Solubility rules and net ionic equations: decide what stays together
Apply solubility exceptions before cancelling ions. Follow precipitation, weak-acid, gas-forming, and no-reaction examples.
Decide which substances remain intact before you cancel any ions. Most incorrect net ionic equations start with the wrong physical state. Perfect arithmetic cannot repair an equation that splits a precipitate into dissolved ions.
Use solubility rules to propose a classroom reaction. Then check whether the question supplies concentrations, temperature, or unusual solution conditions. Those details can require an equilibrium calculation instead of a simple soluble-or-insoluble label.
Read solubility rules as a decision table
The table below summarizes common aqueous rules from OpenStax Chemistry 2e, Section 4.2. “Insoluble” means low solubility under the stated conditions, not absolutely no dissolved material.
| Ion family | Usual classification | Important exceptions |
|---|---|---|
| Group 1 and ammonium salts | Soluble | Use these cation rules before the insoluble families below. |
| Nitrates, acetates, chlorates, bicarbonates | Soluble | No exceptions in this introductory table. |
| Chlorides, bromides, iodides | Soluble | Silver, lead(II), and mercury(I) salts. |
| Fluorides | Soluble | Group 2, lead(II), and iron(III) salts. |
| Sulfates | Soluble | Silver, barium, calcium, mercury(I), lead(II), and strontium salts. |
| Carbonates, chromates, phosphates | Insoluble | Group 1 and ammonium salts. |
| Hydroxides | Insoluble | Group 1 and barium hydroxides; review slight solubility separately. |
Fluoride does not share the chloride rule. Mercury(I) means the paired ion Hg₂²⁺, not mercury(II). Preserve the oxidation-state label when choosing an exception.
Textbook tables simplify borderline cases differently. Purdue's solubility reference marks calcium hydroxide as slightly soluble and strontium hydroxide as soluble. Its sulfide exceptions include sodium, potassium, ammonium, magnesium, calcium, strontium, and barium. Do not extend a short table to every metal sulfide without checking the particular compound.
The calculator applies its own stated classroom rules. It treats several slightly soluble salts as precipitates and reports that choice. Its fluoride list does not cover every exception in the textbook table. Read the displayed rule when the assignment involves a less familiar salt.
Build the ionic equation in a fixed order
- Identify the ions in each reactant. Build electrically neutral product formulas before balancing.
- Assign physical states using the problem's conditions and the relevant solubility rules.
- Balance the molecular equation with coefficients. Keep each compound's formula unchanged.
- Split dissolved strong electrolytes into ions. Keep solids, liquids, gases, and weak electrolytes intact.
- Cancel equal quantities of identical ions on opposite sides. Keep the smallest whole-number coefficients in the remaining equation.
- Count atoms and total charge on both sides. Both checks must pass.
OpenStax Section 4.1 explains molecular, complete ionic, and net ionic representations. The molecular equation preserves the identities of the starting compounds. The net ionic equation isolates the chemical change after spectator cancellation.
The symbol (aq) means dissolved in water. It does not mean that every substance forms separate ions. OpenStax Section 11.2 distinguishes strong electrolytes from weak electrolytes and nonelectrolytes.
Silver chloride: inspect the cancellation
Enter AgNO3 + NaCl in the net ionic equation calculator. The calculator predicts silver chloride as the solid product. Nitrate and sodium remain dissolved under these classroom rules.
Balanced molecular equation
AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq)Complete ionic equation
Ag⁺(aq) + NO₃⁻(aq) + Na⁺(aq) + Cl⁻(aq) → AgCl(s) + Na⁺(aq) + NO₃⁻(aq)Net ionic equation
Ag⁺(aq) + Cl⁻(aq) → AgCl(s)The calculator identifies these spectators: NO₃⁻, Na⁺. Remove them because their forms and quantities match across the complete ionic equation. Keep AgCl(s) intact; it represents the precipitate.
Notice what the shorter equation loses. It no longer names the original nitrate and sodium salts. Keep the molecular equation when you need to explain which solutions supplied the reacting ions.
Barium sulfate: let the exception control the product
Enter BaCl2 + Na2SO4. Barium sulfate belongs to the sulfate exceptions. The calculator balances the sodium chloride coefficient before it cancels the spectators.
BaCl₂(aq) + Na₂SO₄(aq) → BaSO₄(s) + 2NaCl(aq)Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s)The spectator list is Cl⁻, Na⁺. Do not split BaSO₄(s) simply because its formula contains recognizable ions. The state symbol controls that decision.
Acetic acid: dissolved does not mean fully dissociated
Enter CH3COOH + NaOH. The calculator keeps acetic acid molecular because it is a weak acid. It splits the aqueous sodium hydroxide and sodium acetate.
CH₃COOH(aq) + NaOH(aq) → H₂O(l) + NaCH₃COO(aq)CH₃COOH(aq) + OH⁻(aq) → CH₃COO⁻(aq) + H₂O(l)The spectator list contains Na⁺. Acetate remains in the net equation because the reaction forms it from acetic acid. Cancelling acetate against the acid would cancel different chemical species.
Use this example to check your splitting step. If your equation starts with free hydrogen ions alone, you discarded the weak acid's identity too early.
Carbonate and acid: account for the gas
Enter Na2CO3 + HCl. The calculator predicts water and carbon dioxide from the carbonic-acid intermediate. It then balances the acid requirement.
Na₂CO₃(aq) + 2HCl(aq) → 2NaCl(aq) + H₂O(l) + CO₂(g)CO₃²⁻(aq) + 2H⁺(aq) → H₂O(l) + CO₂(g)The spectator list is Na⁺, Cl⁻. Keep carbon dioxide as a gas and water as a liquid. This result describes the calculator's gas-forming classroom model, not a calculation of gas volume or final solution pH.
No net reaction is a complete answer
Enter NaCl + KNO3. The calculator returns No net reaction. Its spectator list includes every dissolved ion: Na⁺, Cl⁻, K⁺, NO₃⁻.
Exchanging salt names on paper does not create a chemical change. Under this model, the same ions remain in solution. Do not invent a precipitate merely because the exercise resembles a double-replacement reaction.
Common mistakes and the limit of these rules
- Changing subscripts to balance atoms changes the compound. Change coefficients instead.
- Cancelling ions before balancing can remove unequal amounts. Compare coefficients as well as formulas and charges.
- Splitting every aqueous formula destroys weak-acid equations. Check electrolyte strength separately from solubility.
- Using a solubility table to predict precipitate quantity ignores the amounts mixed. The calculator does not accept those quantities.
For concentration-dependent precipitation, compare the reaction quotient Q with the solubility product Ksp. Use concentrations after mixing, including dilution. OpenStax Section 15.1 explains that precipitation occurs when Q exceeds Ksp.
The calculator does not solve that equilibrium problem. It also does not model redox reactions or arbitrary complex-ion chemistry. Use its equations to inspect a supported aqueous reaction, then check the assumptions that the problem supplies.