A combination of two dissolved ionic compounds will produce a precipitate when the ions swap partners and form a new compound that is insoluble in water. The solid that crashes out of solution is the precipitate. To predict whether this happens, you need to know which ion pairs dissolve and which don’t, and that comes down to a set of solubility rules you can memorize or reference.
How Precipitates Form
When you mix two solutions of ionic compounds, the positive ions (cations) and negative ions (anions) are all floating freely in water. They can recombine in new ways. If one of those new pairings creates a compound that doesn’t dissolve, it clumps together as a solid and drops out of the liquid. This is a precipitation reaction, and the formation of that insoluble solid is what drives the whole reaction forward.
For example, if you mix a solution of potassium iodide with a solution of lead nitrate, the lead ions pair up with the iodide ions to form lead iodide, a bright yellow solid that falls to the bottom of the container. The potassium and nitrate ions stay dissolved and don’t participate in the reaction at all. Those bystanders are called spectator ions.
The Solubility Rules
These rules tell you which ionic compounds dissolve in water (soluble) and which don’t (insoluble). When a rule has exceptions, the exceptions take priority. Here’s the practical breakdown:
- Always soluble: Compounds containing sodium, potassium, lithium, rubidium, cesium, or ammonium ions dissolve in water with almost no exceptions. Nitrate salts are also generally soluble.
- Chlorides, bromides, and iodides: Soluble, except when paired with silver, lead, or mercury(I). So silver chloride (AgCl), lead bromide (PbBr₂), and mercury(I) chloride (Hg₂Cl₂) are all insoluble. Lead chloride is an interesting case because it dissolves in hot water but not at room temperature.
- Sulfates: Soluble, except when paired with barium, lead, calcium, or mercury(I). Barium sulfate (BaSO₄) and lead sulfate (PbSO₄) are classic insoluble sulfates. Silver sulfate is only slightly soluble.
- Hydroxides: Most are insoluble. The exceptions are hydroxides of sodium, potassium, and the other Group I metals, which dissolve freely. Calcium, strontium, and barium hydroxides are slightly soluble. Transition metal hydroxides like iron(III) hydroxide and aluminum hydroxide are insoluble.
- Sulfides: Most transition metal sulfides are highly insoluble, including cadmium sulfide, iron sulfide, zinc sulfide, and silver sulfide. Lead sulfide is also insoluble.
- Carbonates and phosphates: Generally insoluble, except when paired with Group I metals or ammonium.
Step-by-Step: Predicting the Precipitate
Here’s how to work through any combination and figure out whether a precipitate forms.
Step 1: Identify the ions. Write out the cation and anion from each compound. For example, if you’re mixing silver nitrate (AgNO₃) with sodium chloride (NaCl), your ions are Ag⁺, NO₃⁻, Na⁺, and Cl⁻.
Step 2: Swap partners. Pair each cation with the other compound’s anion. Ag⁺ pairs with Cl⁻ to make AgCl. Na⁺ pairs with NO₃⁻ to make NaNO₃.
Step 3: Check solubility. Look up each new compound. Sodium nitrate is soluble (sodium compounds always are). Silver chloride is insoluble (silver is one of the exceptions to the “chlorides are soluble” rule). A precipitate forms.
Step 4: Write the net ionic equation. Remove the spectator ions, the ones that appear dissolved on both sides. What’s left shows only the ions that actually react: Ag⁺(aq) + Cl⁻(aq) → AgCl(s). That’s the net ionic equation, and it captures the essence of what happened.
Common Combinations That Produce Precipitates
Some pairings come up repeatedly in chemistry courses and lab settings because they reliably produce visible solids:
- Silver nitrate + any chloride, bromide, or iodide: Produces insoluble silver halides. Silver chloride is a white solid. Silver iodide is pale yellow.
- Lead nitrate + potassium iodide: Produces lead iodide, a striking bright yellow precipitate once used as an artist’s pigment called iodine yellow.
- Barium chloride + sodium sulfate: Produces barium sulfate, a white insoluble powder so unreactive it’s used in medical imaging.
- Any soluble iron(III) salt + sodium hydroxide: Produces iron(III) hydroxide, a rust-brown solid.
- Copper sulfate + sodium sulfide: Produces copper sulfide, a black precipitate.
- Strontium nitrate + potassium sulfate: Produces strontium sulfate, a white solid that falls out of solution.
Notice the pattern. You’re always pairing one ion that’s normally well-behaved in solution with the specific partner that makes it insoluble.
Combinations That Won’t Produce a Precipitate
If both possible products are soluble, nothing precipitates. Mixing sodium chloride with potassium nitrate, for instance, gives you sodium nitrate and potassium chloride, both of which dissolve easily. The ions just float around together with no reaction. Anytime both products involve Group I metals, ammonium, or nitrate, you’re almost certainly not going to see a solid form.
Temperature Changes the Outcome
Most ionic compounds become more soluble as temperature increases, which means a precipitate that forms at room temperature might dissolve if you heat the solution. Lead chloride is the textbook example: it’s insoluble in cold water but dissolves readily in hot water. This property is actually used in labs to separate lead from silver and mercury(I), since heating the mixture dissolves the lead chloride while the other two chlorides stay solid.
A few compounds behave in the opposite direction. Calcium acetate, for example, is less soluble at higher temperatures. Heating a calcium acetate solution past about 80°C causes a precipitate to form, and cooling it dissolves the solid again.
Quantitative Prediction With Ksp
The solubility rules give you a yes-or-no answer, but chemistry also offers a more precise tool. Every sparingly soluble compound has a solubility product constant, abbreviated Ksp, which is a number representing how much of it can dissolve. When you mix two solutions, you can calculate a value called Q (the ion product) from the actual concentrations of the ions present. If Q is greater than Ksp, the solution has more dissolved ions than it can hold, and a precipitate forms. If Q is less than Ksp, everything stays dissolved. If Q equals Ksp exactly, the solution is at its limit, saturated, with no net change.
This matters when you’re working with dilute solutions. Two compounds might have an insoluble product according to the basic rules, but if the ion concentrations are extremely low, Q might never exceed Ksp and no visible precipitate appears.
Precipitation in Water Treatment
Precipitation reactions aren’t just a classroom exercise. Municipal water treatment plants use them daily to remove dissolved metals and soften hard water. Hard water contains dissolved calcium and magnesium compounds, and adding lime (calcium oxide) triggers the formation of calcium carbonate, an insoluble solid that settles out and pulls other impurities with it. Lime is cheap and widely available, though it generates large volumes of sludge, so it’s often combined with other chemicals like ferrous sulfate to improve efficiency.
For removing heavy metals from industrial wastewater, treatment facilities add hydroxide compounds that convert soluble metal ions into insoluble metal hydroxides or carbonates. Aluminum-based compounds like alum react with natural alkalinity in wastewater to form insoluble aluminum hydroxide, which acts as a coagulant to sweep contaminants out of solution. The same core principle applies: pair the dissolved ion with a partner that makes it insoluble, and it drops out of the water as a solid you can physically remove.