Acids, Bases & Strength
In NH₃ + H₂O ⇌ NH₄⁺ + OH⁻, identify the Brønsted–Lowry acid and base, and both conjugate pairs.
H₂O is the acid; NH₃ is the base. Conjugate pairs: NH₄⁺/NH₃ and H₂O/OH⁻.
Calculate the pH and pOH of 0.0200 M KOH.
pOH = −log(0.0200) = 1.699. pH = 12.301.
Which pair forms a buffer: HNO₃/KNO₃, HCN/KCN, or NaOH/NaCl? Explain.
HCN/KCN, because it supplies a weak acid and its conjugate base.
At equivalence, is the solution acidic, basic, or neutral for each titration?
Which process loses electrons? Which gains electrons? Is the oxidizing agent oxidized or reduced?
Oxidation loses electrons; reduction gains electrons. The oxidizing agent is reduced.
H₂PO₄⁻ can either gain or lose a proton. Write its conjugate acid and conjugate base. What term describes this ability?
Conjugate acid: H₃PO₄. Conjugate base: HPO₄²⁻. H₂PO₄⁻ is amphiprotic.
Calculate the pH of 0.0040 M Sr(OH)₂. Assume complete dissolution.
[OH⁻] = 0.0080 M. pOH = 2.10, so pH = 11.90.
KNO₂ is added to an HNO₂ solution. What happens to HNO₂ ionization, [H⁺], and pH?
Added NO₂⁻ shifts HNO₂ ⇌ H⁺ + NO₂⁻ left. Ionization decreases, [H⁺] decreases, and pH increases.
A 20.0 mL sample of 0.150 M weak base is titrated with 0.100 M HCl. What HCl volume reaches equivalence? What is the total solution volume at half-equivalence?
Initial base = 0.00300 mol. Equivalence requires 30.0 mL HCl. Half-equivalence requires 15.0 mL HCl, giving 35.0 mL total volume.
Find the oxidation number of:
Cl = +1; N = +5; O = −1.
Which acid is stronger in each pair? Explain why.
Is 0.100 M NaCH₃COO acidic or basic? Calculate its pH. Kₐ(CH₃COOH) = 1.8 × 10⁻⁵.
Basic. pH = 8.87.
Mix 30.0 mL of 0.100 M NH₄Cl with 10.0 mL of 0.100 M NaOH. Does a buffer form? Give the final moles of NH₄⁺ and NH₃.
OH⁻ + NH₄⁺ → NH₃ + H₂O. Initially: 3.00 mmol NH₄⁺ and 1.00 mmol OH⁻. Finally: 2.00 mmol NH₄⁺ and 1.00 mmol NH₃. A buffer forms.
Methyl orange has a red protonated form and a yellow deprotonated form. Which form dominates well below its pKₐ? At pH = pKₐ, what is the ratio of the two forms?
The red, protonated form dominates below pKₐ. At pH = pKₐ, the protonated and deprotonated forms have equal concentrations: 1:1.
For Zn + 2Ag⁺ → Zn²⁺ + 2Ag, identify the species oxidized, species reduced, oxidizing agent, and reducing agent.
Zn is oxidized and is the reducing agent. Ag⁺ is reduced and is the oxidizing agent.
Arrange equal-concentration solutions of HNO₃, HClO, NH₃, KOH, and Ba(OH)₂ by increasing pH. Assume complete dissolution.
HNO₃ < HClO < NH₃ < KOH < Ba(OH)₂. Ba(OH)₂ supplies two OH⁻ per formula unit.
Calculate the pH and percent ionization of 0.100 M HA with Kₐ = 1.0 × 10⁻³. Check the small-x approximation.
The approximation predicts 10% ionization, so it fails the 5% check. Solve x²/(0.100 − x) = 0.00100. [H⁺] = 0.00951 M. pH = 2.02; ionization = 9.51%.
A buffer contains 0.100 mol CH₃COOH and 0.100 mol CH₃COO⁻. Calculate its pH after adding 0.0200 mol HCl. Kₐ = 1.8 × 10⁻⁵. Assume negligible volume change.
HCl converts acetate into acetic acid. Final amounts: 0.0800 mol acetate and 0.120 mol acid. pH = 4.745 + log(0.0800/0.120) = 4.57.
A 25.0 mL sample of 0.100 M HCOOH is titrated with 0.100 M NaOH. Calculate the pH after adding 10.0 mL NaOH. Kₐ(HCOOH) = 1.8 × 10⁻⁴.
Initially 2.50 mmol acid. Added OH⁻ = 1.00 mmol. Finally: 1.50 mmol HCOOH and 1.00 mmol HCOO⁻. pH = 3.745 + log(1.00/1.50) = 3.57.
Write the balanced half-reaction for NO₃⁻ being reduced to NO in acidic solution.
NO₃⁻ + 4H⁺ + 3e⁻ → NO + 2H₂O.
For a triprotic acid, rank Kₐ₁, Kₐ₂, and Kₐ₃, then rank their pKₐ values. During titration with a strong base, where does pH equal each pKₐ?
Kₐ₁ > Kₐ₂ > Kₐ₃; pKₐ₁ < pKₐ₂ < pKₐ₃. The pH equals each pKₐ at its corresponding half-equivalence point, halfway through that proton’s neutralization.
Calculate the pH of a 0.200 M NH₄Cl solution at 25°C. Is the solution acidic, basic, or neutral?
Given: Kᵦ(NH₃) = 1.8 × 10⁻⁵.
NH₄⁺ is a weak acid; Cl⁻ is neutral.
ph= 4.98
Two 1.00 L buffers have equal acid-to-conjugate-base ratios. Buffer A contains 0.010 mol of each component; Buffer B contains 0.100 mol of each. Compare their initial pH values and their resistance to added HCl.
Their initial pH values are approximately equal, because their ratios are equal. Buffer B has greater capacity and experiences a smaller pH change for the same HCl addition.
A 20.0 mL sample of 0.150 M weak base B is titrated with 0.100 M HCl. Kᵦ(B) = 1.0 × 10⁻⁵. Calculate the pH:
Balance the following reaction in basic solution and identify the oxidizing agent:
Cr(OH)₃ + ClO⁻ → CrO₄²⁻ + Cl⁻
2Cr(OH)₃ + 3ClO⁻ + 4OH⁻ → 2CrO₄²⁻ + 3Cl⁻ + 5H₂O.
ClO⁻ is the oxidizing agent: chlorine is reduced from +1 to −1.