Reaction Types, Solubility Guidelines, and Net Ionic Equations

Three driving forces behind aqueous reactions

When two solutions are mixed, a reaction occurs only if something removes ions from solution or moves electrons. Ask three questions. Does a cation from one solution and an anion from the other form a compound with low solubility? Does an acid transfer a proton to a base, forming water or another weak electrolyte? Do any atoms change oxidation number, so that electrons pass from one species to another?

A yes to the first question means a precipitation reaction, a yes to the second an acid–base reaction, and a yes to the third a redox reaction. Gas-forming reactions such as HCO₃⁻(aq) + H⁺(aq) → H₂O(l) + CO₂(g) are usually filed under acid–base because the first step is proton transfer. Combustion and the displacement of one metal by another are redox reactions.

The families are not exclusive. When aqueous barium hydroxide neutralizes sulfuric acid, Ba(OH)₂(aq) + H₂SO₄(aq) → BaSO₄(s) + 2H₂O(l), water forms and an insoluble sulfate precipitates in the same reaction.

Deciding which species to write as ions

A net ionic equation depends on knowing which substances exist as separated ions in water. Strong electrolytes dissociate essentially completely: soluble ionic compounds, strong acids, and strong bases. Write these as ions when they are labeled (aq). Everything else keeps its formula: solids, liquids, gases, weak acids, weak bases, and molecular solutes such as glucose.

The common strong acids are HCl, HBr, HI, HNO₃, HClO₄, and H₂SO₄ (for its first proton). Treat almost every other acid you meet in general chemistry, including HF, acetic acid (CH₃COOH), and carbonic acid, as weak, and keep it as a molecule. The strong bases are the group 1 hydroxides and the heavier group 2 hydroxides. Ammonia is a weak base and stays as NH₃(aq).

Polyatomic ions stay intact when a compound dissociates. Aqueous potassium chromate gives 2K⁺(aq) and CrO₄²⁻(aq), not separate potassium, chromium, and oxygen ions.

Applying the solubility guidelines

The solubility guidelines summarize observed patterns for ionic compounds in water near room temperature. Soluble means a useful amount dissolves; insoluble means the solubility is low enough that a solid forms when the ions meet at ordinary laboratory concentrations. No ionic compound has a solubility of exactly zero. The list below is based on Table 4.1 in OpenStax Chemistry 2e, section 4.2.

To predict a precipitate, list every ion present after mixing, then pair each cation with each anion it did not start with. If any new pairing falls in an insoluble category, that compound precipitates. If every possible pairing is soluble, as when potassium nitrate and sodium chloride solutions are mixed, the ions simply coexist and there is no net reaction.

• Compounds of group 1 cations (Li⁺, Na⁺, K⁺, Rb⁺, Cs⁺) and of NH₄⁺ are soluble.
• Nitrates, acetates, hydrogen carbonates, and chlorates are soluble.
• Chlorides, bromides, and iodides are soluble except those of Ag⁺, Hg₂²⁺, and Pb²⁺.
• Fluorides are soluble except those of Mg²⁺, Ca²⁺, Sr²⁺, Ba²⁺, Pb²⁺, and Fe³⁺.
• Sulfates are soluble except those of Ag⁺, Ba²⁺, Ca²⁺, Hg₂²⁺, Pb²⁺, and Sr²⁺.
• Carbonates, chromates, phosphates, and sulfides are insoluble except with group 1 cations and NH₄⁺.
• Hydroxides are insoluble except with group 1 cations and Ba²⁺; Ca(OH)₂ and Sr(OH)₂ are slightly soluble.

Building the net ionic equation step by step

Each step has a reason. Balancing the molecular equation first guarantees the atom counts. Splitting strong electrolytes shows what is really present in the beaker. Cancelling spectator ions removes species that appear unchanged on both sides and take no part in the change.

The charge check catches errors that an atom count misses. A net ionic equation can carry a nonzero total charge, but that charge must be the same on both sides, because charge is conserved just as atoms are.

• Write the balanced molecular equation with a state symbol on every formula.
• Rewrite each (aq) strong electrolyte as separate ions, multiplying by both the subscript and the coefficient.
• Cross out ions that appear unchanged on both sides. These are the spectator ions.
• Reduce the remaining coefficients to the smallest whole numbers.
• Check that each element and the total charge balance.

Worked example: precipitating silver chromate

You mix 18.0 mL of 0.0850 mol/L potassium chromate with excess aqueous silver nitrate. The new pairings are Ag⁺ with CrO₄²⁻ and K⁺ with NO₃⁻. Silver is not a group 1 cation, so red-brown silver chromate precipitates; potassium nitrate stays dissolved. Molecular equation: 2AgNO₃(aq) + K₂CrO₄(aq) → Ag₂CrO₄(s) + 2KNO₃(aq). Complete ionic equation: 2Ag⁺(aq) + 2NO₃⁻(aq) + 2K⁺(aq) + CrO₄²⁻(aq) → Ag₂CrO₄(s) + 2K⁺(aq) + 2NO₃⁻(aq). Cancelling K⁺ and NO₃⁻ gives the net ionic equation 2Ag⁺(aq) + CrO₄²⁻(aq) → Ag₂CrO₄(s). Charge check: 2(+1) + (−2) = 0 on the left, and 0 on the right.

To find the mass of precipitate, use IUPAC standard atomic weights rounded to two decimal places: Ag 107.87, Cr 52.00, O 16.00. The molar mass of Ag₂CrO₄ is 2(107.87) + 52.00 + 4(16.00) = 331.74 g/mol. Moles of chromate = 0.0180 L × 0.0850 mol/L = 1.53 × 10⁻³ mol, and the net ionic equation gives 1 mol Ag₂CrO₄ per 1 mol CrO₄²⁻. Mass = 1.53 × 10⁻³ mol × 331.74 g/mol = 0.508 g. Both measured quantities have three significant figures, so the answer does too.

Worked examples: a weak acid and a metal in solution

Aqueous acetic acid reacts with aqueous sodium hydroxide: CH₃COOH(aq) + NaOH(aq) → CH₃COONa(aq) + H₂O(l). Acetic acid is weak, so it stays molecular, while NaOH and sodium acetate are strong electrolytes. After Na⁺ cancels, the net ionic equation is CH₃COOH(aq) + OH⁻(aq) → CH₃COO⁻(aq) + H₂O(l). Compare nitric acid with potassium hydroxide, where both reactants are strong and the net ionic equation shrinks to H⁺(aq) + OH⁻(aq) → H₂O(l).

A strip of nickel placed in silver nitrate solution becomes coated with silver crystals. Nickel goes from 0 to +2 and silver from +1 to 0, so this is redox; nitrate is a spectator. The net ionic equation is Ni(s) + 2Ag⁺(aq) → Ni²⁺(aq) + 2Ag(s). The coefficient 2 is required by charge, not by atoms: Ni(s) + Ag⁺(aq) → Ni²⁺(aq) + Ag(s) balances every element but has a total charge of +1 on the left and +2 on the right.

Common mistakes in ionic equations and how to catch them

Most errors come from writing the wrong species, not from arithmetic. Before cancelling anything, ask whether each formula really exists as separate ions in water.

As a last check, read the net ionic equation as a physical event. Two silver ions and one chromate ion forming a solid is something you could watch; an equation that only swaps dissolved ions between partners describes nothing.

• Splitting a weak acid: HF(aq) stays molecular. Compare any acid against the six strong acids.
• Splitting a precipitate: a formula labeled (s) stays whole even though the solid is ionic.
• Breaking a polyatomic ion apart: keep NO₃⁻, SO₄²⁻, and CrO₄²⁻ intact.
• Misplacing subscripts: K₂CrO₄ gives 2K⁺; there is no K₂⁺ ion.
• Treating mercury(I) as Hg⁺: it exists as the diatomic Hg₂²⁺ cation.
• Leaving coefficients unreduced, such as 4Ag⁺ + 2CrO₄²⁻ → 2Ag₂CrO₄.
• Writing a net ionic equation when every pairing is soluble and no water, weak electrolyte, or gas forms, and nothing is oxidized or reduced. The correct answer is no reaction.