Quiz · Hard
pH and pOH, weak acid equilibria, buffers with Henderson-Hasselbalch, titrations and acid-base definitions. Each numeric problem is worked step by step.
Which of these is NOT a strong acid?
HCl
HNO₃
H₂SO₄
HF
Hydrofluoric acid is weak: the strong H–F bond and the small, highly solvated F⁻ ion limit how much it ionizes, even though HF is very dangerous. HCl, HBr, HI, HNO₃, H₂SO₄ and HClO₄ are the common strong acids.
Which species acts as a Lewis acid?
BF₃
NH₃
OH⁻
H₂O
A Lewis acid accepts an electron pair. Boron in BF₃ has only six valence electrons and an empty orbital, so it readily accepts a lone pair (for example from NH₃). The other three are electron-pair donors (Lewis bases).
A weak acid is titrated with a strong base. The pH at the equivalence point is:
Exactly 7
Greater than 7
Less than 7
Equal to the pKa
At equivalence the flask holds the conjugate base (for example acetate), which reacts with water to make OH⁻, so the pH is above 7. pH equals pKa at the half-equivalence point, not the equivalence point.
What is the pH of 0.010 M HCl?
1.00
2.00
0.010
12.00
HCl is a strong acid, so . . 12.00 would be the pH of 0.010 M NaOH.
What is the pH of 0.0010 M NaOH at 25 °C?
11.00
3.00
10.00
7.00
NaOH is a strong base: , so and . Answering 3.00 means you stopped at pOH.
A solution has pH 4.50. What is [H+]?
. is for this solution (pOH 9.50).
How does [H+] at pH 3 compare with [H+] at pH 5?
2 times greater
10 times greater
100 times greater
100 times smaller
pH is a base-10 log scale. A difference of 2 pH units is a factor of . Lower pH means more acidic, so pH 3 has the higher [H+].
In , what role does water play?
Brønsted-Lowry base
Brønsted-Lowry acid
Conjugate base
Spectator
Water donates a proton to ammonia, leaving OH⁻, so here it is a proton donor (acid). Water is amphoteric: with HCl it acts as a base instead.
What is the conjugate base of H₂PO₄⁻?
H₃PO₄
HPO₄²⁻
PO₄³⁻
H₂PO₄
A conjugate base is the acid minus one H+, with the charge lowered by 1: H₂PO₄⁻ → HPO₄²⁻. H₃PO₄ is its conjugate acid, and PO₄³⁻ is two protons away.
What is the pH of 0.10 M acetic acid? ()
1.00
4.74
5.74
2.87
For a weak acid, , so and . 1.00 treats acetic acid as strong, and 4.74 is its pKa.
A buffer contains 0.20 M acetic acid and 0.20 M sodium acetate (pKa = 4.74). What is its pH?
4.74
2.87
7.00
9.26
Henderson-Hasselbalch: . When acid and conjugate base are equal, pH equals pKa. 9.26 is the pKb of acetate.
A buffer has [acetate] = 10 × [acetic acid]. With pKa = 4.74, what is its pH?
3.74
4.84
14.74
5.74
. More base than acid raises the pH above the pKa. 3.74 comes from flipping the ratio.
What volume of 0.125 M NaOH neutralizes 25.0 mL of 0.100 M HCl?
31.3 mL
25.0 mL
12.5 mL
20.0 mL
Moles HCl = 0.0250 L × 0.100 M = 0.00250 mol. With a 1 : 1 ratio, NaOH volume = 0.00250 / 0.125 = 0.0200 L = 20.0 mL. 31.3 mL comes from swapping the two concentrations.
A solution of ammonium chloride (NH₄Cl) in water is:
Neutral, because it is a salt
Slightly acidic
Slightly basic
Strongly basic
NH₄+ is the conjugate acid of the weak base NH₃ and donates protons to water. Cl⁻, the conjugate base of a strong acid, does not react. Not every salt is neutral.
A solution at 25 °C has pOH 3.2. What is its pH?
3.2
10.8
6.3
17.2
At 25 °C, pH + pOH = 14.00, so pH = 14.0 − 3.2 = 10.8. The solution is basic, which matches a low pOH.
Which indicator best suits titrating acetic acid with NaOH?
Methyl orange (changes around pH 3.1–4.4)
Methyl red (changes around pH 4.4–6.2)
Phenolphthalein (changes around pH 8.2–10)
No indicator works for this titration
The equivalence point of a weak acid with a strong base is above 7 (about 8.7 here), so you need an indicator that changes in the basic range. Methyl orange would change far too early.
Pure water at 50 °C has a pH of about 6.6, so it is acidic.
True
False
False. Kw grows with temperature, so both [H+] and [OH⁻] rise. They are still equal, and that equality is what defines neutral. pH 7 means neutral only at 25 °C.
