Solubility Rules: Which Compounds Dissolve in Water?

Understanding how substances behave in water is essential in chemistry. What determines this behavior is solubility, and chemists rely on solubility rules to quickly decide whether an ionic compound will dissolve in water or remain as a solid. These guidelines are handy for predicting the water solubility of ionic compounds without the need for an experiment.

Of course, you could memorize all the chemistry solubility rules by heart — but that’s rarely the most efficient approach. This article guides you through the logic behind the rules, offers a clear overview, and makes them easier to remember.

Scroll down to learn:

• What is the solubility definition?
• How chemistry solubility rules actually work?
• How to read the solubility rules chart?
• How to predict the solubility of salt in water?

You’ve probably encountered questions about solubility in everyday life. Sugar that dissolves instantly in coffee, or chalky limescale left behind after hard water evaporates, are simple examples of substances with different solubilities in water. These effects are especially noticeable when soluble salts or other ionic compounds are involved.

In chemistry, solubility describes how much of a substance (solute) can dissolve in a solvent. The solubility definition is the maximum possible concentration of a solute that can dissolve under specific conditions, such as room temperature and standard pressure. There are three possible outcomes:

• If the amount of solute is below the solubility limit, the solution is dilute.
• If the maximum amount is dissolved, the solution is saturated.
• Any extra solute remains undissolved and may separate as a solid (precipitate). To see how to write the full reaction, try our net ionic equation calculator.

Solubility is equally important in laboratory chemistry and medicine. Many reactions depend on whether a substance stays dissolved or forms a solid precipitate. In biological and medical contexts, some ions are safe or useful when bound within a compound but may become reactive or harmful once they dissolve and move freely in solution.

Solubility isn’t a fixed property and can vary with hydrogen ions and pH, as well as temperature and pressure. In this article, we focus on solubility in water at room temperature (about 25 °C or 77 °F) and standard pressure (1 atm, or 101.325 kilopascals), as these are the conditions you’re most likely to encounter in everyday situations and introductory chemistry problems.

In most cases, to determine whether a compound will dissolve in water, it’s enough to look at the chemical formula of the substance whose solubility you want to predict. Once you identify the ions that make up the ionic compound, you can apply a chemistry solubility rule to decide whether it will dissolve or form a solid:

1. Compounds containing alkali metal ions (Li⁺, Na⁺, K⁺, Cs⁺, Rb⁺) or the ammonium ion (NH₄⁺) are almost always soluble, with very few exceptions.
2. Salts containing nitrate ions (NO₃^-) are generally soluble in water.
3. Salts containing chloride (Cl^-), bromide (Br^-), or iodide (I^-) ions are generally soluble, except when paired with Ag⁺, Pb²⁺, or Hg₂²⁺.
4. Most sulfate salts are soluble, but important exceptions include CaSO₄, BaSO₄, SrSO₄, PbSO₄, and Ag₂SO₄.
5. Most hydroxide salts are insoluble or slightly soluble. Hydroxides of alkali metals are soluble, hydroxides of Ca²⁺, Sr²⁺, and Ba²⁺ are slightly soluble, while hydroxides of transition metals and Al³⁺ are insoluble.
6. Most silver salts are insoluble, with notable exceptions such as AgNO₃ and Ag(C₂H₃O₂).
7. Most carbonates, phosphates, and chromates are insoluble unless paired with alkali metal ions or ammonium. Examples include CaCO₃, PbCO₃, Ca₃(PO₄)₂, Ag₃PO₄, PbCrO₄, and BaCrO₄.
8. Most sulfides of transition metals (e.g., CdS, FeS, ZnS, Ag₂S) and many fluorides (such as BaF₂, MgF₂, and PbF₂) are insoluble in water.

Using solubility rules follows a simple process:

1. First, identify the cation and anion in the compound.
2. Then compare each ion against the solubility rules for ionic compounds, checking both the general rule and any listed exceptions.
3. If two rules seem to conflict, the more specific rule takes precedence.

In real solutions, solubility can also be affected by ionic strength of a solution.

A solubility rules chart condenses multiple chemistry solubility rules into a single, easy-to-scan reference. Instead of memorizing long lists, you can quickly predict the solubility of salt in water by identifying the anion in the compound, as most solubility patterns are grouped this way. Then check the corresponding row to see whether the compound is classified as soluble or insoluble, and finally look for any listed exceptions involving specific cations:

You can observe several common patterns that appear repeatedly in the table. Salts containing alkali metals (such as Li⁺, Na⁺, K⁺) or NH₄⁺ ions are typically soluble salts, while compounds such as carbonates (CO₃^2-), phosphates (PO₄^3-), and many hydroxides (OH^-) are usually insoluble unless paired with those same ions. Cl^-, Br^-, I^-, and SO₄^2- ions often dissolve as well, but they include well-known exceptions that are clearly marked in the solubility rules chart.

Once you’re familiar with the patterns and the solubility rules chart, predicting the water solubility of ionic compounds becomes straightforward. Let’s work through the examples below to gain some practice in applying chemistry solubility rules:

Match the following compounds to their likely solubility in water: KF, Cl₂, MgF₂

KF is soluble. Potassium is an alkali metal, so according to solubility rule #1, salts containing K⁺ are almost always soluble.

Chlorine gas (Cl₂) is a molecular element, not an ionic compound, so solubility rules don’t apply.

MgF₂ is insoluble (forms a precipitate). Many fluorides (including MgF₂) are poorly soluble according to solubility rule #8.

Determine whether BaSO₄ is soluble in water.

According to the solubility rule #4, most sulfates are soluble, but BaSO₄ is a listed exception. Therefore, BaSO₄ is not soluble and forms a solid precipitate.

What forms when AgNO₃ is mixed with NaCl?

The reaction produces AgCl and NaNO₃, as shown in the chemical equation.
According to the solubility rule #2, nitrates such as NaNO₃ are soluble and remain in solution.

According to solubility rule #3 and rule #6, chloride salts are usually soluble, but silver salts are insoluble, so AgCl forms a precipitate.