Solubility Rules and Precipitation Reactions

• Group I Elements and Ammonium: All salts containing $\text{Li}^{+}$, $\text{Na}^{+}$, $\text{K}^{+}$, $\text{Cs}^{+}$, $\text{Rb}^{+}$, or $\text{NH}_{4}^{+}$ are soluble.
• Nitrates: Salts containing the nitrate ion ($\text{NO}_{3}^{-}$) are generally soluble.
• Halides: Salts containing $\text{Cl}^{-}$, $\text{Br}^{-}$, or $\text{I}^{-}$ are generally soluble, except when paired with $\text{Ag}^{+}$, $\text{Pb}^{2+}$, or $\text{Hg}_{2}^{2+}$.
• Sulfates: Most sulfate salts ($\text{SO}_{4}^{2-}$) are soluble, with key exceptions: $\text{CaSO}_{4}$, $\text{BaSO}_{4}$, $\text{PbSO}_{4}$, $\text{Ag}_{2}\text{SO}_{4}$, and $\text{SrSO}_{4}$.
• Hydroxides: Group I hydroxides are soluble. Group II hydroxides (such as $\text{Ca}(\text{OH})_{2}$, $\text{Sr}(\text{OH})_{2}$, and $\text{Ba}(\text{OH})_{2}$) are only slightly soluble, while hydroxides of transition metals and $\text{Al}^{3+}$ are insoluble.

Identifying Insoluble Compounds

An insoluble compound is a substance that does not dissolve significantly in water. When two aqueous solutions react and produce an insoluble product, that product is called a precipitate.

If your predicted product falls into one of these categories, it will likely form a precipitate:

• Silver Salts: Most silver salts are insoluble, though $\text{AgNO}_{3}$ and $\text{Ag}(\text{C}_{2}\text{H}_{3}\text{O}_{2})$ are notable soluble exceptions.
• Sulfides: Transition metal sulfides (e.g., $\text{CdS}$, $\text{FeS}$, $\text{ZnS}$) and sulfides of arsenic, antimony, bismuth, and lead are highly insoluble.
• Carbonates, Chromates, and Phosphates: These are frequently insoluble. Examples include $\text{CaCO}_{3}$, $\text{PbCrO}_{4}$, and $\text{Ag}_{3}\text{PO}_{4}$.
• Fluorides: Many fluorides, such as $\text{BaF}_{2}$ and $\text{MgF}_{2}$, are frequently insoluble.

Exam Tip: Conflict Resolution Sometimes, two rules will seem to contradict each other. For example, "Halides are soluble" and "Silver salts are insoluble." If you are asked about $\text{AgBr}$, the rule regarding silver takes precedence. Therefore, $\text{AgBr}$ is insoluble and will form a precipitate.

Precipitation Reactions and Net Ionic Equations

A precipitation reaction occurs when a soluble cation combines with a soluble anion to form an insoluble solid. To accurately describe what is happening at the molecular level, chemists use net ionic equations.

Step-by-Step Example

Consider the reaction between potassium chloride ($\text{KCl}$) and silver nitrate ($\text{AgNO}_{3}$):

1. Molecular Equation: $\text{KCl}(aq) + \text{AgNO}_{3}(aq) \rightarrow \text{KNO}_{3}(aq) + \text{AgCl}(s)$ By applying solubility rules, we see that $\text{KNO}_{3}$ is soluble (contains Group I and Nitrate), but $\text{AgCl}$ is insoluble (Silver halide).

2. Total Ionic Equation: We break all soluble strong electrolytes into their constituent ions: $\text{K}^{+}(aq) + \text{Cl}^{-}(aq) + \text{Ag}^{+}(aq) + \text{NO}_{3}^{-}(aq) \rightarrow \text{K}^{+}(aq) + \text{NO}_{3}^{-}(aq) + \text{AgCl}(s)$

3. Net Ionic Equation: We remove the spectator ions—ions that appear unchanged on both sides of the equation ($\text{K}^{+}$ and $\text{NO}_{3}^{-}$ in this case). The remaining species show the actual chemical change: $\text{Ag}^{+}(aq) + \text{Cl}^{-}(aq) \rightarrow \text{AgCl}(s)$

Summary Checklist for Predicting Precipitates

When presented with a reaction, follow this workflow:

1. Identify all possible products by swapping the cations and anions of the reactants.
2. Apply the solubility rules to each potential product.
3. If any product is identified as insoluble, a precipitate will form.
4. Write the net ionic equation by focusing only on the ions that form the precipitate.