Ionic Compounds in Solution

In our previous discussion of solution formation, we only talked about molecular compounds. However, solutions of ionic compounds will play a very important role in the rest of this lesson, so we will now expand on that discussion to talk about what happens when an ionic compound dissolves, and why some do and some don't.

Formation of Ionic Solutions

You should be familiar from CH 104 with the structure of ionic compounds: they form lattices with regular, repeating patterns of ions in three dimensions. On their own, they are extremely "clumpy" - even a small crystal of NaCl can be thought of as a clump of trillions of Na⁺ and Cl^- ions, connected to each other by strong ionic bonds. How does that structure come apart to form a homogeneous mixture with no clumps.

If you look closely, you can see that the oxygen atoms in water point toward the positive sodium ions, while the hydrogen atoms point toward the negative chloride ions. This places the negative partial charge on oxygen toward the sodium cation and vice versa for hydrogen. This type of interaction creates a very strong intermolecular force, called an ion-dipole force. The strength of these IMFs, and the fact that a large number of them can be formed for each ion, helps compensate for the tremendous energy needed to break the ionic bonds in NaCl.

As mentioned above, NaCl is an ionic compound that dissolves in water, while CaCO₃ is not. The factors that make one compound soluble and another insoluble in water are complex, relating to the sizes and charges of the ions, among other factors. You will now learn to apply a set of rules for determining whether a given ionic compound is soluble in water.

Solubility Rules

Every ionic compound is composed of a cation and an anion, as you should remember from CH 104. To determine if a particular compound is soluble, you identify the cation and anion that make it up, locate one or both of them on the table of rules below, and apply the rule as you read it.

Let's illustrate how to use the table with the two compounds we've encountered, NaCl and CaCO₃.

NaCl is composed of the Na⁺ and Cl^- ions; as you can see, both these ions are listed as ions that generally form soluble compounds. You should always check the exceptions column, but in this case there are no exceptions listed for Na⁺, and while Cl^- has some exceptions, sodium is not one of them. So NaCl is soluble, as we have observed.

CaCO₃ is composed of the Ca²⁺ and CO₃^2- ions. While Ca²⁺ does not appear on the table, CO₃^2- is listed as an ion that generally forms insoluble compounds. There are exceptions, such as in sodium carbonate, potassium carbonate, etc. but calcium is not one of them. Thus, CaCO₃ is insoluble, once again consistent with our observations.

Practice determining the solubility of ionic compounds now with the following examples. Don't forget the names of ions, especially the difference between sulfide and sulfate.

sodium hydroxide

iron (III) sulfide

magnesium sulfate

lead (II) chloride

barium nitrate

aluminum phosphate

The answers to these practice problems are given below. For more practice, you can do the practice problems in your lab workbook.

sodium hydroxide - soluble

iron (III) sulfide - insoluble

magnesium sulfate - soluble

lead (II) chloride - insoluble

barium nitrate - soluble

aluminum phosphate - insoluble