10^–5. What is the pKa of its conjugate acid?
(A) 4.74
(B) 5.74
(C) 9.74
(D) 10.74
Answer
(D) — pKa = 14 – pKb = 14 – (–log(1.8 × 10^–5)) = 10.74
32. In a titration curve of a weak acid with a strong base, where is the buffer region?
(A) Before any base is added
(B) Around the equivalence point
(C) Just before the equivalence point
(D) After excess base is added
Answer
(C) — The buffer region occurs before the equivalence point where both weak acid and its conjugate base are present.
33. What is the [OH–] in a solution with pH = 3.00?
(A) 1.0 × 10^–11 M
(B) 1.0 × 10^–3 M
(C) 1.0 × 10^–7 M
(D) 1.0 × 10^–6 M
Answer
(A) — pOH = 14 – 3 = 11 → [OH–] = 10^–11 M
34. Which of the following acids would be strongest based on structure and inductive effect?
(A) CH3COOH
(B) CCl3COOH
(C) ClCH2COOH
(D) CH3CH2COOH
Answer
(B) — CCl3COOH is the strongest because the highly electronegative Cl atoms withdraw electrons, weakening the O–H bond.
35. What is the role of a primary standard in titration?
(A) To neutralize weak bases
(B) To serve as the analyte
(C) To standardize titrant solutions
(D) To control temperature
Answer
(C) — A primary standard is a pure substance used to determine the exact concentration of titrants.
36. Write the balanced chemical equation for the reaction between HNO3 and NaOH.
Answer
HNO3 + NaOH → NaNO3 + H2O — This is a typical neutralization reaction between a strong acid and a strong base.
37. What is the conjugate base of HCOOH (formic acid)?
Answer
HCOO– — The conjugate base forms when HCOOH loses a proton (H+).
38. A 0.100 M solution of NH3 has a pH of 11.2. What is the [OH–]?
Answer
6.3 × 10^–4 M — pOH = 14 – 11.2 = 2.8 → [OH–] = 10^–2.8 = 6.3 × 10^–4 M
39. What does it mean when Ka × Kb = Kw?
Answer
It shows the relationship between a conjugate acid–base pair. — The stronger the acid (larger Ka), the weaker its conjugate base (smaller Kb), and vice versa.
40. What is the effect on pH when a buffer is diluted with water?
Answer
Minimal change — The pH of a buffer is determined by the ratio of acid to base, which remains constant upon dilution.
41. Identify a conjugate acid–base pair from this reaction: NH3 + H2O ⇌ NH4+ + OH–
Answer
NH3 and NH4+ — NH3 is the base, and NH4+ is its conjugate acid.
42. What makes HClO4 a stronger acid than HClO?
Answer
More electronegative oxygen atoms withdraw electron density — This weakens the O–H bond, making it easier to lose a proton.
43. What is the definition of a Lewis base?
Answer
An electron pair donor — Lewis bases donate electron pairs to form covalent bonds.
44. What is the pH of a solution with [H+] = 3.2 × 10^–5 M?
Answer
4.49 — pH = –log(3.2 × 10^–5) ≈ 4.49
45. What type of compound is SO3 and what does it form when dissolved in water?
Answer
Acid anhydride; forms H2SO4 — SO3 reacts with water to form sulfuric acid.
46. How does electronegativity affect acid strength?
Answer
Greater electronegativity of the central atom increases acid strength — It pulls electron density away from the O–H bond.
47. Why is HCl considered a strong acid?
Answer
It dissociates completely in water — Almost 100% of HCl molecules ionize to form H+ and Cl–.
48. Give one example of an amphiprotic substance.
Answer
HCO3– — It can act as both a proton donor and acceptor.
49. What is the purpose of a buffer in a biological system?
Answer
To maintain a stable pH — Buffers resist sudden changes in pH from added acids or bases.
50. What happens to the pH of pure water when temperature increases?
Answer
pH decreases slightly — Kw increases with temperature, so [H+] increases, lowering the pH even though the solution remains neutral.