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Honors Chemistry I
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Chapter 8 Study Guide
Important:
1) Read pages 260-288…everything in this packet builds off of this information, so don’t ignore it!
2) Complete all practice problems, vocabulary, and conceptual sections in this packet.
Section 8.1: Describing Chemical Reactions
Vocabulary:
1) Chemical reaction
2) Chemical equation
Evidence of a chemical change:
Physical change examples:
1) Changes of state (e.g., solidliquid)
2) Breaking of a substance into smaller substance (e.g., rocks smashed into gravel or sand)
Important: Physical changes do NOT involve chemical reactions!
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About chemical reactions:
1) A chemical reaction will either release energy or absorb energy.
a. Bond breaking ALWAYS REQUIRES ENERGY! The overall process of bond breaking
(which requires energy) and bond formation (which releases energy) determines whether a
reaction absorbs energy (endothermic change) or releases energy (exothermic change)
b. Example: CH4 + O2  CO2 + H2O + ENERGY
i. This is an EXOTHERMIC reaction because it releases energy as heat
c. Example: N2O4 + ENERGY  2NO2
i. This is an ENDOTHERMIC reaction because energy must be absorbed by N2O4 in order
to form NO2
2) Molecules and atoms must come into contact for them to chemically react.
a. As an example: A match won’t ignite by itself…but when struck against a phosphorus striking
surface, the KClO3 in the match head can ignite!
b. Remember that atoms and molecules possess kinetic energy (energy of MOTION), and when
they “slam” into one another a chemical reaction can occur.
Writing chemical equations:
1) A chemical equation indicates the things reacting, as well as the amount of things reacting.
a. Think of a chemical equation as being like a baking recipe.
2) Chemical equations also indicate the state of the products and reactants (see the table below)
Components of a chemical equation
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Examples:
1) For each chemical equation given, write a sentence that describes the reaction. Be sure to include
physical states and reaction conditions:
a. Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g)
b. NaOH (aq) + HCl(aq)  NaCl(aq) + H2O(l)
c. CaCl2(aq) + Na2CO3(aq)  CaCO3(s) + 2NaCl (aq)
d. CaCO3(s) + ENERGY  CaO(s) + CO2(g)
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Section 8.2: Balancing Chemical Equations
Law of conservation of mass:
Reactions rearrange atoms:


A balanced chemical equation correctly shows the law of conservation of mass in action!
o You will use coefficients to balance chemical reactions. Coefficients are placed in front of the
atom/molecule/compound present in a chemical equation.
Never change the subscripts involved in a chemical equation!! If you change a subscript (e.g., if you
change H2O to H2O2), you’ve completely changed the substances involved in the chemical reaction!
Water is NOT THE SAME THING as hydrogen peroxide!
Balancing chemical reactions:
1) Be sure to identify the reactants (on the left side) and the products (on the right side)
2) Count the total number of atoms of each element on each side of the equation
a. If there is an imbalance of an atom, start planning a strategy to bring balance to the equation
b. You will use coefficients to balance equations.
3) It’s best to focus on balancing one element at a time, usually starting with those that appear in only one
product and one reactant.
4) Generally, it’s smart to worry about balancing hydrogen and oxygen until the end, because they can
sometimes appear in multiple products and reactants.
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A Detailed Example:
Consider the chemical equation below. What can be done to balance the chemical equation and be sure that the
law of conservation of mass is satisfied?
____ Na(s) + ____ H2O(l)  ____ NaOH (aq) + ____ H2(g)
Examples:
Balance all of the following chemical equations:
1) ____ P4 + ____O2  ____ P2O5
2) ____ C3H8 + ____ O2  ____ CO2 + ____H2O
3) ____Ca2Si + ____ Cl2  ____CaCl2 + ____SiCl4
4) Ethane (C2H6) burns in the presence of oxygen gas to form carbon dioxide gas and water vapor. Write a
balanced chemical equation for the reaction described.
5) Magnesium metal and water vapor react to form aqueous magnesium hydroxide and hydrogen gas.
Write a balanced chemical equation for the reaction described.
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6) Iron metal and aluminum oxide form when aluminum metal reacts with iron (III) oxide. Write a
balanced chemical equation for the reaction described.
7) Solid silicon reacts with carbon dioxide gas to form silicon carbide (SiC) and silicon dioxide (sand).
Write a balanced chemical equation for the reaction described.
8) _____ Al(NO3)3 + _____ (NH4)3PO4 → _____AlPO4 + _____ NH4NO3
9) _____ AgF + _____ CaCl2 → _____ AgCl + _____ CaF2
10) _____ C2H4O2 + _____ O2 → _____ CO2 + _____ H2O
11) _____ AgNO3 + _____ Ga → _____ Ag + _____ Ga(NO3)3
12) _____ Li2SO4 + _____ K3PO4 → _____ Li3PO4 + _____ K2SO4
13) _____ N2 + _____ H2 → _____ NH3
14) _____ KMnO4 → _____ K2MnO4 + _____ MnO2 + _____ O2
15) _____ S8 + _____ O2 → _____ SO2
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16) _____ H2O2 → _____ O2 + _____ H2O
17) _____ HgCl2 + _____AgNO3  _____Hg(NO3)2 + _____AgCl
18) _____ Al + _____ Hg(CH3COO)2  _____ Al(CH3COO)3 + Hg
19) _____ Pb(NO3)2 + _____ Na2CrO4  _____ PbCrO4 + _____ NaNO3
20) _____ (NH4)3PO4 + _____Mg(CH3COO)2  _____ Mg3(PO4)2 + _____ NH4CH3COO
21) _____ MoO3 + _____ H2SO4 + _____ Zn  _____ Mo2O3 + _____ H2O + _____ ZnSO4
22) Calcium phosphate and water are produced when calcium hydroxide reacts with phosphoric acid
(H3PO4). Write a balanced chemical equation for this reaction.
23) Ammonia (NH3) reacts with oxygen gas to produce nitrogen monoxide and water vapor. Write a
balanced equation for this reaction.
24) Chlorine gas reacts with potassium bromide to form potassium chloride and bromine gas. Write a
balanced equation for this reaction.
Please note: Additional practice can be found on page 274 (see questions 5-6) and page 293 (tons of practice).
Please feel free to try any and all of these examples, and check with me if you’d like to check your answers.
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Section 8.3: Classifying Chemical Reactions
Types of reactions:
1) Combustion Reaction: An oxidation reaction of an organic (carbon-containing) compound, in which
heat energy is released and the products carbon dioxide and water made.
Example: _____C8H18 + _____O2  _____CO2 + _____H2O + ENERGY
2) Synthesis (Combination) Reaction: A reaction in which two or more substances combine to form a
new compound.
Example: _____ K(s) + _____ Cl2(g)  _____ KCl(s)
3) Decomposition Reaction: A reaction in which a single compound breaks down to form two or more
simpler substances.
Example: _____H2CO3 (aq)  _____CO2(g) + _____H2O(l)
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4) Single-Displacement Reaction: A reaction in which a chemically-similar substance “replaces” the
substance that its similar to in a chemical compound.
Example: _____Al(s) + _____Pb(NO3)2(aq)  _____ Pb(s) + ______Al(NO3)3(aq)
We use the “activity series” to predict whether or not a single-displacement reaction will occur or not.
Steps to using the activity series:
 Identify the potential reactants in the table shown
 Locate the two elements that might be involved in the single-replacement.
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 If the atom to replace the atom included in the compound is HIGHER up the activity series list, then the
reaction will occur.
 If the atom to replace the atom included in the compound is LOWER on the activity series list, no
reaction will occur.
 General tips:
o Alkali metals (group 1) almost always will react to form compounds, and they are usually never
alone after a chemical reaction
o Group 2 metals (alkaline earth metals) are less reactive than alkali metals, but MORE reactive
than transition metals.
Examples: For each of the situations below, write a balanced chemical reaction if the reaction happens. If the
reaction does not happen, write “no reaction.”
 Aluminum is dipped into an aqueous solution of zinc (II) nitrate
 A chunk of solid sodium is placed into cold water
 Gold is added to a solution of calcium chloride
 Lead is placed into an iron (III) nitrate solution
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 Zinc metal is added to a solution of copper (II) sulfate
 Magnesium ribbon is dipped into a nickel (II) chloride solution
5) Double-Displacement Reaction: A reaction in which a gas, a solid precipitate, or a molecular
compound forms from the apparent exchange of atoms or ions between two compounds. This reaction
resembles a “trading places” situation.
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Balancing and classifying reactions:
For each of the problems provided below, create a balanced chemical reaction and classify the reaction as either
combustion, synthesis, combination, or single-displacement (no double displacement reactions are included).
1) _____ H2 + _____ Cl2  _____ HCl
Type: __________________
2) _____Mg + ______O2  _____MgO
Type: __________________
3) _____KI + _____ Br2  _____ KBr + _____ I2
Type: __________________
4) _____C2H6 + _____ O2  _____ CO2 + _____H2O
Type: __________________
5) _____ Ca + _____ H2O  _____ Ca(OH)2 + _____ + H2
Type: __________________
6) _____KClO3  _____ KCl + _____ O2
Type: __________________
7) _____CH4 + _____ O2  _____CO2 + _____H2O
Type: __________________
8) _____ Zn + _____ HCl 
Type: __________________
9) _____C4H10 + _____O2 
Type: __________________
10) _____ BaO + _____ H2O 
Type: __________________
11) _____C6H14 + _____ O2  _____CO2 + H2O
Type: __________________
12) _____ Fe + _____ O2  _____Fe2O3
Type: __________________
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Practice Problem Set A: For each problem given below, balance the chemical equation and indicate the type of
reaction.
1) Chlorine gas reacts with aqueous sodium bromide to produce aqueous sodium chloride and bromine gas.
2) A solid sample of calcium oxide reacts with liquid water to produce an aqueous solution of calcium
hydroxide
3) A solid sample of calcium chlorate forms solid calcium chloride and oxygen gas when heated
4) Aqueous silver (I) nitrate reacts with aqueous potassium sulfate to form aqueous silver (I) sulfate and
aqueous potassium nitrate.
5) _____ Zn(s) + _____ CuBr2(aq)  _____ZnBr2 (aq) + _____Cu(s)
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Practice Problem Set B: Predict whether a reaction would occur in each case when the materials are brought
together. If a reaction would occur, complete and balance the chemical equation. If no reaction would occur,
write “no reaction occurs.”
1) Ag(s) + H2O(l)
2) Mg(s) + Cu(NO3)2
3) Al(s) + O2(g)
4) H2SO4(aq) + KOH (aq)
Practice Set C: Predict the products, write a balanced chemical equation, and identify the type of reaction for
each reaction below:
1) _____Zn + _____CuSO4 
2) _____BaCl2 + _____Na2SO4 
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3) _____Zn + _____F2 
4) _____C5H10 + _____O2 
MIXED REVIEW: Balance the following equations
1) _____Ca(OH)2 (aq) + _____H3PO4(aq) _____ H2O(l) + _____Ca3(PO4)2 (s)
2) _____Al(OH)3(s) + _____HCl(aq)  _____AlCl3(aq) + _____H2O(l)
3) _____ KO2(s) + _____ H2O(l)  _____ KOH(aq) + _____ O2(g) + _____H2O2 (aq)
4) _____ Eu(s) + _____ HF(g)  _____EuF3(s) + _____H2(g)
5) _____C12H22O11(s) + _____O2(g)  _____CO2(g) + _____H2O(g)
6) Solid Iron (III) sulfide reactions with gaseous hydrogen chloride to form solid Iron (III) chloride and
hydrogen sulfide gas
7) Liquid carbon disulfide reacts with ammonia gas to produce hydrogen sulfide gas and solid ammonium
thiocyanate
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Section 8.4: Writing Net Ionic Equations
Ionic Equations:
 When ionic compounds dissolve in water, the anion and cation separate from one another and spread out
uniformly throughout the aqueous solution.
 What does this mean? Simply, whenever we write that there’s an aqueous substance present (e.g.,
KI(aq)), it actually exists as shown below:
Examples: NaI(aq) =
Pb(NO3)2 (aq) =
(NH4)3PO4(aq) =
Total ionic equations:
Write the complete ionic equation for the reaction below. FYI, this is the exact process performed in our lab
activity that produced lead (II) iodide (aka “the solid yellow stuff”)
Step 1: Balance the equation
_____KI(aq) + _____Pb(NO3)2(aq)  _____PbI2(s) + _____KNO3(aq)
Step 2: Separate all of the aqueous ionic substances into their ions:
At this point, pay close attention to the fact that you have some of the same ions on each side of the equation.
Just like in basic algebra, when you have the same thing on each side of an equation, it cancels out! We call
these ions that “cancel out” on both sides of the equation “SPECTATOR IONS.” Spectator ions are ions present
in a solution that don’t actually participate in the chemical reaction…just like spectators at a sporting event!
Re-Write Step 2 from above and cancel all spectator ions present:
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Examples: Write both the TOTAL ionic equation and the NET ionic equation for all of the following chemical
reactions. You should start by writing a balanced chemical equation. Write the TOTAL ionic equation first, and
then rewrite it showing the cancellation of spectator ions to arrive at the NET ionic equation.
1) Ca(OH)2 (aq) + HCl(aq)  CaCl2(aq) + H2O(l)
2) Br2(l) + NaI(aq)  NaBr(aq) + I2(s)
3) Mg(s) + AgNO3(aq)  Ag(s) + Mg(NO3)2(aq)
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4) AgNO3(aq) + KBr(aq)  AgBr(s) + KNO3(aq)
5) Ni(s) + Pb(NO3)2(aq)  Ni(NO3)2 (aq) + Pb(s)
6) Ca(s) + H2O(l)  Ca(OH)2(aq) + H2(g)
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Solubility Rules & Predicting Solubility:
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 “Soluble” means that it will dissolve and ionize in water
 “Insoluble” means that it will NOT dissolve and ionize in water.
o If you drop an “insoluble” substance in water, it will remain solid and you should typically be
able to see the little solid pieces rather clearly
o Important to ionic equations, insoluble compounds will NOT ionize!!!
Examples:
Label each compound as SOLUBLE or INSOLUBLE
1) NaCl
____________________________
2) Ba(NO3)2
____________________________
3) BaSO4
____________________________
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4) MgCO3
____________________________
5) (NH4)2CO3
____________________________
6) PbCl2
____________________________
7) LiClO3
____________________________
8) NH4Br
____________________________
9) AgCl
____________________________
10) NaOH
____________________________
11) NaHCO3
____________________________
12) (NH4)3PO4
____________________________
13) Ca3(PO4)2
____________________________
14) Mg(OH)2
____________________________
15) CaS
____________________________
16) Na2CrO4
____________________________
17) CaCO3
____________________________
18) PbSO4
____________________________
19) (NH4)2SO4
____________________________
20) Hg2I2
____________________________
21) KI
____________________________
22) BaBr2
____________________________
23) Li2S
____________________________
24) Ba(OH)2
____________________________
25) KOH
____________________________
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Examples: (2 additional example problems are included in test review section):
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1) Aqueous solutions of barium chloride and potassium sulfate are mixed.
a. Write the formulas for the ionic compounds that are reactants for this chemical reaction
b. What compound will precipitate (form a solid)?
c. Write a balanced chemical equation for the reaction
d. Write the TOTAL ionic equation
e. What is/are the spectator ion(s)?
f. Write the NET ionic equation
2) Aqueous solutions of Iron (III) sulfate and lithium hydroxide are mixed
a. Write the formulas for the ionic compounds that are reactants for this chemical reaction
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b. What compound will precipitate?
c. Write a balanced chemical equation for this reaction
d. Write the TOTAL ionic equation for this reaction
e. What is/are the spectator ion(s)?
f. Write the NET ionic equation for this reaction
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Test Review and Supplemental Problems
The following problems are intended to be additional practice for quizzes, tests, and exams. If you need
additional practice/examples, complete the chapter 8 problems provided in your textbook on pages 292-295.
1) Balance all of the chemical equations given below:
a. PCl5(l) + H2O(l)  H3PO4(aq) + HCl(g)
b. C(s) + CaO(s)  CaC2(s) + CO2(g)
c. FeCO3(s) + H2CO3(aq)  Fe(HCO3)2(aq)
d. Fe(s) + O2(g)  Fe2O3(s)
e. FeO(s) + O2(g)  Fe2O3 (s)
f. Cr(s) + S8(s)  Cr2S3(s)
g. NaHCO3(s)  Na2CO3(s) + CO2(g) + H2O(g)
h. Potassium reacts with water yielding potassium hydroxide and hydrogen gas
i. Solutions of silver (I) nitrate and lithium bromide are mixed. (predict the products and balance
the equation)
j. Solutions of zinc (II) sulfate and sodium phosphate are mixed. (predict the products and balance
the equation)
k. Liquid silicon tetrachloride is reacted with pure solid magnesium, producing solid silicon and
solid magnesium chloride
l. Na2SiF6(s) + Na(s)  Si(s) + NaF(s)
m. The combustion of ethanol (C2H5OH) in the presence of oxygen gas forms carbon dioxide and
water vapor.
n. Solid zinc metal reacts with aqueous hydrogen chloride to form aqueous zinc chloride and
hydrogen gas.
o. Aqueous strontium hydroxide reacts with aqueous hydrogen bromide (aka “hydrobromic acid”)
to produce liquid water and aqueous strontium bromide.
p. Aqueous solutions of lead (II) nitrate and sodium phosphate are mixed, resulting in the
precipitate formation of lead (II) phosphate with aqueous sodium nitrate as the other product
2) Iron oxide ores, commonly a mixture of FeO and Fe2O3, are given the general formula Fe3O4. They yield
elemental iron when heated to extremely high temperature with either carbon monoxide or elemental
hydrogen gas. Balanced the following equations for these two processes:
a. Fe3O4(s) + H2(g)  Fe(s) + H2O(g)
b. Fe3O4(s) + CO(g)  Fe(s) + CO2(g)
3) Chromium compounds exhibit a wide variety of bright colors. When solid ammonium dichromate—a
vivid orange compound—is ignited, a spectacular reaction occurs. Although the reaction is somewhat
more complex, let’s assume here that the products are solid chromium (III) oxide, nitrogen gas, and
water vapor. Write a balanced equation for this chemical reaction.
4) Write a balanced chemical equation for each of the following examples and classify each reaction as a
synthesis, decomposition, combustion, single displacement, or double displacement:
a. zinc chloride + ammonium sulfide
b. zinc metal + copper (II) sulfate
c. magnesium bromide + chlorine
d. aluminum oxide decomposes
e. silver nitrate + sodium chloride
f. magnesium + copper (II) nitrate
g. sodium hydroxide + sulfuric acid
h. lead (II) nitrate + potassium bromide
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i. copper + tin (IV) chloride
j. C10H22 + O2
k. potassium sulfide + iron (III) nitrate
l. zinc metal + silver nitrate
m. silver metal + oxygen gas
n. tin(IV) chloride decomposes
o. calcium hydroxide + phosphoric acid
p. magnesium iodide + chlorine gas
q. barium nitrate + sodium phosphate
r. manganese(IV) oxide decomposes
s. C11H24 + O2
t. aluminum chloride + potassium sulfide
5) For each of the reactions listed, write the balanced equation, the total ionic equation, and the net ionic
equation:
a. Aqueous potassium chloride is added to aqueous silver nitrate to form a silver chloride
precipitate and aqueous potassium nitrate
b. Aqueous potassium hydroxide is mixed with aqueous iron (III) nitrate to form a precipitate of
iron (III) hydroxide and aqueous potassium nitrate
3) Aqueous solutions of calcium chloride and sodium carbonate are mixed
a. Write the formulas for the ionic compounds that are reactants for this chemical reaction
b. What compound will precipitate?
c. Write a balanced chemical equation for this reaction
d. Write the TOTAL ionic equation for this reaction
e. What is/are the spectator ion(s)?
f. Write the NET ionic equation for this reaction
4) Aqueous solutions of silver (I) nitrate and potassium phosphate are mixed
a. Write the formulas for the ionic compounds that are reactants for this chemical reaction
b. What compound will precipitate?
c. Write a balanced chemical equation for this reaction
d. Write the TOTAL ionic equation for this reaction
e. What is/are the spectator ion(s)?
f. Write the NET ionic equation for this reaction
6) When the following solutions are mixed, what precipitate (if any) is formed? Be sure to reference the
solubility rules listed in section 8.4.
a. KNO3(aq) + BaCl2(aq)
b. Na2SO4(aq) + Pb(NO3)2(aq)
c. KOH(aq) + Fe(NO3)3(aq)
d. FeSO4(aq) + KCl(aq)
e. CaCl2(aq) + Na2SO4(aq)
f. K2S(aq) + Ni(NO3)2(aq)
g. Al(NO3)3(aq) + Ba(OH)2(aq)
7) For all of the above questions (a-g), write the complete balanced chemical equation, the total ionic
equation, and the net ionic equation.
8) Write the net ionic equations for the reaction (if any) that occurs when aqueous solutions of the
following are mixed:
a. Ammonium sulfate and barium nitrate
b. Lead (II) nitrate and sodium chloride
c. Sodium phosphate and potassium nitrate
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