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Net Ionic Equations
Net Ionic Equations
Making Sense of Chemical Reactions
OBJECTIVE
Students will be able to write net ionic equations from balanced molecular equations.
LEVEL
Chemistry
T E A C H E R
P A G E S
NATIONAL STANDARDS
UCP.1, UCP.2, B.3
CONNECTIONS TO AP
AP Chemistry:
III. Reactions A. Reaction Types 1. Acid-base reactions 2. Precipitation reactions 3. Oxidationreduction reactions B. Stoichiometry 1. Ionic and molecular species present in chemical systems:
net ionic equations
TIME FRAME
45 minutes
MATERIALS
periodic table
solubility rules
student white boards (optional)
TEACHER NOTES
This lesson should follow a study of the five main types of chemical reactions. Having students write
net ionic equations will enhance their ability to understand “what is really happening in the beaker” type
questions. The student lecture notes can provide the basis for your lecture over this material. Guide
students through the proper steps of writing net ionic equations using several examples until the students
feel confident. It often helps to demonstrate the example chemical reactions. This is easily
accomplished by mixing the stated solutions in a test tube or beaker and asking students to explain in
chemical terms what has taken place. Did the solution change color? Did a gas form? (Your
demonstration solutions may need to be fairly concentrated in some cases so that enough gas is formed
to actually cause bubbling.) Did a precipitate form? Follow this lesson with The Eight Solution Problem
lab to reinforce observing chemical reactions and give students more practice writing net ionic
equations.
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Train students to write the state symbols for each species. Doing so will reduce the chance of incorrect
spectator cancellations. For example, if zinc metal reacts with hydrochloric acid, the following net ionic
equation results:
Zn(s) + 2 H+(aq) → Zn2+(aq) + H2(g)
Students often cancel out the zinc terms even though one is a solid metal while the other is an ion in
solution. Requiring students to write the state symbols makes it a bit more evident that the two terms are
not identical and therefore do not cancel.
Students should memorize the strong acids and the strong bases. An abbreviated table of the most
common solubility rules is presented within this lesson. It is strongly recommended that your students
begin memorizing the solubility rules during the first year of chemistry. The solubility rules must be
memorized for the AP* exam.
Answers to the examples from the notes: Demonstrate these reactions whenever possible as you
guide students through the example reactions. Recognition of state symbols is much easier when
students actually see the reaction take place. Example 1 is worked out on the student pages. The
answers to examples 2−5 follow.
P A G E S
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)
•
Balanced net ionic equation:
Mg(s) + 2 H+(aq) → Mg2+(aq) + H2(g)
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T E A C H E R
Continue writing net ionic reactions throughout the year when explaining the chemical reactions
associated with demonstrations as well as during pre- and post-lab discussions to continually reinforce
this concept.
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Net Ionic Equations
T E A C H E R
P A G E S
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) → Ca2+(aq) + 2 Cl−(aq) + Mg2+(aq) + 2 NO3−(aq)
•
Balanced net ionic equation:
No net ionic equation—everything cancels so there is no driving force—this was just a physical
combination of two different solutions with no chemical reaction taking place.
ANSWERS TO THE CONCLUSION QUESTIONS
1. Solutions of lead(II) nitrate and lithium chloride are mixed.
• Pb(NO3)2(aq) + 2 LiCl(aq) → 2 LiNO3(aq) + PbCl2(aq)
• 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) → Cu2+(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)
• 2 KClO3(s) → 2 KCl(s) + 3 O2(g)
• 2 KClO3(s) → 2 KCl(s) + 3 O2(g)
NOTE: Solids and gases do not ionize. All three reactions are therefore the same. Students do not
need to write three identical equations to receive full credit.
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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)
• 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 in order 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) + Mg2+(aq) + 2 Br−(aq) → Mg2+(aq) + 2 Cl−(aq) + Br2(ℓ)
• Cl2(g) + 2 Br−(aq) → 2 Cl−(aq) + Br2(ℓ)
NOTE: All reactants and products are either gases or molecular compounds, so this is the net ionic
equation.
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)
P A G E 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) + Ba2+(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) + CO32−(aq) → SnCO3(s) + 2 Li+(aq) + 2 F−(aq)
• Sn2+(aq) + CO32−(aq) → SnCO3(s)
NOTE: Carbonate was not listed directly in the solubility rules so it is considered insoluble.
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T E A C H E R
6. Methane gas is burned in the presence of oxygen gas.
• CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g)
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Net Ionic Equations
Net Ionic Equations
Making Sense of Chemical Reactions
Now that you have mastered balancing chemical equations it is time to focus only on the reacting
species. 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 it 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.
MATERIALS
periodic table
solubility rules
student white boards (optional)
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 in order for the reaction to occur since 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 while all others are written as
undissociated molecules.
How do you know if a substance will extensively ionize in solution? All strong electrolytes will ionize
extensively in solution. Electrolytes are comprised of three classes of compounds—strong acids, strong
bases, and soluble salts. Using the information that follows you should be able to determine whether or
not to ionize a particular substance.
1. Strong Acids: HCl, HBr, HI, H2SO4, HNO3, and HClO4
• Notice HF is missing from this list, so it is classified as a weak acid. Why? The fluoride atom is
very small and as a result it is highly attracted to the hydrogen atom. As a result, it does not
dissociate completely in aqueous solution.
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2. Strong Bases: Hydroxides of group IA and IIA, but…
• Be and Mg are classified as weak since they do not dissociate completely. This is due to the fact
that they are small atoms and as a result 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 since they are not as soluble as IA
hydroxides. For these IIA hydroxides, to the little extent that they do dissolve, they dissociate
100%.
3. Soluble Salts: Use the solubility rules presented in Table 1 to determine whether or not a
compound containing a metal and non-metal will dissolve in aqueous solution and thus behave as
a strong electrolyte.
Table 1: Solubility Rules
Always Soluble if in a Compound
Except With
NO3−, Group IA, NH4+, C2H3O2−, ClO4−, ClO3−
No Exceptions
Cl−, Br−, I−
Pb2+, Ag+, Hg22+
SO4
Pb2+, Ag+, Hg22+
Ca2+, Sr2+, Ba2+
2−
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 won’t always be correct, but
will cover most of the situations you will encounter.)
A few other important points:
• Gases, pure liquids, and solids are non-electrolytes — do not ionize
• H2CO3(aq) decomposes into H2O(ℓ) and CO2(g)
• NH4OH(aq) decomposes into H2O(ℓ) and NH3(g)
• If a driving force is absent, no gas forms, no precipitate forms nor does any molecular substance
such as water form. If you encounter such a situation, write “NO REACTION”.
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)
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Net Ionic Equations
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 equation above: Na+ and NO3−
Example 2
Solutions of iron(III) nitrate and potassium hydroxide are mixed.
Balanced formula equation:
Total ionic equation:
Balanced net ionic equation:
Example 3
Magnesium ribbon reacts with hydrochloric acid.
Balanced formula equation:
Total ionic equation:
Balanced net ionic equation:
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Example 4
Solutions of acetic acid and lithium bicarbonate are mixed.
Balanced formula equation:
Total ionic equation:
Balanced net ionic equation:
Example 5
Solutions of magnesium chloride and calcium nitrate are mixed.
Balanced formula equation:
Total ionic equation:
Balanced net ionic equation:
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Net Ionic Equations
Name______________________________________
Period _____________________________________
Net Ionic Equations
Making Sense of Chemical Reactions
CONCLUSION QUESTIONS
Using your own paper, 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 in order 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.
3. Solid potassium chlorate decomposes upon heating.
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.
8. Solutions of sodium bicarbonate and hydrochloric acid are mixed.
9. Solutions of ammonium perchlorate and barium hydroxide are mixed.
10. Solutions of tin(II) fluoride and lithium carbonate are mixed.
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