Basicity Calculators

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Basicity refers to the tendency of a substance to accept protons (Brønsted-Lowry definition) or donate electron pairs (Lewis definition). The strength of a base is quantified by its base dissociation constant Kb: the larger the Kb (smaller pKb), the stronger the base. For a weak base B: B + H₂O ⇌ BH⁺ + OH⁻; Kb = [BH⁺][OH⁻]/[B]. Strong bases (NaOH, KOH) dissociate completely. Weak bases (NH₃, amines, pyridine) partially ionize. The relationship between pKa and pKb of a conjugate acid-base pair: pKa + pKb = pKw = 14.

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Brønsted-Lowry Bases

A Brønsted-Lowry base is a proton acceptor. B + H₂O ⇌ BH⁺ + OH⁻. Kb = [BH⁺][OH⁻]/[B]. pKb = −log(Kb). Stronger base: larger Kb, smaller pKb.

Kb and pKb Values

  • NaOH: strong base; complete dissociation → Kb = ∞
  • Ammonia (NH₃): Kb = 1.8 × 10⁻⁵; pKb = 4.74
  • Methylamine (CH₃NH₂): Kb = 4.4 × 10⁻⁴; pKb = 3.36 (stronger than NH₃)
  • Aniline (C₆H₅NH₂): Kb = 4.0 × 10⁻¹⁰; pKb = 9.4 (very weak; lone pair on N delocalized into ring)
  • Pyridine: Kb = 1.8 × 10⁻⁹; pKb = 8.75

Conjugate Acid-Base Relationship

pKa(conjugate acid) + pKb(base) = 14 at 25°C. NH₄⁺ (pKa = 9.26) is the conjugate acid of NH₃ (pKb = 4.74): 9.26 + 4.74 = 14.00 ✓. Stronger base → weaker conjugate acid.

Factors Affecting Basicity

  • Electronegativity: O is less basic than N (O holds electrons more tightly)
  • Induction: electron-donating groups increase basicity; electron-withdrawing groups decrease it
  • Resonance: lone pair delocalized into adjacent π system → reduced basicity (aniline vs. cyclohexylamine)
  • Hybridization: sp³ N > sp² N > sp N in basicity

Glossary

Basicity
The tendency to accept protons (Brønsted-Lowry) or donate electron pairs (Lewis); measured by Kb and pKb; stronger base = larger Kb, smaller pKb; pKa(conjugate acid) + pKb = 14 at 25°C.
Kb (Base Dissociation Constant)
Kb = [BH⁺][OH⁻]/[B] for B + H₂O ⇌ BH⁺ + OH⁻; larger Kb = stronger base; pKb = −log(Kb); NH₃ Kb = 1.8×10⁻⁵; strong bases have effectively infinite Kb.
Lewis Base
An electron pair donor; broader than Brønsted-Lowry (all Brønsted bases are Lewis bases); includes ligands that coordinate to metal ions without accepting protons; examples: NH₃, H₂O, CO, F⁻.

Frequently Asked Questions

Basicity measures the tendency of a substance to accept protons (Brønsted-Lowry) or donate electron pairs (Lewis). It is quantified by the base dissociation constant Kb: for B + H₂O ⇌ BH⁺ + OH⁻, Kb = [BH⁺][OH⁻]/[B]. Larger Kb → stronger base (more OH⁻ produced). pKb = −log(Kb): smaller pKb → stronger base. Range: pKb < 0 = very strong; pKb 4–6 = moderate (amines); pKb 8–12 = weak; strong bases (NaOH) have essentially infinite Kb. The acid strength of the conjugate acid (BH⁺) is related: pKa(BH⁺) + pKb(B) = 14 at 25°C.

Amines are the most common organic bases. Basicity is affected by: (1) Inductive effects: alkyl groups are electron-donating → push electron density toward N → increase lone pair availability → increase basicity. Methylamine (Kb = 4.4×10⁻⁴) > ammonia (Kb = 1.8×10⁻⁵). But steric effects can reduce basicity for bulky secondary/tertiary amines. (2) Resonance: aniline (C₆H₅NH₂): lone pair on N is delocalized into the aromatic ring → reduced basicity (Kb = 4×10⁻¹⁰ vs. 1.8×10⁻⁵ for NH₃). (3) Hybridization: sp³ nitrogen (alkylamines) > sp² nitrogen (pyridine, aromatic) > sp nitrogen (nitriles); sp orbital holds electrons more tightly → less available for protonation.

For weak base B with concentration C and Kb: B + H₂O ⇌ BH⁺ + OH⁻. If x = [OH⁻] formed: Kb = x²/(C − x) ≈ x²/C (valid when x << C). x = [OH⁻] = √(Kb × C). pOH = −log[OH⁻]. pH = 14 − pOH. Example: 0.10 M NH₃ (Kb = 1.8×10⁻⁵): [OH⁻] = √(1.8×10⁻⁵ × 0.10) = √(1.8×10⁻⁶) = 1.34×10⁻³ M. pOH = −log(1.34×10⁻³) = 2.87. pH = 14 − 2.87 = 11.13. Check: x/C = 1.34×10⁻³/0.10 = 1.34% < 5% → approximation valid.

Brønsted-Lowry base: accepts a proton (H⁺); must have lone pair to bond with proton; examples: NH₃, OH⁻, F⁻, H₂O, CH₃NH₂. Lewis base: donates an electron pair; broader definition — includes all Brønsted-Lowry bases plus many others that don't accept protons. Examples: NH₃ (Lewis base to BF₃); H₂O (Lewis base to Cu²⁺); CO (Lewis base to Fe in hemoglobin); F⁻ (Lewis base to BF₃); alkenes (Lewis base to acid catalysts). Lewis acids are electron pair acceptors (BF₃, AlCl₃, metal ions). All Brønsted-Lowry bases are Lewis bases (lone pairs needed for both), but not vice versa — CO coordinates to metal ions as a Lewis base without acting as a Brønsted-Lowry base.