When two clear solutions go cloudy

Mix silver nitrate with sodium chloride and you get an instant white cloud — silver chloride precipitating out the moment Ag⁺ meets Cl⁻. The chemistry is just the ion product crashing past the solubility ceiling: when [Ag⁺][Cl⁻] exceeds the Ksp of AgCl, the equilibrium can’t hold the ions in solution and they crystallize. Predicting whether a precipitate will form is bedrock for qualitative analysis, water treatment, and any synthesis where you separate products by selective precipitation.

You can predict most outcomes qualitatively from a short list of solubility rules. When concentrations are borderline or you need a quantitative answer, you switch to comparing the reaction quotient Q against Ksp.

The qualitative procedure

Step 1: Write out the ions actually in solution

Mix NaCl(aq) with AgNO₃(aq) and you have four ions floating around: Na⁺, Cl⁻, Ag⁺, NO₃⁻. The salts dissociate completely (they were soluble enough to be in solution to begin with).

Step 2: Pair the cations with the new anions

Skip the original pairings (you already know those are soluble). The new combinations are:

Na⁺ + NO₃⁻ → NaNO₃
Ag⁺ + Cl⁻ → AgCl

Step 3: Run each through the solubility rules

Generally soluble:

All nitrates (NO₃⁻)
All Group 1 (Li⁺, Na⁺, K⁺, etc.) and ammonium (NH₄⁺) salts
Most chlorides, bromides, iodides — except Ag⁺, Pb²⁺, Hg₂²⁺
Most sulfates — except BaSO₄, PbSO₄, CaSO₄, SrSO₄

Generally insoluble:

Most hydroxides — except Group 1, plus Ca(OH)₂, Sr(OH)₂, Ba(OH)₂
Most carbonates, phosphates, sulfides — except Group 1 and NH₄⁺

NaNO₃ is soluble (nitrate rule). AgCl is insoluble (silver chloride exception). Precipitate confirmed.

Step 4: Strip the spectators and write the net ionic equation

Spectator ions appear unchanged on both sides. Cancel them.

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

Full ionic: Na⁺ + Cl⁻ + Ag⁺ + NO₃⁻ → AgCl(s) + Na⁺ + NO₃⁻

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

Na⁺ and NO₃⁻ are along for the ride.

When you need numbers: Q vs. Ksp

The solubility rules are pass/fail. For borderline cases (slightly soluble salts, dilute solutions), compare the reaction quotient Q to Ksp.

Will a precipitate form when 50.0 mL of 0.0020 M Pb(NO₃)₂ mixes with 50.0 mL of 0.0040 M NaI?

Possible precipitate: PbI₂ (Ksp = 9.8 × 10⁻⁹).

After mixing, total volume is 100.0 mL — both concentrations halve:

[Pb²⁺] = 0.0010 M
[I⁻] = 0.0020 M

Q = [Pb²⁺][I⁻]² = (0.0010)(0.0020)² = 4.0 × 10⁻⁹

Q (4.0 × 10⁻⁹) < Ksp (9.8 × 10⁻⁹) → no precipitate forms.

If Q had exceeded Ksp, PbI₂ would crash out until Q dropped back down to Ksp.

Selective precipitation: a three-ion example

Add Na₂CO₃ to a solution containing both Ca²⁺ and Mg²⁺.

Possible precipitates:

CaCO₃, Ksp = 3.4 × 10⁻⁹
MgCO₃, Ksp = 6.8 × 10⁻⁶

CaCO₃ has the smaller Ksp, so as you slowly add CO₃²⁻, calcium precipitates first. Magnesium stays in solution until Q for MgCO₃ exceeds 6.8 × 10⁻⁶ — three orders of magnitude higher. This is the principle behind selective precipitation in qualitative analysis: tune the concentration of the common ion so one cation drops out cleanly while the other remains dissolved.

Where students get tripped up

Forgetting dilution. When you mix two solutions, the total volume is the sum, and every ion concentration drops accordingly. Use the post-mixing concentrations in your Q calculation, not the original ones.

Treating “insoluble” as absolutely insoluble. PbCl₂ and CaSO₄ get labeled insoluble in introductory courses, but they’re actually slightly soluble. For borderline situations, run the Ksp calculation rather than relying on the rule alone.

Forgetting stoichiometric exponents in Ksp. For PbI₂, Ksp = [Pb²⁺][I⁻]². The iodide concentration gets squared because two I⁻ ions appear in the formula. Every ion concentration is raised to its coefficient in the dissolution equation.

Practice

Try these and check your work with our Equilibrium Calculator:

Will a precipitate form when equal volumes of 0.10 M BaCl₂ and 0.10 M Na₂SO₄ are mixed? Ksp(BaSO₄) = 1.1 × 10⁻¹⁰.
Write the net ionic equation for Pb(NO₃)₂(aq) + KI(aq).
Does a precipitate form when 25.0 mL of 0.0050 M AgNO₃ mixes with 25.0 mL of 0.0030 M NaCl? Ksp(AgCl) = 1.8 × 10⁻¹⁰.
A solution holds 0.010 M Ca²⁺ and 0.010 M Ba²⁺. Adding Na₂SO₄ slowly — which sulfate precipitates first?
Predict the products and write the net ionic equation for FeCl₃(aq) + NaOH(aq).

How to Predict Precipitation Reactions