Precipitation Reactions

• Precipitation reactions occur when dissolved substances react to form solid products (precipitate).

• Often involve exchange of ions in aqueous solutions (double-displacement reactions).

• Common examples include natural formations (e.g., coral reefs, kidney stones).

• Used in industries and chemical analysis techniques.

Solubility

• Solubility: Maximum concentration that can be achieved for a substance under specific conditions.

• Soluble: Substances with high solubility.

• Insoluble: Substances that precipitate when concentration exceeds solubility.

Solubility Guidelines

Always Soluble:

• Nitrates (NO3−), chlorates (ClO3−), acetates (C2H3O2−), bicarbonate (HCO3−)

• Alkali metal compounds (Li+, Na+, K+, Rb+, Cs+)

• Ammonium compounds (NH4+)

Soluble with Exceptions:

• Chlorides, bromides, iodides (Cl−, Br−, I−) - Exceptions: Ag+, Hg22+, Pb2+

• Sulfates (SO42−) - Exceptions: Ag+, Hg22+, Pb2+, Ca2+, Sr2+, Ba2+

Insoluble Except with Always Soluble Cations:

• Carbonates (CO32−)

• Phosphates (PO43−)

• Chromates (CrO42−)

• Sulfides (S2−)

Insoluble Except with Always Soluble Cations and Other Exceptions:

• Hydroxides (OH−) - Exceptions: Ca2+, Sr2+, Ba2+

Predicting Precipitation Reactions

1. Identify ions in reactants.

2. Consider possible cation/anion combinations.

3. Check solubility of products formed.

4. If at least one product is insoluble, a precipitation reaction occurs.

Example Equations

• Mixing potassium iodide (KI) and lead(II) nitrate (Pb(NO3)2) forms lead(II) iodide (PbI2) precipitate:

  • $2KI(aq) + Pb(NO3)2(aq) \rightarrow PbI2(s) + 2 KNO3(aq)$

• Net ionic equation: $Pb^{2+}(aq) + 2I^{-}(aq) \rightarrow PbI2(s)$

Further Examples

• AgNO3 + NaF:

  • $AgNO3(aq) + NaF(aq) \rightarrow AgF(s) + NaNO3(aq)$

• BaSO4 formation:

  • $Ba^{2+}(aq) + SO4^{2-}(aq) \rightarrow BaSO4(s)$

• No reaction with: Na2CO3 and NH4NO3 yielding no precipitate.