Most salts formed from alkali metals (lithium, sodium, potassium, rubidium, cesium) and those containing nitrate or ammonium dissolve readily in water. Beyond those near-universal categories, solubility depends on the specific combination of positive and negative ions in the salt. A handful of reliable rules, along with their exceptions, cover the vast majority of common salts you’ll encounter.
Salts That Are Almost Always Soluble
Some categories of salts dissolve in water with virtually no exceptions. These are the safest starting points:
- Alkali metal salts. Any salt containing sodium, potassium, lithium, rubidium, or cesium as the positive ion will dissolve. Sodium chloride (table salt), potassium nitrate, and sodium sulfate are all examples.
- Ammonium salts. Salts built around the ammonium ion dissolve reliably. Ammonium chloride, ammonium sulfate, and ammonium nitrate all go into solution easily.
- Nitrate salts. Every common nitrate is soluble. Silver nitrate, for instance, dissolves freely even though most other silver salts do not.
If a salt falls into any of these three groups, you can treat it as soluble without worrying about exceptions.
Salts That Are Usually Soluble, With Exceptions
Several other groups of salts dissolve in water most of the time, but specific combinations create insoluble compounds. These exceptions matter because they come up frequently in chemistry and in everyday situations like water treatment.
Chlorides, bromides, and iodides (collectively called halides) are generally soluble. The important exceptions are halide salts of silver, lead, and mercury(I). Silver chloride, lead bromide, and mercury(I) chloride do not dissolve. This is why adding table salt to a solution containing silver ions produces a white solid that drops out of solution.
Sulfate salts are soluble in most cases, but barium sulfate, lead sulfate, silver sulfate, and strontium sulfate are not. Calcium sulfate is a borderline case. It’s technically “sparingly soluble,” meaning only a small amount dissolves. You may recognize calcium sulfate in its hydrated form as gypsum, the mineral used in drywall. Its limited solubility is actually useful in industrial processes because it prevents calcium and sulfate from building up in recycled water.
Salts That Are Usually Insoluble
Some categories of salts resist dissolving. If you see one of these negative ions paired with anything other than an alkali metal or ammonium, assume the salt is insoluble unless you know otherwise.
Hydroxides are only slightly soluble as a general rule. The exceptions are hydroxides of alkali metals (sodium hydroxide and potassium hydroxide dissolve easily) and, to a lesser degree, calcium, strontium, and barium hydroxides, which are slightly soluble. Hydroxides of transition metals like iron, cobalt, and aluminum are firmly insoluble.
Carbonates are frequently insoluble. Calcium carbonate is the classic example. It’s the main component of limestone, marble, and the scale that builds up inside water heaters and pipes. Sodium carbonate and potassium carbonate dissolve fine (alkali metal rule), but calcium carbonate, strontium carbonate, barium carbonate, iron carbonate, and lead carbonate do not.
Phosphates follow the same pattern. Calcium phosphate, which makes up much of bone mineral, has an extremely low solubility. At body temperature and neutral pH, the concentrations of calcium and phosphate ions that can coexist in solution with solid calcium phosphate are vanishingly small.
Sulfides of transition metals are insoluble. Cadmium sulfide, iron sulfide, zinc sulfide, and silver sulfide all refuse to dissolve. Silver sulfide is the black tarnish that forms on silverware.
Fluorides are frequently insoluble as well, including barium fluoride, magnesium fluoride, and lead fluoride. This sets fluorides apart from the other halides, which are generally soluble.
Why Some Salts Dissolve and Others Don’t
When a salt dissolves, water molecules pull individual ions away from the crystal. Water is a polar molecule, meaning one end carries a slight positive charge and the other a slight negative charge. The negative end of a water molecule is attracted to positive ions, and the positive end is attracted to negative ions. These attractions, called ion-dipole forces, are what power dissolution.
Whether a salt actually dissolves comes down to a tug-of-war between two forces. The first is the energy holding the ions together in the solid crystal (lattice energy). The second is the energy released when water molecules surround and stabilize the freed ions (hydration energy). If the hydration energy is large enough to overcome the lattice energy, the salt dissolves. If it isn’t, the salt stays solid.
This balance explains some patterns that might otherwise seem random. Small ions with high charges hold onto each other tightly in a crystal, making the lattice hard to break apart. But small ions also interact very strongly with water, which helps pull them into solution. The outcome depends on the specific size match between the positive and negative ion. When both ions are similar in size, especially if one or both carry multiple charges, they tend to form insoluble salts. This is why barium sulfate is insoluble (barium and sulfate are well-matched in size) while magnesium sulfate (Epsom salt) dissolves easily: the much smaller magnesium ion interacts so strongly with water that it overcomes the lattice energy.
The same logic works in reverse for small anions like fluoride and hydroxide. Lithium fluoride is only slightly soluble because lithium and fluoride are both small ions that grip each other tightly. Cesium fluoride, where the positive ion is much larger, dissolves readily. And barium hydroxide is quite soluble, while magnesium hydroxide (milk of magnesia) barely dissolves at all.
How Temperature Changes Solubility
For most solid salts, raising the temperature increases how much will dissolve. The degree of that increase varies enormously. Sodium chloride is famously flat: it dissolves about 36 grams per 100 grams of water at 20°C, and heating the water barely changes that number. Potassium nitrate is the opposite extreme. At 30°C, about 48 grams dissolve per 100 grams of water. At 60°C, that jumps to roughly 107 grams. If you cool a hot, concentrated potassium nitrate solution, the excess salt recrystallizes out as the solubility drops.
Calcium sulfate is an unusual case. Its solubility actually decreases at higher temperatures, which is why hot water heaters and boilers are especially prone to calcium sulfate scale. At temperatures above about 42°C (108°F), the hydrated form of calcium sulfate (gypsum) begins converting to anhydrite, which is even less soluble.
Solubility in Everyday Life
These rules aren’t just academic. Hard water is hard precisely because calcium and magnesium dissolve into groundwater as their soluble salts, particularly calcium bicarbonate. The U.S. Geological Survey defines water hardness as the amount of dissolved calcium and magnesium. When you heat hard water, calcium carbonate (insoluble) precipitates out of solution, forming the white crusty scale inside kettles, water heaters, and pipes. That’s a direct demonstration of the carbonate insolubility rule.
Soap scum is another example. The calcium dissolved in hard water reacts with soap to form calcium stearate, an insoluble salt that leaves a film on sinks and bathtubs. Water softeners work by swapping calcium and magnesium ions for sodium ions, and sodium salts of soap are soluble.
Quick Reference for Solubility Rules
When you need to predict whether a salt dissolves, work through these rules in order. If two rules seem to conflict, the one listed first takes priority.
- Alkali metal and ammonium salts: soluble, almost no exceptions.
- Nitrates: soluble.
- Chlorides, bromides, iodides: soluble, except with silver, lead, or mercury(I).
- Sulfates: soluble, except with barium, lead, silver, strontium, and calcium.
- Hydroxides: insoluble, except alkali metals (soluble) and calcium, strontium, barium (slightly soluble).
- Carbonates, phosphates, sulfides, fluorides: insoluble, except with alkali metals and ammonium.
- Silver salts: insoluble, except silver nitrate and silver acetate.
These rules cover the vast majority of salts you’ll encounter in a chemistry course, a lab, or real-world water chemistry. The exceptions are worth memorizing on their own, since they tend to appear repeatedly in precipitation reactions and practical applications.