Ionic Equilibria

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Ionic equilibrium is the equilibrium between ions and unionized molecules in solution. It is significant because it helps us: 1. Determine equilibrium constants and concentrations of ions 2. Understan

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Strong Electrolytes: - Ionize completely or almost completely in solution - Examples: HCl, HNO₃, H₂SO₄, NaOH, KOH, NaCl - Single arrow (→) used in equations Weak Electrolytes: - Dissociate only parti

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Degree of dissociation (α) is the fraction of total moles of electrolyte that dissociates into ions at equilibrium. Formula: α = (Number of moles dissociated)/(Total number of moles) Percent dissoci

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Arrhenius Theory: Acid: Substance that produces H⁺ ions in aqueous solution Example: HCl(aq) → H⁺(aq) + Cl⁻(aq) Base: Substance that produces OH⁻ ions in aqueous solution Example: NaOH(aq) → Na⁺(aq)

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Brønsted-Lowry Theory: Acid: Proton (H⁺) donor Base: Proton (H⁺) acceptor Example: HCl + NH₃ ⇌ NH₄⁺ + Cl⁻ Acid₁ Base₂ Acid₂ Base₁ Conjugate pairs: - HCl/Cl⁻ (acid/conjugate base) - NH₄⁺

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Lewis Theory: Acid: Electron pair acceptor Base: Electron pair donor Amphoteric Nature of Water: Water can act as both acid and base: As acid: H₂O + NH₃ ⇌ OH⁻ + NH₄⁺ As base: H₂O + HCl ⇌ H₃O⁺ + Cl⁻

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For weak acid HA: HA(aq) ⇌ H⁺(aq) + A⁻(aq) Acid dissociation constant: Kₐ = [H⁺][A⁻]/[HA] Where: - [H⁺] = concentration of hydrogen ions - [A⁻] = concentration of conjugate base - [HA] = concentrat

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Ostwald's Dilution Law relates degree of dissociation to concentration. For weak acid HA with degree of dissociation α: Kₐ = α²c/(1-α) For very weak acids (α << 1): Kₐ ≈ α²c Therefore: α = √(Kₐ/c)