NET IONIC EQUATIONS

Net ionic equations are useful in that they show only those chemical species participating in a chemical reaction. The key to being able to write net ionic equations is the ability to recognize monoatomic and polyatomic ions, and the solubility rules. If you are weak in these areas, a review of these concepts would be helpful before attempting to write net ionic equations. A summary of polyatomic ions and the solubility rules are given at the end of this tutorial.

Let's first start with a complete chemical equation and see how the net ionic equation is derived. Take for example the reaction of lead(II) nitrate with hydrochloric acid to form lead(II) chloride and nitric acid, shown below:

Pb(NO3)2 (aq) + 2 HCl (aq) PbCl2 (s) + 2 HNO3 (aq)

This complete equation may be rewritten in ionic form by using the solubility rules, and by recognizing strong acids. All nitrates are soluble, therefore the lead(II) nitrate will be dissociated. Both hydrochloric acid and nitric acid are strong acids and will therefore be dissociated. The lead(II) chloride, however is insoluble--all halides are soluable except silver, lead, copper(I) and mercury(I). The above equation written in dissociated form is:

Pb2+(aq) + 2 NO3-(aq) + 2 H+(aq) + 2 Cl-(aq) PbCl2 (s) + 2 H+(aq) + 2 NO3-(aq)

At this point, one may cancel out those ions which have not participated in the reaction. Notice how the nitrate ions and hydrogen ions remain unchanged on both sides of the reaction.

Pb2+(aq) + 2 NO3-(aq) + 2 H+(aq) + 2 Cl-(aq) PbCl2 (s) + 2 H+(aq) + 2 NO3-(aq)

What remains is the net ionic equation, showing only those chemical species participating in a chemical process:

Pb2+(aq) + 2 Cl-(aq) PbCl2 (s)

After a while, you should be able to predict the net ionic equation given only the reactants. For instance, suppose you had to determine the net ionic equation resulting from the reaction of solutions of sodium sulfate and barium bromide:

BaBr2 (aq) + Na2SO4 (aq) ?

One way to tackle this problem is to examine what ions are floating together in solution: Ba2+, Br-, Na+, and SO42-. We knew that barium ions and bromide ions were soluble together, but will sodium ions or sulfate ions combine with barium ions to form an insoluble compound? Barium ions and sodium ions, both being positive in charge will repel each other, so no compound is expected to form between them. On the other hand, sulfate ions and barium ions could easily form barium sulfate. Now it is just a matter of consulting the solubility rules to see if barium sulfate is soluble or insoluble. According the these rules, all sulfates are soluble except strontium, barium, and lead; silver and calcium sulfate are partially soluble. As you can see, barium sulfate will be insoluble. The sodium ions must therefore combine with bromide ions to form sodium bromide. According to the solublility rules, sodium bromide should be soluble--all compounds of sodium are soluble. Now we can write a complete balanced equation:

BaBr2 (aq) + Na2SO4 (aq) BaSO4 (s) + 2 NaBr (aq)

As before, the above equation can be rewritten showing the soluble species as ions in solution:

Ba2+(aq) + 2 Br-(aq) + 2 Na+(aq) + SO42-(aq) BaSO4 (s) + 2 Na+(aq) + 2 Br-(aq)

Next, cross out any species which have not changed on both sides of the reactions--these are the spectator ions:

Ba2+(aq) + 2 Br-(aq) + 2 Na+(aq) + SO42-(aq) BaSO4 (s) + 2 Na+(aq) + 2 Br-(aq)

What remains is the balanced, net ionic equation:

Ba2+(aq) + SO42-(aq) BaSO4 (s)

For those of you who need to review polyatomic ions, here is a list of the polyatomic ions you probably will encounter in this course:

NO3- nitrate ion , MnO4- permanganate ion , NO2- nitrite ion , CrO42- chromate ion , C2H3O2- acetate ion , Cr2O72- dichromate ion , HSO4- hydrogen sulfate ion or bisulfate ion, PO43- phosphate ion HCO3- hydrogen carbonate ion or bicarbonate ion , HPO42- hydrogen phosphate ion, SO42- sulfate ion , H2PO4- dihydrogen phosphate ion, SO32- sulfite ion, SCN- thiocyanate ion , CO32- carbonate ion , CN- cyanide ion, OH- hydroxide ion, S2O32- thiosulfate ion, H3O+ hydronium ion, NH4+ ammonium ion

The following solubility rules were taken from "Laboratory Experiments in General Chemistry", Alan Pribula, Scientific American Books,1989.

1. All salts of Na+, K+, and NH4+ are soluble.

2. All salts of NO3-, ClO3-, and C2H3O2- are soluble.
(AgC2H3O3 is partially soluble.)

3. All salts of halides and SCN- are soluble except those of Ag +, Cu+, Pb2+, and Hg22+. (HgI2 is also insoluble. PbCl2 is soluble in hot water.)

4. All salts of SO42- are soluble except for BaSO4 , PbSO4 , and SrSO4 . (CaSO4 , Ag2SO4 and Hg2S4 are partially soluble.

5. All salts of CO32-, PO43-,SiO32-, C2O42-, and SO32- are insoluble, except for those of Na+, K+, and NH4+. (MgC2O4 and Cr2(C2O4)3 are also soluble.)

6. All oxides and hydroxides are insoluble except for those of Na+, K+, and Ba2+. (Ca(OH)2 is partially soluble.)

7. All salts of S2- and insoluble except for those of Group IA and IIa elements and of NH4+.

While determining net ionic equations, cite aloud the solubility rule you are applying--this reinforces the rules and you won't forget them!