Chapter 19: Ionic Equilibria in Aqueous Systems

Flashcards covering buffers, titrations, and solubility equilibria concepts from Chapter 19.

What is the Henderson–Hasselbalch equation for a buffer containing a weak acid (HA) and its conjugate base (A−)?

pH = pKa + log([A−]/[HA])

What does buffering capacity measure in a buffer solution?

The ability of a buffer to resist changes in pH when small amounts of acid or base are added; capacity increases with concentration and is greatest when [acid] ≈ [base].

What is an acid–base buffer (in general terms)?

A solution that minimizes pH changes upon addition of acid or base, usually containing a weak acid and its conjugate base, or a weak base and its conjugate acid.

Name three common buffer pairs listed in the notes.

HC2H3O2/NaC2H3O2 (acetic acid/acetate), NH4Cl/NH3 (ammonium/ammonia), H2CO3/NaHCO3 (carbonic acid/bicarbonate)

Which of the following is NOT a buffer? A. HF/KF B. HNO2/NaNO2 C. HCl/NaCl D. NH4NO3/NH3 E. NaH2PO4/Na2HPO4

C. HCl/NaCl is not a buffer.

In preparing an HC2H3O2/CH3COO− buffer, what are the conjugate acid and conjugate base?

Conjugate acid: CH3COOH; conjugate base: CH3COO−.

What happens to [H3O+] and pH when a strong acid is added to a buffer?

The added H+ reacts with the conjugate base (A−) to form HA, causing only a small increase in [H3O+] and a small drop in pH.

What happens to [H3O+] and pH when a strong base is added to a buffer?

The acid (HA) reacts with OH− to form A− and H2O, causing a small decrease in [H3O+] and a small rise in pH.

What is the buffer range in terms of pH and pKa?

The usable pH range is approximately pKa ± 1.

Compute the pH of a buffer with 0.25 M HC2H3O2 and 0.35 M NaC2H3O2 given Ka = 1.8×10−5.

pH ≈ pKa + log([A−]/[HA]) = 4.75 + log(0.35/0.25) ≈ 4.89.

What is the general relationship between pH and pKa when [acid] ≈ [base]?

pH is approximately equal to pKa.

How is the ratio [base]/[acid] obtained from pH and pKa?

[base]/[acid] = 10^(pH − pKa).

For the blood buffer H2CO3/HCO3−, with pKa ≈ 6.35, what ratio gives pH = 7.40?

[HCO3−]/[H2CO3] ≈ 10^(7.40 − 6.35) ≈ 11.2.

What is the effect of adding a common ion on the solubility of a sparingly soluble salt?

It shifts the dissolution equilibrium to the left, decreasing solubility.

Define Ksp.

The solubility product constant; the equilibrium constant for the dissolution of a sparingly soluble solid: for AB(s) ⇌ A+(aq) + B−(aq), Ksp = [A+][B−].

How does Ksp relate to solubility?

Smaller Ksp generally indicates a less soluble salt under similar conditions.

What is the effect of a common ion on the solubility of Ag3PO4 in water (vs in 0.10 M AgNO3)?

In water, solubility is higher; in 0.10 M AgNO3 (common ion), solubility decreases dramatically due to Le Châtelier’s principle.

If Ag3PO4 has Ksp = 2.6×10−18, what is its molar solubility in pure water?

Let s be solubility: Ksp = (3s)^3(s) = 27s^4, so s ≈ (2.6×10−18/27)^(1/4) ≈ 1.8×10−5 M.

What is the molar solubility of Pb(IO3)2 given Ksp and solubility 4.0×10−5 M?

Pb(IO3)2 ⇌ Pb2+ + 2 IO3−; Ksp = [Pb2+][IO3−]^2 = s(2s)^2 = 4s^3; solving gives Ksp ≈ 2.56×10−13 when s ≈ 4.0×10−5 M.

What is the equivalence point in a titration?

The point at which chemically equivalent amounts of acid and base have reacted; the reaction is complete.

What is an endpoint in a titration?

The point at which a physical change (often a color change) indicates the equivalence point has been reached.

Name the three main types of titration covered in the notes.

Strong acid–strong base, Weak acid–strong base, Strong acid–weak base.

What is a pH curve?

A plot of pH versus the volume of titrant added.

What is the pH at the equivalence point for a strong acid–strong base titration?

pH ≈ 7 (neutral salt solution).

In a weak acid–strong base titration, what happens at the equivalence point?

The solution contains the conjugate base and is typically basic (pH > 7).

What is the key principle for choosing an indicator for a titration?

Choose an indicator whose pKa (color change range) is near the pH at the equivalence point so the color change occurs near equivalence.

What governs the color change of an acid–base indicator?

An indicator is a weak acid HIn that ⇌ H3O+ + In−; the ratio [HIn]/[In−] changes with pH, producing a color change.

Name two common indicators and their typical color changes at relevant ranges.

Examples: phenolphthalein (colorless in acid, pink in base; range ~8.2–10.0), methyl red (red in acid, yellow in base; range ~4.8–6.0).

What is the buffer range in terms of pH and pKa?

pH ≈ pKa ± 1.

What is the effect of adding acid to a buffer in terms of HA and A− concentrations?

[HA] increases slightly, [A−] decreases slightly; pH decreases slightly.

What is the effect of adding base to a buffer in terms of HA and A− concentrations?

[HA] decreases slightly, [A−] increases slightly; pH increases slightly.

What is the relationship between pH and pKa in a buffer when [HA] = [A−]?

pH = pKa.

What is the Ka and pKa relationship?

Ka is the acid dissociation constant; pKa = −log(Ka); smaller Ka means a weaker acid, larger pKa indicates weaker acid.

What is the formula for pH of a buffer using Ka and concentrations?

pH = pKa + log([A−]/[HA]); pKa = −log(Ka).

What is meant by spectator ions in buffer contexts?

Ions that do not participate in the acid–base equilibrium (e.g., Na+ or Cl− in many buffers).

What is the solubility product constant (Ksp) used for?

Describes the equilibrium between a sparingly soluble solid and its ions in solution; helps predict solubility.

Which statement correctly describes the common-ion effect on solubility?

Adding a common ion decreases solubility by shifting the dissolution equilibrium left.

What is the effect of the common ion on Ag3PO4 dissolving in 0.10 M AgNO3?

Solubility decreases substantially due to the common Ag+ ion.

What is the typical pH at equivalence for a strong acid–strong base titration and why?

Approximately 7, because the products are neutral salt and water, with no excess H+ or OH− in the solution.

Which buffer pair would best buffer around pH 7?

A pair with pKa close to 7 (for example, a conjugate acid–base system with pKa ≈ 7.0); in the notes, a near-7 buffer example is NaHSO3/Na2SO3 with pKa ≈ 7.19.

What is the general rule for choosing an indicator relative to the equivalence pH?

The indicator’s pKa (color change range) should bracket the pH at equivalence, yielding a visible endpoint near equivalence.

What is the typical pH range for blood buffer (H2CO3/HCO3−) at pH 7.4 and why is there more base than acid?

pH 7.4; ratio [HCO3−]/[H2CO3] ≈ 11.2, indicating more base (HCO3−) than acid (H2CO3) to stabilize pH during metabolism.

If a buffer has [HA] = [A−], what is the expected pH relative to pKa?

pH is approximately equal to pKa. If the ratio deviates, pH moves above or below pKa accordingly.