We said previously that soluble ionic compounds are strong electrolytes because they break apart to produce ions in a solution. But not all ionic compounds are soluble in water. We can generally categorize ionic compounds as soluble or insoluble in water. However, as we'll learn when we discuss equilibrium, even insoluble compounds will dissolve to a relatively small extent, so solubility is a continuum. We can follow these general rules to categorize ionic compounds as soluble or insoluble. In the top of this table, ionic compounds containing the listed ions are generally soluble with the exceptions listed when the ionic compound is insoluble. For example, salts with the chloride anion are typically soluble, but silver chloride is insoluble because silver is one of chloride's exceptions. In the bottom of the table, ionic compounds containing the listed ions are generally insoluble, with the exceptions listed when the ionic compound is soluble. For example, sulfide salts are generally insoluble, but calcium sulfide is soluble because calcium is one of sulfide's exceptions. Let's practice determining if ionic compounds are soluble or insoluble. For each of these examples, we'll first split the ionic compound into its component ions. Then we'll apply our solubility rules. In our first example, iron (III) chloride splits apart into iron (III) ions and chloride ions. Iron (III) ions are not in our table, but chloride is. Chloride salts are generally soluble and iron (III) ions are not exceptions, so iron (III) chloride should be soluble. Next, let's consider lead (II) sulfate. When we split apart this ionic compound, we have lead (II) ions and sulfate ions. Lead (II) ions do not have their own rule. Sulfate is normally soluble, but lead (II) ions are exceptions, so lead (II) sulfate is insoluble. Barium sulfide breaks apart into barium ions and sulfide ions. Barium does not have its own rule. Sulfide is normally insoluble except when paired with cations that are always soluble or when paired with calcium, strontium, or barium. So barium sulfide is soluble. We'll apply these solubility rules in writing equations for precipitation reactions. In the precipitation reactions, we add together two solutions of soluble ionic compounds and form an insoluble ionic compound, which comes out of solution as a precipitate. The cations and anions of the two soluble ionic compounds change partners. If one of the resulting ionic compounds is insoluble, the precipitation reaction occurs. For example, when potassium iodide solution is added to lead (II) nitrate solution, our potassium switches partners and pairs with nitrate, while the lead (II) ions pair with iodide. Since lead (II) iodide is insoluble, the precipitation reaction occurs. We can see this reaction here. Remember that soluble ionic compounds are strong electrolytes, so the potassium iodide and lead (II) nitrate initially have both broken apart into ions in solution. When the two solutions are mixed, the lead (II) ions combine with the iodide ions to produce a yellow solid. The potassium and nitrate ions, however, remain in solution. If, however, an insoluble ionic compound is not formed, no reaction will occur. For example, if we add together potassium iodide and sodium chloride, our possible ionic compounds after switching partners are potassium chloride and sodium iodide. Since both of these are still soluble ionic compounds, no reactions occur; the ions all remain in solution. Let's practice writing the equations for precipitation reactions. In this example, we're adding together calcium acetate and ammonium chloride. First, we'll have our ions switch partners to see what compounds are formed. We can form calcium chloride and ammonium acetate. Next, we need to evaluate if either of these ionic compounds is insoluble. According to our solubility rules, both of these compounds are soluble: chloride salts are normally soluble and calcium is not an exception, and all ammonium salts and acetate salts are soluble. So ammonium acetate is also soluble. Since we produce two soluble ionic compounds, no reaction occurs; all of our ions remain in solution. In our next example, we're adding together potassium phosphate and calcium nitrate. Switching partners, we produce calcium phosphate and potassium nitrate. According to our solubility rules, phosphate salts are generally insoluble and calcium is not an exception, so calcium phosphate is our precipitate. Both potassium salts and nitrate salts are soluble, so potassium nitrate remains in solution. Next, we'll balance our equation; since we have all ionic compounds, we'll balance ions instead of atoms. If we start with one calcium phosphate, we can add a three in front of calcium nitrate to balance our calcium ions, and a two in front of potassium phosphate to balance our phosphate ions. Then we just need a six in front of potassium nitrate to balance both potassium and nitrate ions. Now we have a balanced chemical equation for our precipitation reaction.