Net Ionic Reactions
Model 1 – Net Ionic Reactions.
Net ionic reactions are frequently used when strong electrolytes react in solution to form nonelectrolytes or weak electrolytes. These equations let you focus on the chemical reactions that are
occurring, and can emphasize general chemical trends.
The reaction between strong acids and strong bases is one example:
Reaction:
HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l)
Complete: H+ (aq) + Cl- (aq) + Na+ (aq) + OH- (aq)  Na+ (aq) + Cl- (aq) + H2O (l)
Net:
H+ (aq) + OH- (aq)  H2O (ℓ)
The formation of a precipitate from solutions of strong electrolytes is another:
Reaction: 2Na3PO4 (aq) + 3Ba(NO3)2 (aq)  6NaNO3 (aq) + Ba3(PO4)2 (s)
Complete:
6Na+ (aq) + 2PO43- (aq) + 3Ba2+ (aq) + 2NO3- (aq)  Ba3(PO4)2 (s) + 6Na+ (aq) +6 NO3- (aq)
Net:
3Ba2+ (aq) + 2PO43- (aq)  Ba3(PO4)2 (s)
Critical Thinking Questions
1. Verify that all of the reactions in Model 1 are balanced.
2. Is the following statement true or false? Explain.
If the same chemical shows up on both sides of a chemical reaction, it doesn’t react.
3. Do the moles of the chemical species that are crossed out in the reactions change?
Information
The net ionic reaction for a strong acid reacting with a strong base is:
H+ (aq) + OH- (aq)  H2O (ℓ)
Critical Thinking Question
4. Write balanced reactions for the following strong acid-strong base reactions. Verify that the net
ionic equation given above is correct.
a) Hydrochloric acid and barium hydroxide
b) Sulfuric acid (1st proton) and potassium hydroxide
c) Hydrobromic acid and sodium hydroxide
5. Do the concentrations of the spectator ions in these reactions increase, decrease, or stay the same?
Information – Solubility Rules
The solubility rules for ionic compounds are given in Table 10.9 (pg. 329) of McQuarrie, etal.
1. Most alkali metal (Group 1) and ammonium salts are soluble.
2. Most nitrates, perchlorates, and acetates are soluble.
3. Most halides (Group 17) are soluble, except silver, lead, and mercury (I) halides.
4. Most sulfates are soluble, except sulfates of calcium, barium, and strontium.
5. Most carbonates, chromates, sulfides, oxides, phosphates, and hydroxides are
insoluble, except hydroxides of Group 2.
Insoluble compounds will form precipitates from the ions in solution. For example, when potassium
iodide is mixed with lead(II) nitrate, solid lead(II) iodide forms:
Molecular Reaction:
2KI (aq) + Pb(NO3)2 (aq)  2KNO3 (aq) + PbI2 (aq)
Complete Ionic Reaction:
2K+ (aq) + 2I- (aq) + Pb2+ (aq) + 2NO3- (aq)  PbI2 (s) + 2K+ (aq) + 2NO3- (aq)
Net Ionic Reaction:
Pb2+ (aq) + 2I- (aq)  PbI2 (s)
Critical Thinking Questions
6. Precipitates form when the following aqueous solutions are mixed. Write the balanced chemical
and net ionic reactions for each. Label all phases.
a)
Sodium chloride and silver chloride:
b)
Ammonium chromate and iron(III) nitrate:
c)
Potassium phosphate and copper(II) nitrate
d)
Copper(II) iodide and ammonium carbonate
Model 2 – Putting it all together.
When 50.00 mL of 0.100 M sodium chloride is mixed with 50.00 mL of 0.100 M lead(II) nitrate,
solid lead(II) chloride (278.1 g/mole) forms a precipitate.
2NaCl (aq) + Pb(NO3)2(aq)  2NaNO3 (aq) + PbCl2 (s)
Before mixing:
After mixing:
0.00500 moles Na+
0.00500 moles Cl0.00500 moles Pb2+
0.0100 moles NO3-
0.0050 moles Na+
0 moles Cl0.00500 – 0.00250 moles
= 0.00250 moles Pb2+
0.0100 moles NO30.00250 moles PbCl2 or 0.695 g
Critical Thinking Questions.
7. Write the net ionic equation for this reaction.
8. Verify the moles of each reactant before mixing.
9. Verify the moles of all the species, and the mass of solid, after mixing.
10. What are the molar concentrations of the ions remaining in solution after mixing?
Exercises
1. Complete and balance each of the following acid-base neutralization reactions. Name the
products, and write the net ionic equations.
a) HBr (aq) + KOH (aq) 
b) HNO3 (aq) + Ba(OH)2 (aq) 
2. Determine if a precipitate forms for the following. If one does, write balanced chemical and net
ionic reactions:
a) NiCl2 (aq) + NH4NO3 (aq) 
b) FeCl2 (aq) + Na2S (aq) 
c) MgSO4 (aq) + AgNO3 (aq) 
3. How many milliliters of 0.100 M HNO3 are needed to neutralize each of the following solutions?
a) 45.0 mL 0.667 M KOH
b) 58.5 mL 0.0100 M Al(OH)3
c) 34.7 mL 0.775 M NaOH
4. Calculate the mass of MgCO3 precipitated by mixing 10.0 mL of a 0.200 M Na2CO3 solution with
5.00 mL of 0.0500 M Mg(NO3)2 solution.
Problems
1. The solubility of magnesium hydroxide (Mg(OH)2) in water is 9.0 × 10-4 g/100.0 mL. What
volume of 0.00100 M HCl is needed to neutralize 1.00 L of a saturated Mg(OH)2 solution?
2. Phosphate can be removed from drinking water supplies by treating the water with Ca(OH)2. How
much Ca(OH)2 is required to remove 90% of the PO43- from 4.5 × 106 L of drinking water containing
25 mg/L PO43-?
5 Ca(OH)2 (aq) + 3 PO43- (aq)  Ca5(PO4)3OH (s) + 9 OH- (aq)