Common Ka and Kb values

Look up the pKa of a weak acid or the pKb of a weak base at a glance. Values are aqueous, 25 °C, ionic strength I = 0 (thermodynamic constants) unless a row says otherwise. Cross-check high-precision work against Harris QCA Appendix G or the CRC Handbook — pKa’s shift measurably with temperature and ionic strength.

Weak acids — pKa order

HF pKa varies across sources. Harris and IUPAC critical compilations give 3.17; some older textbooks list 3.20 (Ka 6.3×10⁻⁴). Both fall inside 3.14–3.20 depending on ionic-strength correction, so quote with a hedge (≈ 3.17) unless your class picked a value.

Polyprotic acids

For H₃PO₄, the second dissociation (pKa₂ ≈ 7.20) is what makes phosphate the physiological buffer — pKa near blood pH 7.4 gives maximum buffer capacity there.

Weak bases — pKb order

For a base B, Kb = [BH⁺][OH⁻] / [B]. At 25 °C, pKa(BH⁺) + pKb(B) = pKw = 14.00, so you can convert either way from the conjugate-acid pKa.

Pattern: aliphatic amines are stronger bases than NH₃; aromatic amines (pyridine, aniline) are much weaker because the nitrogen lone pair delocalizes into the ring.

Biological-buffer favorites

Tris pKa is unusually temperature-dependent (dpKa/dT ≈ −0.028 per °C). A Tris buffer titrated to pH 8.0 on the bench drifts to about pH 7.5 in a 37 °C incubator — a real problem for cell culture and enzyme assays.

Ka / Kb and pKa / pKb — the arithmetic

• pKa = −log₁₀(Ka). Smaller pKa → stronger acid.
• pKb = −log₁₀(Kb). Smaller pKb → stronger base.
• For any conjugate pair, Ka · Kb = Kw = 1.0×10⁻¹⁴ (25 °C), so pKa + pKb = 14.00.
• Kw itself is temperature-dependent: Kw ≈ 1.0×10⁻¹⁴ at 25 °C, ≈ 2.4×10⁻¹⁴ at 37 °C (Harris Appendix G) — which is why “neutral” at body temperature is pH ≈ 6.81, not 7.00.

Worked example — Henderson–Hasselbalch

Prepare an acetate buffer at pH 5.00 from 0.10 M acetic acid and 0.10 M sodium acetate. Use pKa(acetic) = 4.76.

Henderson–Hasselbalch: pH = pKa + log([A⁻] / [HA]).

Solve for the ratio:

• log([A⁻] / [HA]) = pH − pKa = 5.00 − 4.76 = 0.24
• [A⁻] / [HA] = 10^0.24 = 1.74

So mix 1.74 volumes of 0.10 M sodium acetate with 1.00 volume of 0.10 M acetic acid. In a 100 mL total prep: ≈ 63.5 mL acetate + 36.5 mL acid. Check the final pH with a meter — Henderson–Hasselbalch neglects activity corrections and is exact only at I = 0.

Values shift with T and I

The tabulated pKa’s in Harris Appendix G are thermodynamic — extrapolated to zero ionic strength (I = 0) and reported at 25 °C. Real solutions differ:

• Temperature. Most pKa’s drift by roughly 0.01–0.03 units per °C. Amine bases (Tris, ammonia) are on the high end of that drift.
• Ionic strength. Activity coefficients suppress apparent [H⁺] and [conjugate species]. “Conditional” or “apparent” pKa’s tabulated at I = 0.1 M can differ from the I = 0 value by 0.1–0.3 units. Harris flags this whenever a chapter uses apparent constants.
• Solvent. Add methanol, DMSO, or other cosolvent and every constant on this page changes — sometimes by several pKa units. Never carry aqueous pKa’s into non-aqueous work.

Common student mistakes

1. Using pKa where the problem needs pKb (or vice versa). Convert with pKa + pKb = 14.00 at 25 °C.
2. Quoting HF to three significant figures across sources. Different compilations report 3.14, 3.17, and 3.20 — hedge unless your course fixed a value.
3. Ignoring polyprotic behavior. Phosphate has three pKa’s; carbonate has two. Which species dominates depends on pH.
4. Confusing thermodynamic and apparent constants. Bench pH readings correspond to apparent, not thermodynamic, activities.
5. Applying Henderson–Hasselbalch far from pKa. The approximation is best within ±1 pH unit of pKa; outside that window, use the full ICE-table treatment.