Net Ionic Equation Calculator

About the Net Ionic Equation Calculator

Type the reactants of a reaction in water — AgNO3 + NaCl, HCl + NaOH, Na2CO3 + HCl, or names such as silver nitrate + sodium chloride — and the calculator predicts the products, gives every substance its state, balances the molecular equation, writes the complete ionic equation with every dissolved strong electrolyte split into its ions, and cancels the spectator ions to leave the net ionic equation.

It handles double replacements that form a precipitate, neutralizations (strong and weak acids and bases, ammonia, insoluble hydroxides and metal oxides), reactions that make a weak acid, and the acids that break down into a gas as they form: H₂CO₃ → CO₂ + H₂O, H₂SO₃ → SO₂ + H₂O, H₂S, and NH₄OH → NH₃ + H₂O. Each state comes with the rule behind it. Type the products too, after an arrow, and it works out the ionic equations of any reaction in water — a single replacement such as Zn + HCl → ZnCl₂ + H₂ included.

Without a pass the free preview shows the net ionic equation in full, with the first rows of the states table and the first step of the working; the molecular and complete ionic equations are hidden; with a Pro pass, or after unlocking this result, everything shows and you can copy all three equations.

Common uses

• Chemistry homework: molecular, complete ionic and net ionic equations with the reason for every state.
• Predicting whether mixing two solutions in the lab will give a precipitate or a gas.
• Checking a net ionic equation you wrote yourself, including partly cancelled spectators.
• Revising solubility rules and the strong acids and bases.

How the equations are built

• Predict the products. In a double replacement the two compounds swap partners: AB + CD → AD + CB, each new compound balanced by the charges of its ions. An acid gives H⁺, so with a hydroxide it makes water; ammonia takes the acid’s H⁺ and becomes NH₄⁺.
• Give each substance its state by the rules below: a dissolved strong electrolyte is (aq) and splits into ions; a precipitate is (s); water is (l); a weak acid or base stays as molecules; gases are (g).
• Balance the molecular equation with whole numbers.
• Complete ionic equation: every (aq) strong electrolyte is written as its ions, with the coefficients multiplied out.
• Net ionic equation: the ions that appear unchanged on both sides — the spectators — cancel, and the coefficients are divided by any common factor. When every ion cancels, there is no net reaction.
• Something must drive it. A swap happens only when something is taken out of the dissolved ions: a precipitate, a gas, water or a weak acid or base. If all that is left would be an insoluble solid going into solution — BaSO₄ with hydrochloric acid, AgCl with nitric acid — nothing drives it and there is no reaction.

Solubility rules used

• Always soluble: compounds of the group 1 ions (Li⁺, Na⁺, K⁺ …) and of ammonium; nitrates, chlorates, perchlorates and permanganates; acetates too, except silver and mercury(I) acetate, which are only slightly soluble.
• Chlorides, bromides and iodides dissolve, except those of Ag⁺, Pb²⁺ (PbCl₂ is slightly soluble), Hg₂²⁺, Cu⁺ and Tl⁺ (and HgI₂).
• Fluorides dissolve, except those of Mg²⁺, Ca²⁺, Sr²⁺, Ba²⁺ and Pb²⁺.
• Sulfates dissolve, except those of Ba²⁺, Sr²⁺, Pb²⁺ and Hg₂²⁺; CaSO₄ and Ag₂SO₄ are slightly soluble.
• Hydroxides are insoluble, except the group 1 hydroxides and Ba(OH)₂; Ca(OH)₂ and Sr(OH)₂ are slightly soluble.
• Carbonates, phosphates, sulfites, oxalates, chromates and sulfides are insoluble, except with group 1 ions and ammonium (MgCrO₄ dissolves; the sulfides of Mg, Ca, Sr and Ba react with water).
• Metal oxides are written as solids: the oxide ion O²⁻ cannot exist in water (Na₂O and CaO react with it to form hydroxides), and with acids they react as bases.

These are the general rules for dilute solutions at room temperature; “slightly soluble” salts precipitate only from fairly concentrated solutions, and the calculator says so. For the measured solubility of a particular salt, see the IUPAC–NIST solubility database.

Strong and weak electrolytes

• Strong acids (fully ionized): HCl, HBr, HI, HNO₃, HClO₄, HClO₃ and H₂SO₄ (written as 2H⁺ + SO₄²⁻, although its second proton is only partly given up).
• Strong bases: the group 1 hydroxides and Ca(OH)₂, Sr(OH)₂, Ba(OH)₂ (what dissolves is fully ionized).
• Weak acids and bases stay mostly as molecules and are written whole: HF, HNO₂, HClO, HCN, CH₃COOH, HCOOH, H₃PO₄, H₂C₂O₄, NH₃.
• Gas-forming acids break down as they form: carbonic acid into CO₂ and water, sulfurous acid into SO₂ and water, thiosulfuric acid into sulfur, SO₂ and water; H₂S escapes as a gas. NH₄⁺ with OH⁻ gives ammonia and water — choose whether to write the ammonia as a gas or dissolved.

Sources

• IUPAC–NIST Solubility Data Series database: measured solubilities of individual salts
• IUPAC, Nomenclature of Inorganic Chemistry — IUPAC Recommendations: names and formulas of the ions

Limitations

• Products are predicted for double replacements and acid–base reactions only. Redox reactions such as single replacements, combustion and disproportionation are not predicted, and neither are the reactions of hydrides, nitrides and carbides with water: type their products to get the ionic equations.
• Organic compounds other than the acetate and formate ions, typed as such (CH3COOH, CH3COONa), do not form ions here: a molecular formula such as C6H12O6 is treated as a molecular compound.
• Solubility follows the general rules, which describe dilute solutions at room temperature. Borderline salts are flagged as slightly soluble; real solubility depends on concentration, temperature and complex formation (for example AgCl dissolving in ammonia).
• An acid with several hydrogens is assumed to react completely. With less base the salt keeps some hydrogen; type those products to get their equations.
• Sulfuric acid is written as fully ionized (2H⁺ + SO₄²⁻), as in most courses; its second ionization is in fact incomplete.

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