NET IONIC EQUATIONS
Chemical reactions that occur spontaneously do so as a result of several types of driving forces. Writing the net ionic equation for a reaction often makes evident which driving force caused the reaction. The most common driving forces for chemical reactions are formation of a precipitate, formation of a molecular compound (such as water), and formation of a gas.
PURPOSE
In this activity, you will write balanced formula equations, balanced total ionic equations, and balanced net ionic equations to explain chemical reactions.
CLASS NOTES
Net ionic equations represent only the species that are actually reacting in a chemical reaction. The “parts” of the equation that are not shown in the net ionic equation are known as the spectator ions. Spectator ions do just that; they “spectate” as opposed to participate in a chemical reaction. Spectator ions must be present initially for the reaction to occur because compounds are neutral; however, they are not directly involved in the reaction.
The type of equations that you have become familiar with thus far are known as balanced formula equations. In balanced formula equations, all species are included. Each compound is represented by its correct chemical formula, and coefficients are used to balance the equation so that it obeys the law of conservation of matter. In a total ionic equation, substances that ionize extensively in solution are written as separate ions whereas all others are written as undissociated molecules.
STRONG ACIDS
• HCl, HBr, HI, H2SO4, HNO3, and HClO4
• Notice HF is missing from this list—it is classified as a weak acid. Why? The fluoride atom is very small, and it is highly attracted to the hydrogen atom. As a result, it does not dissociate completely in aqueous solution.
STRONG BASES
• Hydroxides of Group 1 and Group 2, but…
• Be and Mg are classified as weak because they do not dissociate completely. This is due to the fact that they are small atoms and are highly attracted to the hydroxide ion. Therefore, they do not dissociate completely in aqueous solution, and are considered weak bases.
• Ba(OH)2, Sr(OH)2, and Ca(OH)2 are marginally strong bases because they are not as soluble as Group 1 hydroxides. For these Group 2 hydroxides, to the little extent that they do dissolve they dissociate 100%.
SOLUBLE SALTS
Use the solubility rules presented in Table 1 to determine whether or not a compound containing a metal and nonmetal will dissolve in aqueous solution and thus behave as a strong electrolyte.
If a substance is not addressed by one of the three rules listed in Table 1, assume the substance is insoluble or it
is a weak electrolyte and does not ionize in solution. (This will not always be correct but will cover most of the situations you will encounter.)
A few other important points:
• Gases, pure liquids, and solids are nonelectrolytes—they do not ionize. • H2CO3(aq) decomposes into H2O(ℓ) and CO2(g).
• NH4OH(aq) decomposes into H2O(ℓ) and NH3(g).
Example 1:
Solutions of sodium chloride and lead(II) nitrate are mixed. First, write the balanced
formula equation.
Balanced formula equation: Write the correct chemical formula for each compound and use coefficients to balance the equation.
2 NaCl(aq) + Pb(NO3)2(aq) → 2 NaNO3(aq) + PbCl2(s)
Total ionic equation: Determine which substances are strong electrolytes and write them in ionic form.
2 Na+(aq) + 2 Cl−(aq) + Pb2+(aq) + 2 NO3 −(aq) → 2 Na +
(aq) + 2 NO3 −(aq) + PbCl2(s)
Cancel ions that are common to both sides of the equation. These are the spectator ions.
2 Na+
(aq) + 2 Cl−(aq) + Pb2+
(aq) + 2 NO3 −
(aq) → 2 Na+(aq) + 2 NO3 −
(aq) + PbCl2(s)
Balanced net ionic equation: Rewrite the equation, focusing on the reacting species.
2 Cl−(aq) + Pb2+(aq) → PbCl2(s)
Identify the spectator ions in the previous equation:
Na+ and NO3 −
Example 2:
Solutions of iron (III) nitrate and potassium hydroxide are mixed. Balanced formula equation:
Total ionic equation:
Example 3:
Magnesium ribbon reacts with hydrochloric acid. Balanced formula equation:
Total ionic equation:
Balanced net ionic equation:
Example 4:
Solutions of acetic acid and lithium bicarbonate are mixed. Balanced formula equation:
Total ionic equation:
Example 5:
Solutions of magnesium chloride and calcium nitrate are mixed. Balanced formula equation:
Total ionic equation:
Balanced net ionic equation:
CONCLUSION QUESTIONS
Write the balanced formula equation, the balanced total ionic equation and, finally, the balanced net ionic
equation for each of the following chemical reactions. You must write all three balanced equations to receive full credit. Be sure to include state symbols for each component and correct charges for ions where appropriate.
1. Solutions of lead(II) nitrate and lithium chloride are mixed.
2. Copper metal is placed into a solution of silver nitrate.
4. Solid sodium metal is placed into distilled water.
5. Chlorine gas is bubbled into a solution of magnesium bromide.
6. Methane gas is burned in the presence of oxygen gas.
7. Solutions of silver acetate and barium chloride are mixed.
9. Solutions of ammonium perchlorate and barium hydroxide are mixed.
Net Ionic Equations
(Answer KEY)
ABOUT THIS LESSON
This guided lecture is intended to teach students the basic procedure for writing net ionic reactions. As students become more comfortable identifying the species involved in the reactions, chemistry may seem more
meaningful. This lesson is often included in a unit on chemical equations but could easily be included in a solutions unit.
CLASS NOTES
Example 2:
Solutions of iron(III) nitrate and potassium hydroxide are mixed.
Balanced formula equation:
Fe(NO3)3(aq) + 3 KOH(aq) → 3 KNO3(aq) + Fe(OH)3(s)
Total ionic equation:
Fe3+(aq) + 3 NO3 −
(aq) + 3 K+(aq) + 3 OH−(aq) → 3 K+(aq) + 3 NO3 −
(aq) + Fe(OH)3(s)
Balanced net ionic equation:
Fe3+
(aq) + 3 OH−(aq) → Fe(OH)3(s)
Example 3:
Magnesium ribbon reacts with hydrochloric acid. Balanced formula equation:
Mg(s) + 2 HCl(aq) → MgCl2(aq) + H2(g)
Total ionic equation:
Mg(s) + 2 H+(aq) + 2 Cl−(aq) → Mg2+(aq) + 2 Cl−(aq) + H2(g)
Mg(s) + 2 H+(aq) → Mg2+(aq) + H2(g)
Example 4:
Solutions of acetic acid and lithium bicarbonate are mixed.
Balanced formula equation:
HC2H3O2(aq) + LiHCO3(aq) → LiC2H3O2(aq) + H2CO3(aq)
Total ionic equation:
HC2H3O2(aq) + Li +
(aq) + HCO3
−
(aq) → Li+(aq) + C2H3O2 −
(aq) + H2O(ℓ) + CO2(g)
Balanced net ionic equation:
HC2H3O2(aq) + HCO3 −
(aq) → H2O(ℓ) + CO2(g) + C2H3O2 −
(aq)
Example 5:
Solutions of magnesium chloride and calcium nitrate are mixed.
Balanced formula equation:
MgCl2(aq) + Ca(NO3)2(aq) → CaCl2(aq) + Mg(NO3)2(aq)
Total ionic equation:
Mg2+(aq) + 2 Cl−(aq) + Ca2+(aq) + 2 NO3 −(aq) → Ca 2+
(aq) + 2 Cl−(aq) + Mg2+(aq) + 2 NO3 −(aq)
Balanced net ionic equation:
CONCLUSION QUESTIONS
1. Solutions of lead(II) nitrate and lithium chloride are mixed.
• Pb(NO3)2(aq) + 2 LiCl(aq) → 2 LiNO3(aq) + PbCl2(s)
• • Pb2+(aq) + 2 NO3
−
(aq) + 2 Li+(aq) + 2 Cl−(aq) → 2 Li+(aq) + 2 NO3 −
(aq) + PbCl2(s)
• Pb2+
(aq) + 2 Cl−(aq) → PbCl2(s)
2. Copper metal is placed into a solution of silver nitrate.
• Cu(s) + 2 AgNO3(aq) → Cu(NO3)2(aq) + 2 Ag(s)
• Cu(s) + 2 Ag+(aq) + 2 NO3 −(aq) → Cu
2+
(aq) + 2 NO3 −(aq) + 2 Ag(s)
• Cu(s) + 2 Ag+(aq) → Cu2+(aq) + 2 Ag(s)
3. Solid potassium chlorate decomposes upon heating.
• 2 KClO3(s) → 2 KCl(s) + 3 O2(g)
• Note: Solids and gases do not ionize; therefore, all three reactions are the same. Students do not need to
write three identical equations to receive full credit.
4. Solid sodium metal is placed into distilled water.
• 2 Na(s) + 2 H2O(ℓ) → 2 NaOH(aq) + H2(g)
• 2 Na(s) + 2 H2O(ℓ) → 2 Na
+
(aq) + 2 OH−(aq) + H2(g)
• Note: There are no spectator ions in this equation; therefore, students need not repeat the second equation
to receive full credit.
5. Chlorine gas is bubbled into a solution of magnesium bromide.
• Cl2(g) + MgBr2(aq) → MgCl2(aq) + Br2(ℓ)
• Cl2(g) + Mg
2+
(aq) + 2 Br−(aq) → Mg2+(aq) + 2 Cl−(aq) + Br2(ℓ)
• Cl2(g) + 2 Br−(aq) → 2 Cl−(aq) + Br2(ℓ)
6. Methane gas is burned in the presence of oxygen gas.
• CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g)
• Note: All reactants and products are either gases or molecular compounds, so this is the net ionic
CONCLUSION QUESTIONS (CONTINUED)
7. Solutions of silver acetate and barium chloride are mixed.
• 2 AgC2H3O2(aq) + BaCl2(aq) → 2 AgCl(s) + Ba(C2H3O2)2(aq)
• 2 Ag+(aq) + 2 C2H3O2
−
(aq) + Ba2+(aq) + 2 Cl−(aq) → 2 AgCl(s) + Ba2+(aq) + 2 C2H3O2 −
(aq)
• Ag+(aq) + Cl−(aq) →AgCl(s)
8. Solutions of sodium bicarbonate and hydrochloric acid are mixed.
• NaHCO3(aq) + HCl(aq) → NaCl(aq) + H2CO3(aq)
• Na+(aq) + HCO3
−
(aq) + H+(aq) + Cl−(aq) → Na+(aq) + Cl−(aq) + H2O(ℓ) + CO2(g)
• HCO3
−
(aq) + H+(aq) → H2O(ℓ) + CO2(g)
• Note: Carbonic acid quickly decomposes into liquid water and carbon dioxide gas.
9. Solutions of ammonium perchlorate and barium hydroxide are mixed.
• 2 NH4ClO4(aq) + Ba(OH)2(aq) → 2 NH4OH(aq) + Ba(ClO4)2(aq)
• 2 NH4
+
(aq) + 2 ClO4 −
(aq) + Ba2+(aq) + 2 OH−(aq) → 2 NH3(g) + 2 H2O(ℓ) + Ba 2+
(aq) + 2 ClO4 −
(aq)
• NH4
+
(aq) + OH−(aq) → NH3(g) + H2O(ℓ)
• Note: Ammonium hydroxide decomposes into ammonia gas and liquid water.
10. Solutions of tin(II) fluoride and lithium carbonate are mixed.
• SnF2(aq) + Li2CO3(aq) → SnCO3(s) + 2 LiF(aq)
• • Sn2+(aq) + 2 F−(aq) + 2 Li+(aq) + CO3
2−
(aq) → SnCO3(s) + 2 Li +
(aq) + 2 F−(aq)
• Sn2+(aq) + CO3
2−
(aq) → SnCO3(s)
