Solubility rules and precipitates

TL;DR:

• A precipitate is a solid that forms when ions in a mixed solution combine into an insoluble compound.

• Solubility rules tell you which combinations stay dissolved and which fall out as solids.

• Silver nitrate with sodium chloride gives solid AgCl. The net ionic equation is Ag⁺(aq) + Cl⁻(aq) → AgCl(s), with Na⁺ and NO₃⁻ as spectators.

Pour two clear solutions together and sometimes the mixture turns cloudy at once, with a white solid drifting down. Other times nothing happens at all. A short list of solubility rules lets you predict which it will be, and a net ionic equation shows what really changed.

What happens when salts dissolve

Soluble ionic compounds split into separate ions in water. Sodium chloride in water is really free Na⁺(aq) and Cl⁻(aq). The label (aq) means dissolved in water, (s) means solid.

When you mix two solutions, all the ions meet. If any pairing of a positive ion and a negative ion forms a compound that does not dissolve, those ions lock together into a solid and leave the solution. That solid is the precipitate.

Common solubility rules

These are general patterns, listed in the order that is easiest to apply. Your course may use a slightly different list, so treat this as the typical version.

• Compounds of Group 1 ions (Na⁺, K⁺, and so on) and ammonium (NH₄⁺) are soluble.

• Nitrates (NO₃⁻) and acetates are soluble.

• Chlorides, bromides and iodides are soluble, except those of Ag⁺, Pb²⁺ and Hg₂²⁺.

• Sulfates are soluble, except those of Ba²⁺, Sr²⁺ and Pb²⁺ (and Ca²⁺ is only slightly soluble).

• Carbonates, phosphates, sulfides and hydroxides are mostly insoluble, except with Group 1 ions and ammonium.

Worked example: silver nitrate and sodium chloride

1. Write the ions in each solution: AgNO₃ gives Ag⁺ and NO₃⁻. NaCl gives Na⁺ and Cl⁻.

2. Swap partners: Ag⁺ with Cl⁻ gives AgCl, and Na⁺ with NO₃⁻ gives NaNO₃.

3. Apply the rules: sodium and nitrate compounds are soluble. Chlorides are soluble except with Ag⁺, so AgCl is insoluble.

4. Molecular equation: AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq).

5. Full ionic equation: Ag⁺(aq) + NO₃⁻(aq) + Na⁺(aq) + Cl⁻(aq) → AgCl(s) + Na⁺(aq) + NO₃⁻(aq).

6. Cancel the ions that appear unchanged on both sides, Na⁺ and NO₃⁻. These are spectator ions.

7. Net ionic equation: Ag⁺(aq) + Cl⁻(aq) → AgCl(s).

The net ionic equation is the real chemistry: silver and chloride ions combine. The sodium and nitrate ions just watch. This is why silver nitrate would make a white precipitate with any soluble chloride, not just sodium chloride.

Check that the equation balances in atoms and charge. Left side charge: (+1) + (−1) = 0. Right side: AgCl is neutral, so 0. Atoms: 1 Ag and 1 Cl on each side.

How much solid forms?

Say 25.0 mL of 0.100 M silver nitrate is mixed with plenty of sodium chloride. Moles of silver ions: 0.0250 L × 0.100 mol/L = 0.00250 mol. The mole ratio of Ag⁺ to AgCl is 1 to 1, so 0.00250 mol of AgCl forms. The molar mass of AgCl is 107.87 + 35.45 = 143.32 g/mol, so the mass is 0.00250 × 143.32 = 0.358 g.

Silver nitrate stains and harms: Silver nitrate stains skin and clothes dark brown or black and can burn the skin and eyes. Use it only in a supervised lab with goggles and gloves, and put waste containing silver in the collection container, not down the sink.

When nothing happens

Try sodium nitrate with potassium chloride. The possible swaps give sodium chloride and potassium nitrate, and every one of those is soluble. All four ions stay in solution, so every ion is a spectator and there is no net ionic equation. The mixture just contains four kinds of ions.

Try a second example for practice: lead(II) nitrate with potassium iodide. The swaps give PbI₂ and KNO₃. Potassium and nitrate compounds are soluble, but iodides of lead are an exception, so PbI₂ is a solid. Net ionic equation: Pb²⁺(aq) + 2 I⁻(aq) → PbI₂(s). Check the charge: left (+2) + 2(−1) = 0, and right 0. The 2 in front of I⁻ is needed to balance the atoms. Lead compounds are toxic, so this is a paper exercise unless your teacher runs it in the lab.

How to check your answer

• Charge balance. The net ionic equation must have equal total charge on both sides.

• One solid at least. If nothing in the products is a solid, a gas or water, there is usually no reaction.

• Spectators. Only ions that appear in identical form on both sides cancel.

• Formulas. Write each compound with its ions' charges balanced before applying the rules.

Mistakes that cost marks

• Writing a precipitate as separate ions.

• Splitting an insoluble solid into ions in the full ionic equation.

• Cancelling an ion that is part of the real reaction.

• Forgetting subscripts or charges when swapping partners, such as Ca²⁺ with Cl⁻ needing CaCl₂.

• Using the rules as laws. Soluble and insoluble are matters of degree.

Where you will use it next

Precipitation is used to remove ions from water and to detect them in analysis. Why some compounds dissolve and others do not links to intermolecular forces and boiling points, and the amounts follow from stoichiometry.

Frequently asked questions

Is 'insoluble' the same as 'does not dissolve at all'?

Not quite. Insoluble compounds dissolve to a very small extent. In most courses the amount is small enough to ignore.

Why do we leave spectator ions out?

They do not take part in the change. Leaving them out shows the real reaction, which is the same whichever soluble salts supply the ions.