Net Ionic Equations in Aqueous Solution
When aqueous solutions are mixed, ions can recombine to form an insoluble ionic solid
(a precipitate) or, for strong acid–strong base mixtures, liquid water. To best represent what actually
changes, we write a net ionic equation: an equation that includes only the species that change during the
reaction; spectator ions are omitted.
What is a Net Ionic Equation?
Start from the molecular equation, dissociate all strong electrolytes (soluble ionic compounds,
strong acids, strong bases) into ions to get the complete ionic equation, then cancel spectators on
both sides to obtain the net ionic equation. For insoluble products we write the solid as a neutral formula
(e.g., \(\ce{AgCl(s)}\)), not separated ions.
Workflow (always in this order)
- Write a plausible double replacement (ion exchange) or recognize acid–base neutralization.
- Use solubility rules (below) to assign states: (aq), (s), (l), (g).
- Write the complete ionic equation by dissociating all aqueous strong electrolytes into ions.
- Cancel spectator ions (identical species on both sides).
- Check atom count and charge balance; report the net ionic equation.
Solubility Guidelines (first matching rule wins)
- Group 1 metal cations (\(\ce{Li+}\), \(\ce{Na+}\), \(\ce{K+}\), \(\ce{Rb+}\), \(\ce{Cs+}\)) and \(\ce{NH4+}\) salts are soluble.
- \(\ce{NO3^-}\) (nitrate), \(\ce{CH3COO^-}\) / \(\ce{C2H3O2^-}\) (acetate), and \(\ce{ClO4^-}\) (perchlorate) are soluble.
- Salts of \(\ce{Ag+}\), \(\ce{Pb^{2+}}\), \(\ce{Hg2^{2+}}\) are insoluble (unless rules 1–2 apply).
- \(\ce{Cl^-}\), \(\ce{Br^-}\), \(\ce{I^-}\) salts are soluble (except with the cations in rule 3).
- \(\ce{CO3^{2-}}\), \(\ce{PO4^{3-}}\), \(\ce{S^{2-}}\), \(\ce{O^{2-}}\), \(\ce{OH^-}\) are insoluble.
Group 2 sulfides are slightly soluble; hydroxides of \(\ce{Ca^{2+}}\), \(\ce{Sr^{2+}}\), \(\ce{Ba^{2+}}\) are slightly soluble.
- \(\ce{SO4^{2-}}\) (sulfates) are soluble except those of \(\ce{Ca^{2+}}\), \(\ce{Sr^{2+}}\), \(\ce{Ba^{2+}}\).
Model molecular patterns
\(\ce{MA(aq) + NB(aq) -> MB(s/l/g) + NA(aq)}\)
If neither product is a solid (or water or gas), no reaction is written at the net-ionic level.
Worked Example 1 — Precipitation
Predict the net ionic equation when \(\ce{AgNO3(aq)}\) is mixed with \(\ce{KBr(aq)}\).
\(\ce{AgNO3(aq) + KBr(aq) -> AgBr(s) + KNO3(aq)}\)
\[
\text{Complete ionic:}\quad
\ce{Ag^+(aq) + NO3^-(aq) + K^+(aq) + Br^-(aq) -> AgBr(s) + K^+(aq) + NO3^-(aq)}
\]
\[
\text{Cancel spectators }(\ce{K^+},\ \ce{NO3^-})\ \Rightarrow\
\boxed{\ce{Ag^+(aq) + Br^-(aq) -> AgBr(s)}}
\]
Rule use: halide rule (4) with the silver exception (3) → \(\ce{AgBr}\) is insoluble.
Worked Example 2 — Strong Acid & Strong Base
Write the net ionic equation for \(\ce{HCl(aq)}\) and \(\ce{NaOH(aq)}\).
\(\ce{HCl(aq) + NaOH(aq) -> NaCl(aq) + H2O(l)}\)
\[
\text{Complete ionic:}\quad
\ce{H^+(aq) + Cl^-(aq) + Na^+(aq) + OH^-(aq) -> Na^+(aq) + Cl^-(aq) + H2O(l)}
\]
\[
\text{Net ionic:}\quad \boxed{\ce{H^+(aq) + OH^-(aq) -> H2O(l)}}
\]
Worked Example 3 — “No Reaction”
Do \(\ce{KNO3(aq)}\) and \(\ce{NaCl(aq)}\) react?
\(\ce{KNO3(aq) + NaCl(aq) -> KCl(aq) + NaNO3(aq)}\)
\[
\text{All four salts are soluble by rules 1–2. No precipitate, water, or gas forms.}
\]
\[
\text{Net ionic: } \boxed{\text{NR (no reaction)}}
\]
Checklist for Students
- Dissociate only strong electrolytes (aq). Keep solids, liquids (e.g., \(\ce{H2O}\)), and gases together.
- Apply the lower-numbered solubility rule first if rules conflict.
- Cancel identical species on both sides (including charge and phase).
- Always check charge balance in the final net ionic equation.
