# How to Balance Net Ionic Equations

These are the steps to write a balanced net ionic equation and a worked example problem.

### Steps To Balance Ionic Equations

1. Write the net ionic equation for the unbalanced reaction. If you are given a word equation to balance, you'll need to be able to identify strong electrolytes, weak electrolytes, and insoluble compounds. Strong electrolytes dissociate entirely into their ions in water. Examples of strong electrolytes are strong acids, strong bases, and soluble salts. Weak electrolytes yield very few ions in solution, so they are represented by their molecular formula (not written as ions). Water, weak acids, and weak bases are examples of weak electrolytes. The pH of a solution can cause them to dissociate, but in those situations, you'll be presented an ionic equation, not a word problem. Insoluble compounds do not dissociate into ions, so they are represented by the molecular formula. A table is provided to help you determine whether or not a chemical is soluble, but it's a good idea to memorize the solubility rules.
1. Separate the net ionic equation into the two half-reactions. This means identifying and separating the reaction into an oxidation half-reaction and a reduction half-reaction.
2. For one of the half-reactions, balance the atoms except for O and H. You want the same number of atoms of each element on each side of the equation.
3. Repeat this with the other half-reaction.
4. Add H2O to balance the O atoms. Add H+ to balance the H atoms. The atoms (mass) should balance out now.
5. Balance charge. Add e- (electrons) to one side of each half-reaction to balance charge. You may need to multiply the electrons by the two half-reactions to get the charge to balance out. It's fine to change coefficients as long as you change them on both sides of the equation.
6. Add the two half-reactions together. Inspect the final equation to make sure it is balanced. Electrons on both sides of the ionic equation must cancel out.
7. Double-check your work! Make sure there are equal numbers of each type of atom on both sides of the equation. Make sure the overall charge is the same on both sides of the ionic equation.
1. If the reaction takes place in a basic solution, add an equal number of OH- as you have H+ ions. Do this for both sides of the equation and combine H + and OH- ions to form H2O.
2. Be sure to indicate the state of each species. Indicate solid with (s), liquid for (l), gas with (g), and an aqueous solution with (aq).
1. Remember, a balanced net ionic equation only describes chemical species that participate in the reaction. Drop additional substances from the equation.

### Example

The net ionic equation for the reaction you get mixing 1 M HCl and 1 M NaOH is:

H+(aq) + OH-(aq) → H2O(l)

Even though sodium and chlorine exist in the reaction, the Cl- and Na+ ions are not written in the net ionic equation because they don't participate in the reaction.

### Solubility Rules in Aqueous Solution

 Ion Solubility Rule NO3- All nitrates are soluble. C2H3O2- All acetates are soluble except silver acetate (AgC2H3O2), which is moderately soluble. Cl-, Br-, I- All chlorides, bromides, and iodides are soluble except Ag+, Pb+, and Hg22+. PbCl2 is moderately soluble in hot water and slightly soluble in cold water. SO42- All sulfates are soluble except sulfates of Pb2+, Ba2+, Ca2+, and Sr2+. OH- All hydroxides are insoluble except those of the Group 1 elements, Ba2+, and Sr2+. Ca(OH)2 is slightly soluble. S2- All sulfides are insoluble except those of the Group 1 elements, Group 2 elements, and NH4+. Sulfides of Al3+ and Cr3+ hydrolyze and precipitate as hydroxides. Na+, K+, NH4+ Most salts of sodium-potassium and ammonium ions are soluble in water. There are some exceptions. CO32-, PO43- Carbonates and phosphates are insoluble, except those formed with Na+, K+, and NH4+. Most acid phosphates are soluble.