Properties
High-Yield Summary
- Kw = [H₃O⁺][OH⁻] = 1.0×10⁻¹⁴ at 25°C (water autoionization); [H₃O⁺]=[OH⁻]=1.0×10⁻⁷ M in pure water. Fixed at a given T, changes with T.
- pH = −log[H⁺], pOH = −log[OH⁻], pH + pOH = 14 at 298 K. pH<7 acidic, pH>7 basic, pH=7 neutral.
- Strong acids/bases dissociate completely — memorize the short lists (HCl, HBr, HI, H₂SO₄, HNO₃, HClO₄, HClO₃; LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)₂, Sr(OH)₂, Ba(OH)₂). Everything else is weak.
- Ka = [H₃O⁺][A⁻]/[HA] for weak acids, Kb = [B⁺][OH⁻]/[BOH] for weak bases. Ka×Kb = Kw for any conjugate pair — one strong forces the other weak.
- Salt hydrolysis after neutralization: strong+strong→neutral (pH7); strong acid+weak base→acidic; weak acid+strong base→basic; weak+weak→depends on relative strengths.
Key Terms
- Autoionization of water
- H₂O + H₂O ⇌ H₃O⁺ + OH⁻ — water reacting with itself, the basis of Kw.
- Kw
- Water dissociation constant = [H₃O⁺][OH⁻] = 1.0×10⁻¹⁴ at 25°C.
- Ka / Kb
- Equilibrium (dissociation) constants for weak acids / weak bases; smaller value = weaker, less dissociation.
- Induction
- Electronegative atoms near an acidic proton pull electron density from the bond, weakening it and increasing acid strength.
- Neutralization
- Acid + Base → Salt + Water.
pH / pOH
pH = −log[H⁺] pOH = −log[OH⁻] pH + pOH = 14 (298 K)
- [H⁺] = hydrogen (hydronium) ion concentration, M
- Fast estimate without a calculator: for [H⁺] = n×10⁻ᵐ, p-value ≈ m − 0.n. Ex: 4×10⁻⁵ M → pH ≈ 5−0.4 = 4.6 (actual 4.40).
- Worked example: 0.010 M HCl (strong acid, full dissociation) → [H⁺]=1.0×10⁻² M → pH=2.00.
Ka × Kb = Kw
Ka(acid) × Kb(conjugate base) = Kw
- Kw = 1.0×10⁻¹⁴ at 25°C
- Worked example: acetic acid Ka=1.8×10⁻⁵ → Kb(acetate)=Kw/Ka=(1.0×10⁻¹⁴)/(1.8×10⁻⁵)≈5.6×10⁻¹⁰ (very weak conjugate base).
Neutralization Outcomes by Strength Pairing
| Acid + Base | Resulting pH |
|---|---|
| Strong + strong (HCl + NaOH) | Neutral (pH 7) — Na⁺, Cl⁻ spectators, no hydrolysis. |
| Strong acid + weak base (HCl + NH₃) | Acidic — NH₄⁺ (conjugate acid of weak base) hydrolyzes to regenerate H₃O⁺. |
| Weak acid + strong base (HClO + NaOH) | Basic — ClO⁻ (conjugate base of weak acid) hydrolyzes to regenerate OH⁻. |
| Weak + weak (HClO + NH₃) | Depends on relative Ka/Kb strengths — no fixed direction. |
Must-Know Points
- Water is never included in Kw, Ka, or Kb expressions — as a pure liquid, its concentration doesn't meaningfully change.
- Ka < 1.0 = weak acid; Kb < 1.0 = weak base. Smaller K always means weaker/less dissociation.
- Worked example: 0.100 M acetic acid (Ka=1.8×10⁻⁵) → x²≈1.8×10⁻⁶ → [H₃O⁺]≈1.3×10⁻³ M → pH≈2.87.
Common MCAT Trap
- Any acid/base NOT on the short strong-list is weak by default — don't assume a common-sounding acid (e.g., HF, H₃PO₄) is strong just because it's familiar.
- Kw is fixed only at a GIVEN temperature — it changes if temperature changes (like any equilibrium constant), so pH+pOH=14 only holds at 298 K.
- In salt hydrolysis, correctly identify which ion is the conjugate acid/base of the WEAK component — the spectator ion (from the strong component) never hydrolyzes.
Quick Recall
What is [OH⁻] in a solution with pH 4?
Is a solution of NH₄Cl acidic, basic, or neutral?
If Ka(acid) is large, what does that say about Kb of its conjugate base?