DAT General Chemistry: Solubility Rules and Net Ionic Equations

DAT High-Yield Concept

Always-Soluble Ions: Group 1 alkali metal cations ( $\ce{Li^+}$ , $\ce{Na^+}$ , $\ce{K^+}$ , $\ce{Rb^+}$ , $\ce{Cs^+}$ ), ammonium ( $\ce{NH4^+}$ ), nitrate ( $\ce{NO3^-}$ ), acetate ( $\ce{C2H3O2^-}$ ), and perchlorate ( $\ce{ClO4^-}$ ) form soluble compounds in water without exception.
Strong Electrolytes Dissociate: Soluble ionic compounds split completely into aqueous ions in complete ionic equations. Insoluble solids, gases, liquids, and weak electrolytes remain intact in molecular form.
Spectator Ions: Ions present in identical chemical states on both reactant and product sides do not participate in the net reaction and are canceled out.
Net Ionic Equations: Show only species undergoing a chemical transformation, conserving both total mass and total electrical charge.

Solubility Rules

Solubility Definition: An ionic compound is classified as soluble if it dissolves significantly in water (typically

> \pu{0.1 M}

) to form solvated ions, and insoluble if it forms a precipitate solid (

\ce{(s)}

Predicting precipitation requires applying general aqueous solubility rules. Water-soluble salts break apart into hydrated cations and anions due to favorable ion-dipole interactions that overcome lattice energy.

Mnemonic — Soluble Ions ("NAG SAG"):

Nitrates ( $\ce{NO3^-}$ )
Acetates ( $\ce{C2H3O2^-}$ )
Group 1 alkali metals ( $\ce{Li^+}$ , $\ce{Na^+}$ , $\ce{K^+}$ , etc.)
Sulfates ( $\ce{SO4^{2-}}$ )*
Ammonium ( $\ce{NH4^+}$ )
Group 17 halides ( $\ce{Cl^-}$ , $\ce{Br^-}$ , $\ce{I^-}$ )**

Exceptions:
** Halide exceptions (PMS): $\ce{Pb^{2+}}$ , $\ce{Hg2^{2+}}$ , $\ce{Ag^+}$
* Sulfate exceptions (Castro Bear / PMS): $\ce{Ca^{2+}}$ , $\ce{Sr^{2+}}$ , $\ce{Ba^{2+}}$ , $\ce{Pb^{2+}}$ , $\ce{Ag^+}$

Ion Type	General Solubility	Key Exceptions (Insoluble/Soluble)
Group 1 cations, $\ce{NH4^+}$ , $\ce{NO3^-}$ , $\ce{C2H3O2^-}$ , $\ce{ClO4^-}$	Soluble	None (Always soluble on the DAT)
Halides ( $\ce{Cl^-}$ , $\ce{Br^-}$ , $\ce{I^-}$ )	Soluble	Insoluble with $\ce{Ag^+}$ , $\ce{Pb^{2+}}$ , $\ce{Hg2^{2+}}$
Sulfates ( $\ce{SO4^{2-}}$ )	Soluble	Insoluble with $\ce{Ca^{2+}}$ , $\ce{Sr^{2+}}$ , $\ce{Ba^{2+}}$ , $\ce{Pb^{2+}}$ , $\ce{Ag^+}$
Hydroxides ( $\ce{OH^-}$ ) & Sulfides ( $\ce{S^{2-}}$ )	Insoluble	Soluble with Group 1, $\ce{NH4^+}$ , $\ce{Ca^{2+}}$ , $\ce{Sr^{2+}}$ , $\ce{Ba^{2+}}$
Carbonates ( $\ce{CO3^{2-}}$ ) & Phosphates ( $\ce{PO4^{3-}}$ )	Insoluble	Soluble with Group 1 cations and $\ce{NH4^+}$

Worked Example: Predicting Salt Solubility

Problem: Classify $\ce{PbCl2}$ , $\ce{K2SO4}$ , and $\ce{BaSO4}$ as soluble or insoluble in aqueous solution.

Reasoning:

$\ce{PbCl2}$ : Halides are generally soluble, but $\ce{Pb^{2+}}$ is a listed PMS exception. Therefore, $\ce{PbCl2}$ is insoluble ( $\ce{(s)}$ ).
$\ce{K2SO4}$ : Contains potassium ( $\ce{K^+}$ ), a Group 1 alkali metal ion. Group 1 salts are always soluble ( $\ce{(aq)}$ ).
$\ce{BaSO4}$ : Sulfates are generally soluble, but $\ce{Ba^{2+}}$ is a listed exception. Therefore, $\ce{BaSO4}$ is insoluble ( $\ce{(s)}$ ).

Common Exam Trap: Assuming all halides or sulfates are soluble without checking for cation exceptions like $\ce{Pb^{2+}}$ or $\ce{Ba^{2+}}$ .

Complete Ionic Equations and Spectator Ions

A molecular equation writes all reactants and products as neutral compounds. When aqueous solutions of two soluble ionic compounds are mixed, they may undergo an ion-exchange process—a double-replacement reaction (covered in depth in its own lesson). To show the actual state of dissolved ions, we write a complete ionic equation.

Complete Ionic Equation: Expresses all strong soluble electrolytes as dissociated aqueous ions (

\ce{(aq)}

), while leaving insoluble precipitates (

\ce{(s)}

), pure liquids (

\ce{(l)}

), gases (

\ce{(g)}

), and weak electrolytes in intact molecular form.

Spectator ions are aqueous ions that appear identically on both sides of the reaction arrow. They do not undergo chemical change, oxidation state changes, or phase changes.

Worked Example: Writing Complete Ionic Equations

Starting Point: Aqueous lead(II) nitrate reacts with aqueous potassium iodide.

1. Write the balanced molecular equation:

\ce{Pb(NO3)2(aq) + 2KI(aq) -> PbI2(s) + 2KNO3(aq)}

2. Dissociate all soluble strong electrolytes into separate ions:

\ce{Pb^{2+}(aq) + 2NO3^-(aq) + 2K^+(aq) + 2I^-(aq) -> PbI2(s) + 2K^+(aq) + 2NO3^-(aq)}

3. Identify spectator ions: $\ce{K^+(aq)}$ and $\ce{NO3^-(aq)}$ remain aqueous ions on both reactant and product sides.

Common Exam Trap: Incorrectly splitting solid precipitates like $\ce{PbI2(s)}$ into aqueous ions in the complete ionic equation.

Net Ionic Equations

A net ionic equation is obtained by canceling all spectator ions from the complete ionic equation. It highlights the essential chemical transformation.

Canceling $\ce{2K^+(aq)}$ and $\ce{2NO3^-(aq)}$ from the complete ionic equation yields the net ionic equation:

\ce{Pb^{2+}(aq) + 2I^-(aq) -> PbI2(s)}

From molecular to net ionic: cross out whatever never changes — A molecular equation is written the way the bottles are labelled, but it hides what is in the beaker. Soluble ionic compounds are already separated into free ions in water, so the complete ionic equation writes them that way, while anything that is not dissolved keeps its formula: solids, pure liquids and gases are never split. Once both sides are written as ions, some species appear identically on the left and on the right. Those are spectator ions, and cancelling them leaves the net ionic equation, which is the reaction itself. Here silver ion and chloride ion combine into insoluble silver chloride, while sodium ion and nitrate ion float through the mixture unchanged.

Every valid net ionic equation must satisfy two conservation principles:

Conservation of Mass: Equal numbers of each atom type on both sides (1 Pb atom, 2 I atoms).
Conservation of Charge: Equal total charge on both sides. Reactants: $+2 + 2(-1) = 0$ ; Products: $0$ (neutral solid).

Critical Thinking Questions

Mixing aqueous solutions of $\ce{Na2CO3}$ and $\ce{CaCl2}$ forms a white precipitate. What are the spectator ions, and what is the net ionic equation?
Why is the net ionic equation for the neutralization of any strong acid with any strong base always $\ce{H^+(aq) + OH^-(aq) -> H2O(l)}$ ?
If aqueous solutions of $\ce{KNO3}$ and $\ce{NaCl}$ are mixed, no precipitate, gas, or weak electrolyte forms. What is the net ionic equation for this mixture?

References

Sources returned by Google Search grounding during lesson verification.

Solubility Rules - Chemistry LibreTexts — chem.libretexts.org
16.18: Net Ionic Equations - Chemistry LibreTexts — chem.libretexts.org

Solubility Rules and Net Ionic Equations

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