Solubility Rules

Rucete ✏ Chemistry In a Nutshell

1. Generally Soluble Compounds

The following ions form soluble compounds, with a few exceptions:

• Alkali metal ions (Li⁺, Na⁺, K⁺, Rb⁺, Cs⁺) – no exceptions
• Ammonium ion (NH₄⁺) – no exceptions
• Nitrates (NO₃⁻) – no exceptions
• Acetates (C₂H₃O₂⁻) – except AgC₂H₃O₂ (slightly soluble)
• Chlorates (ClO₃⁻) – no exceptions
• Perchlorates (ClO₄⁻) – no exceptions
• Halides (Cl⁻, Br⁻, I⁻) – soluble except with Ag⁺, Hg₂²⁺, or Pb²⁺
• Sulfates (SO₄²⁻) – soluble except with Ag⁺, Ca²⁺, Sr²⁺, Ba²⁺, Hg₂²⁺, or Pb²⁺

2. Generally Insoluble Compounds

The following compounds are insoluble, unless paired with alkali metals or ammonium:

• Carbonates (CO₃²⁻) – insoluble except with alkali metals or NH₄⁺
• Phosphates (PO₄³⁻) – same exceptions
• Chromates (CrO₄²⁻) – same exceptions
• Sulfides (S²⁻) – soluble with alkali metals, alkaline earth metals, or NH₄⁺
• Hydroxides (OH⁻) – insoluble except with alkali metals, and slightly soluble with Ca²⁺, Sr²⁺, Ba²⁺
• Sulfites (SO₃²⁻) – same exceptions as carbonates and phosphates

3. Solubility Application in Net Ionic Equations

• Use the solubility rules to determine which compounds form precipitates (solids) in double-replacement reactions.
• For example, when mixing Pb(NO₃)₂ and HCl:
  • Nitrates and H⁺ are soluble, but PbCl₂ is insoluble → forms a precipitate.
  • Net Ionic Equation:
    $${\text{Pb}}^{2+}(aq)+2{\text{Cl}}^{-}(aq)\to {\text{PbCl}}_2(s)$$