Electrolyte Solutions & Acid–Base / Solubility Essentials

Electrolytes & Body Relevance

• Body fluids contain multiple ions: Na^+, K^+, Ca^{2+}, Mg^{2+}, Cl^- , SO4^{2-}, HPO4^{2-}, HCO_3^-
• Maintain osmotic pressure, acid-base balance; blood pH ≈ 7.35\text{–}7.45

Strong, Weak & Non-Electrolytes

• Electrolyte = substance whose aqueous solution conducts electricity (forms free ions)
  • Strong electrolyte: ~100 % dissociation (e.g. HCl, NaOH, NaCl)
  • Weak electrolyte: partial dissociation (e.g. CH3COOH, NH3)
• Non-electrolyte: no ions in solution (e.g. glucose)

Brønsted–Lowry Acid–Base Theory

• Acid = proton donor; Base = proton acceptor
• Conjugate pairs differ by one H^+: HB \rightleftharpoons H^+ + B^-
• Amphoteric species act as acid or base (e.g. H2O, HCO3^-)

Acid/Base Strength & Direction of Reaction

• Reaction proceeds from stronger acid/base pair to weaker pair
• Stronger acid → weaker conjugate base & vice-versa

Classification of Salt Solutions

• Neutral ions: cations of strong bases & anions of strong acids (except HSO_4^-)
• Basic anions = conjugate bases of weak acids; acidic cations = conjugate acids of weak bases/metal ions
  • Neutral cation + basic anion → basic solution (e.g. NaCN)
  • Acidic cation + neutral anion → acidic solution (e.g. NH_4Cl)

Water Auto-Ionization & pH Scale

• 2H2O \rightleftharpoons H3O^+ + OH^-
• Ion-product: K_w=[H^+][OH^-]=1.0\times10^{-14} (25 °C)
• Neutral: [H^+]=[OH^-]=10^{-7}\,\text{M}
• pH=-\log[H^+] ; pOH=-\log[OH^-] ; pH+pOH=14
  • Acidic: pH

Ionization Constants for Weak Species

• Weak acid: HA+H2O \rightleftharpoons H3O^+ + A^- ; K_a
• Weak base: B^- + H2O \rightleftharpoons HB + OH^- ; Kb
• pKa=-\log Ka ; pKb=-\log Kb
  • Larger Ka / smaller pKa → stronger acid
  • KaKb=Kw ; pKa+pK_b=14 (25 °C)

Strong vs Weak (must know)

• Strong acids: HCl, HNO3, HClO4, H2SO4
• Strong bases: NaOH, KOH, Ca(OH)2, Ba(OH)2
• Everything else is weak in water

Solubility Product K_{sp}

• For AmBn(s) \rightleftharpoons mA^{n+}+nB^{m-} : K_{sp}=[A^{n+}]^m[B^{m-}]^n (eq. concentrations)
• Ion product Q{sp} predicts precipitation • Q{sp}
• Common-ion effect: added common ion ↓ solubility (Le Châtelier)
• Salt effect (inert ions) can slightly ↑ solubility via ionic strength

Fractional Precipitation & Dissolution

• Salts with smaller K_{sp} precipitate first when same ion involved
• To dissolve a ppt: remove one ion (complexation, protonation, etc.) so Q{sp}

Quick Reference Equations

• pH=-\log[H^+]
• [H^+]=10^{-pH}
• pH+pOH=14
• KaKb=K_w
• K_{sp}=[A^{n+}]^m[B^{m-}]^n
• Strong acid/base ≈ strong electrolyte (≈100 % ionization)