Precipitation Reactions and Net Ionic Equations

Predicting Precipitates: Solubility Guidelines

How do we know if mixing two solutions will create a solid? We use solubility guidelines to predict whether the resulting combinations of cations and anions will stay dissolved (aqueous, $aq$) or fall out of solution as a solid (precipitate, $s$).

The "Always Soluble" Rules

If a compound contains any of the following, it is generally soluble and will not form a precipitate:

1. Group I Alkali Metals: $\text{Li}^+$, $\text{Na}^+$, $\text{K}^+$, $\text{Rb}^+$, $\text{Cs}^+$.
2. Ammonium: $\text{NH}_4^+$.
3. Nitrates: $\text{NO}_3^-$.
4. Acetates: $\text{CH}_3\text{COO}^-$.

The "Usually Soluble" Rules (With Exceptions)

Some ions are generally soluble, but they have specific "enemies" that turn them into solids:

• Halides (Group 17: $\text{Cl}^-$, $\text{Br}^-$, $\text{I}^-$): Generally soluble, except when paired with $\text{Ag}^+$, $\text{Cu}^+$, $\text{Pb}^{2+}$, or $\text{Hg}_2^{2+}$.
• Sulfates ($\text{SO}_4^{2-}$): Generally soluble, except when paired with $\text{Pb}^{2+}$, $\text{Hg}_2^{2+}$, $\text{Ca}^{2+}$, $\text{Ba}^{2+}$, or $\text{Sr}^{2+}$.

Insoluble Compounds

If a combination of ions does not fall into the "always soluble" or "usually soluble" categories, it is likely to be an insoluble compound. Common insoluble groups include:

• Carbonates ($\text{CO}_3^{2-}$) and Phosphates ($\text{PO}_4^{3-}$).
• Sulfides ($\text{S}^{2-}$) of transition metals.
• Most Silver salts (except silver nitrate and silver acetate).

Exam Tip: Rule Precedence Sometimes rules seem to contradict each other. For example, halides are "generally soluble," but silver is "generally insoluble." If you are asked about $\text{AgBr}$, the "insoluble" rule for silver takes precedence. Therefore, $\text{AgBr}$ will form a precipitate.

Representing the Reaction

When a precipitation reaction occurs, chemists represent the process in three distinct stages: the molecular equation, the total ionic equation, and the net ionic equation.

1. The Molecular Equation

This shows the complete chemical formulas of all reactants and products as if they were intact molecules. Example: $\text{Pb}(\text{NO}_3)_2(aq) + 2\text{NaI}(aq) \rightarrow \text{PbI}_2(s) + 2\text{NaNO}_3(aq)$

2. The Total (Complete) Ionic Equation

Since the reactants are strong electrolytes, they exist in the water as dissociated ions. The total ionic equation breaks every soluble compound into its constituent ions, while keeping the solid precipitate intact. $\text{Pb}^{2+}(aq) + 2\text{NO}_3^-(aq) + 2\text{Na}^+(aq) + 2\text{I}^-(aq) \rightarrow \text{PbI}_2(s) + 2\text{Na}^+(aq) + 2\text{NO}_3^-(aq)$

3. The Net Ionic Equation

In the example above, notice that $\text{Na}^+$ and $\text{NO}_3^-$ appear exactly the same on both sides of the equation. These are called spectator ions. They are present in the solution, but they do not actually participate in the chemical change that forms the solid.

To find the net ionic equation, we "cancel out" the spectator ions. This leaves us with only the species directly involved in the formation of the precipitate: $\text{Pb}^{2+}(aq) + 2\text{I}^-(aq) \rightarrow \text{PbI}_2(s)$

Summary Checklist for Precipitation

To master this topic, follow these steps for any given reaction:

1. Identify the products: Determine which cation and anion pair up to form each new compound.
2. Apply solubility rules: Check if any of those products are insoluble based on the guidelines.
3. Write the molecular equation: Ensure it is balanced.
4. Write the total ionic equation: Dissociate all $(aq)$ strong electrolytes into ions.
5. Identify spectator ions: Find ions that are identical on both sides.
6. Write the net ionic equation: Remove the spectators to show the real chemistry.