Reactions in solution and net ionic equations

Precipitation reactions and the solubility rules, acid–base and acid–carbonate reactions, and writing molecular, full ionic and net ionic equations with spectator ions, plus an investigation to predict and test the products of reactions in solution.

Mix a clear solution of lead(II) nitrate with a clear solution of potassium iodide and a brilliant yellow solid appears instantly. Four kinds of ion went into the beaker, yet only two of them did anything. Writing an equation that shows just those two is one of the most useful habits in chemistry, because it tells you what actually happened.

What is really in an ionic solution

When a soluble ionic compound dissolves, its ions separate and spread through the water, each surrounded by water molecules. A bottle labelled "0.10 mol L⁻¹ NaCl" contains no NaCl units at all: it holds Na⁺(aq) and Cl⁻(aq) ions, each at 0.10 mol L⁻¹. Strong acids behave the same way: HCl(aq) is fully ionised to H⁺(aq) and Cl⁻(aq). (You'll refine H⁺ to H₃O⁺ in Year 12.)

So mixing two ionic solutions just mixes ions. A reaction happens only if some pair of ions can be removed from solution by forming one of:

an insoluble solid: a precipitate;
a molecular substance such as water;
a gas that escapes, such as carbon dioxide.

If none of these can form, nothing reacts and the ions simply stay mixed.

Precipitation and the solubility rules

You met precipitates in Year 10 (Science lesson 3E). To predict one at Stage 6, pair each cation with the anion from the other solution and check the solubility rules. You need to know these:

Ions	Solubility	Main exceptions
Na⁺, K⁺, NH₄⁺ (and all group 1)	all soluble	none at this level
NO₃⁻, CH₃COO⁻	all soluble	none at this level
Cl⁻, Br⁻, I⁻	mostly soluble	Ag⁺, Pb²⁺ (PbCl₂ is slightly soluble)
SO₄²⁻	mostly soluble	Ba²⁺, Pb²⁺, Sr²⁺; Ca²⁺ and Ag⁺ slightly soluble
CO₃²⁻, PO₄³⁻	mostly insoluble	group 1 and NH₄⁺ are soluble
OH⁻	mostly insoluble	group 1, NH₄⁺ and Ba²⁺ soluble; Ca²⁺ slightly soluble

Colours that turn up in exams: AgCl white, AgBr cream, AgI pale yellow, PbI₂ bright yellow, Cu(OH)₂ pale blue, Fe(OH)₂ green, Fe(OH)₃ red-brown, BaSO₄ white. In Year 12 (Chemistry lesson 5I) you'll put numbers on "slightly soluble" with the solubility product, Ksp.

Molecular, full ionic and net ionic equations

The same reaction can be written three ways. Each tells you something different.

From molecular equation to net ionic equation for lead(II) nitrate and potassium iodide. The struck-out ions are spectators.

Molecular equation: write every substance as a full formula. Useful for calculating masses, because it names the actual chemicals you measured out.
Full ionic equation: split every soluble ionic compound and every strong acid into its ions. Keep solids, liquids, gases and molecular substances (water, CO₂) as formulas.
Net ionic equation: cross out the spectator ions, the ones that appear unchanged on both sides, and write what is left.

Writing a net ionic equation for a precipitation

Solutions of iron(III) chloride and sodium hydroxide are mixed. Predict the precipitate and write the net ionic equation.

The possible new pairs are Fe³⁺ with OH⁻ and Na⁺ with Cl⁻. Sodium salts are soluble, but iron(III) hydroxide is not, so a red-brown precipitate of Fe(OH)₃ forms.

Molecular: FeCl₃(aq) + 3NaOH(aq) → Fe(OH)₃(s) + 3NaCl(aq)

Full ionic: Fe³⁺(aq) + 3Cl⁻(aq) + 3Na⁺(aq) + 3OH⁻(aq) → Fe(OH)₃(s) + 3Na⁺(aq) + 3Cl⁻(aq)

Cancel 3Na⁺ and 3Cl⁻: $$\text{Fe}^{3+}\text{(aq)} + 3\text{OH}^-\text{(aq)} \rightarrow \text{Fe(OH)}_3\text{(s)}$$ Check that charge balances: +3 − 3 = 0 on the left and 0 on the right.

A net ionic equation is more general than the molecular one. Fe³⁺(aq) + 3OH⁻(aq) → Fe(OH)₃(s) describes iron(III) nitrate with potassium hydroxide just as well as iron(III) chloride with sodium hydroxide.

Acid–base reactions

An acid reacts with a base to form a salt and water. The salt's cation comes from the base and its anion from the acid. For an acid with a metal hydroxide:

$$\text{H}_2\text{SO}_4\text{(aq)} + 2\text{KOH(aq)} \rightarrow \text{K}_2\text{SO}_4\text{(aq)} + 2\text{H}_2\text{O(l)}$$

Write out the ions and everything cancels except the formation of water. Every strong acid–strong base neutralisation has the same net ionic equation:

$$\text{H}^+\text{(aq)} + \text{OH}^-\text{(aq)} \rightarrow \text{H}_2\text{O(l)}$$

H⁺ and OH⁻ combine to form water molecules; Na⁺ and Cl⁻ are left in solution unchanged.

There is usually nothing to see, but the temperature rises because neutralisation is exothermic. Metal oxides are bases too: CuO(s) + 2HCl(aq) → CuCl₂(aq) + H₂O(l), and the black powder dissolves to give a blue-green solution. Because CuO is a solid, it stays as a formula in the ionic equation: CuO(s) + 2H⁺(aq) → Cu²⁺(aq) + H₂O(l).

Acid–carbonate reactions

Acids react with carbonates (and hydrogen carbonates) to form a salt, water and carbon dioxide. The fizzing is the giveaway.

Marble chips in dilute hydrochloric acid. The gas is identified as CO₂ with limewater.

Acid and a solid carbonate

Write the molecular and net ionic equations for marble chips (calcium carbonate) reacting with dilute nitric acid.

Molecular: CaCO₃(s) + 2HNO₃(aq) → Ca(NO₃)₂(aq) + H₂O(l) + CO₂(g)

CaCO₃ is a solid, so it is not split. Full ionic: CaCO₃(s) + 2H⁺(aq) + 2NO₃⁻(aq) → Ca²⁺(aq) + 2NO₃⁻(aq) + H₂O(l) + CO₂(g)

Only NO₃⁻ is a spectator: $$\text{CaCO}_3\text{(s)} + 2\text{H}^+\text{(aq)} \rightarrow \text{Ca}^{2+}\text{(aq)} + \text{H}_2\text{O(l)} + \text{CO}_2\text{(g)}$$

With a soluble carbonate the carbonate ion appears separately: 2H⁺(aq) + CO₃²⁻(aq) → H₂O(l) + CO₂(g). Hydrogen carbonates give H⁺(aq) + HCO₃⁻(aq) → H₂O(l) + CO₂(g), which is why baking soda neutralises spilt acid.

InteractiveSpectator ion sieve

Try this

Choose Pb²⁺ and I⁻. The net ionic equation is Pb²⁺(aq) + 2I⁻(aq) → PbI₂(s). Turn on Sieve out spectators: which two ions disappear?
Choose Fe³⁺ and OH⁻ and compare the full ionic equation with the worked example above. Why are there 3NO₃⁻ here instead of 3Cl⁻?
Choose Ba²⁺ and Cl⁻. The readout lists four spectator ions. What does that tell you?
Choose H⁺ with OH⁻, then H⁺ with CO₃²⁻. Which product removes the ions in each case?
Try Cu²⁺ with I⁻. Why doesn't the simple precipitation pattern work here?

Investigation: predicting and testing products in solution

Aim: to predict the products of six reactions in solution, then test the predictions.

Reactions: (1) AgNO₃ + NaCl; (2) CuSO₄ + NaOH; (3) BaCl₂ + Na₂SO₄; (4) HCl + NaOH with a few drops of universal indicator; (5) HCl + Na₂CO₃; (6) KNO₃ + NaCl. Use 0.1 mol L⁻¹ solutions, 2 mL of each in a test tube.

Method: write a predicted molecular and net ionic equation for each before you mix. Mix, observe against a dark background, and record colour, precipitate, gas and temperature change (feel the tube or use a thermometer). For reaction 5, pass the gas into limewater using a delivery tube.

Typical results: (1) white precipitate; (2) pale blue precipitate; (3) white precipitate; (4) indicator changes from red to green and the tube warms slightly; (5) fizzing, limewater turns milky; (6) no visible change, because both possible new compounds are soluble.

Validity: use solutions of the same concentration and volume, clean tubes rinsed with distilled water, and a control tube of each solution on its own so a "cloudy" result isn't just a contaminated tube. Repeating each test confirms your observations are reliable.

Safety: safety glasses. Silver nitrate stains skin; barium compounds are toxic; sodium hydroxide is corrosive. Put silver and barium waste in the residue container.

Test yourself

Which mixture of 0.1 mol L⁻¹ solutions produces a precipitate?
What is the net ionic equation for the reaction of sodium carbonate solution with dilute sulfuric acid?
Silver nitrate solution is added to calcium chloride solution. Which ions are spectators?
Which is the correct full ionic equation for magnesium hydroxide (a solid) reacting with hydrochloric acid?
In the net ionic equation 3Ca²⁺(aq) + 2PO₄³⁻(aq) → Ca₃(PO₄)₂(s), what is the total charge on the left-hand side?

Name each form of the equation.

Predict the result of mixing each pair of solutions.

Ions that are present but take no part in a reaction are called ions. A net ionic equation must balance in both atoms and . The net ionic equation for any strong acid with a strong base is H⁺(aq) + → H₂O(l). An acid reacting with a carbonate produces a salt, water and . All are soluble.

5 marks

A student mixes 20 mL of 0.10 mol L⁻¹ copper(II) sulfate with 20 mL of 0.10 mol L⁻¹ barium hydroxide. Predict all the products, describe what the student would see, and write the full ionic and net ionic equations.

Your answer