Acids and Bases Quiz
Questions: 16 · 10 minutes
1. Which species can act as a Lewis acid by accepting an electron pair?
Cl⁻
NH₃
OH⁻
BF₃
2. Equal-concentration solutions of HCl and acetic acid are prepared. Which statement is generally correct?
Acetic acid has the lower pH because weak acids retain more protons.
Both have the same pH because their concentrations are equal.
Their pH values cannot differ unless one solution contains a base.
HCl has the lower pH because it ionizes more completely in water.
3. Sodium fluoride dissolves in water, and F⁻ reacts slightly with water to form HF and OH⁻. How should the resulting solution be classified?
Basic, because F⁻ acts as the conjugate base of a weak acid.
Acidic, because every fluoride compound releases protons in water.
Neutral, because all salts produce solutions with pH 7.
Acidic, because OH⁻ is consumed when the salt dissolves.
4. Which comparison between acid strength and acid concentration is correct?
Strength describes ionization tendency, while concentration describes the amount dissolved per unit volume.
Diluting an acid necessarily changes it from strong to weak.
Acid strength and concentration are equivalent measures.
A strong acid must always be more concentrated than a weak acid.
5. Why can bicarbonate, HCO₃⁻, behave as an amphiprotic species?
It can release either a sodium ion or a chloride ion.
It can donate a proton to form CO₃²⁻ or accept a proton to form H₂CO₃.
It produces equal amounts of every acid and base in water.
It cannot participate in proton-transfer reactions.
6. A 0.010 M HCl solution is diluted tenfold with water. Assuming complete dissociation and ideal behavior, how does its pH change?
It increases from 2 to 3.
It decreases from 2 to 1.
It stays at 2 because HCl remains a strong acid.
It increases from 2 to 12.
7. Ammonia, NH₃, gains a proton during a Brønsted–Lowry reaction. Which species is produced?
The amide ion, NH₂⁻
Hydrazine, N₂H₄
The ammonium ion, NH₄⁺
The hydronium ion, H₃O⁺
8. Four weak acids are compared under the same conditions. Which one is the strongest acid?
The acid with the smallest molar mass
The acid with the largest Ka value
The acid with the smallest Ka value
The acid with the lowest concentration
9. What is the conjugate base of H₂PO₄⁻?
H₃PO₄
HPO₄²⁻
PO₄³⁻
H₃O⁺
10. Equal amounts in moles of HCl and NaOH react completely in aqueous solution, with neither reactant in excess. What are the principal products?
Na₂O and HCl
NaH and Cl₂
NaCl and H₂O
H₂ and NaClO
11. At 25°C, a solution has a hydronium-ion concentration of 1.0 × 10⁻³ M. What is its pH?
3
2
1
11
12. An acid–base indicator has a pKa close to 7. Over which approximate pH interval would its main color transition usually occur?
Around pH 0-2
Only at exactly pH 7
Around pH 10-12
Around pH 6-8
13. At 25°C, a strong monoprotic acid is titrated with a strong base. What is the approximate pH at the equivalence point under ideal conditions?
1
5
13
7
14. Which mixture is most likely to act as a buffer?
A strong acid mixed with excess strong acid
A strong base dissolved in pure water
A weak acid mixed with a significant amount of its conjugate base
Equal moles of a strong acid and strong base after complete neutralization
15. During the titration of a weak acid with a strong base, substantial amounts of both the weak acid and its conjugate base are present before the equivalence point. How does the solution behave?
It behaves as a strong acid whose pH cannot change.
It becomes neutral immediately after any base is added.
It acts as a buffer and resists moderate pH changes.
It behaves as pure water because the two species cancel completely.
16. Which statement best describes an Arrhenius acid in water?
It accepts an electron pair from another species.
It increases the concentration of hydronium ions in the solution.
It increases the concentration of hydroxide ions in the solution.
It must contain a metal ion that dissociates in water.