Ch. 11 Chemical
Reactions
11.1 Describing
Chemical Reactions
I. Equation Basics
• A.
Fe (s) + O2 (g)  Fe2O3 (s)
Products
Reactants
• B. Skeleton equation: does not tell amounts of
each part
• C. Symbols for states of matter:
• solid (s)
gas (g)
liquid (l)
aqueous (aq)
Substance dissolved in water
• D. Catalyst: substance that speeds up rate of reaction
II. Balancing
Equations
• A. Atoms are never lost
or gained in a reaction,
they are just rearranged.
• B. Rules:
– 1. Only use coefficients (#’s in front of substances)
– 2. Must have same # of each element on each side
– 3. Equation must be reduced
III. Examples
• A.
N2 (g) + O2 (g)  N2O5 (g)
• B. __MnO2 + HCl  MnCl2 + H2O + Cl2
• C.
FeCl3 + Ca(OH)2 Fe(OH)3 + CaCl2
11.2 Types of
Chemical Reactions
I. Combination Rxns
• A. Two or more substances combine to one
• B. Na (s) + Cl2 (g) --> NaCl (s)
+
(
)

• C. If metal combines with non-metal, ionic
compound produced
• D. If non-metals combined, covalent compound
produced
II. Decomposition Rxns
• A. Single compound broken into two or more
products
Heat
• B. HgO (s)  Hg (l) + O2 (g)
Means
Heat
Added
• C. Most decomp. rxns require energy (heat) as
a catalyst
• ***Demo: H2O2 (l)  H2O (l) + O2 (g)
• ***Demo:
NI3 
N2 +
I2
***
III. Single-Replacement
• A. Atoms of one element replace atoms of a
second element in a compound
• B. Mg + Zn(NO3)2  Mg(NO3)2 + Zn
+

+
• C. Reactivity of metals determines whether one
atom will replace another
• D. Mg more reactive than Zn, removes Nitrate
***Demo Thermite Rxn***
IV. Double Replacement
• A. Two atoms switch places (often forming solid)
• B. K2CO3 (aq) + BaCl2 (aq)  2 KCl (aq) + BaCO3 (s)
• *** Demo: Silver Nitrate and Magnesium Chloride***
V. Combustion Rxns
• A. Element or compound reacts with O2, usually
forming energy
• B. Mg (s) + O2 (g)  MgO (s)
• C. If compound has C and H, products usually
H2O and CO2
• D. CxHy + O2  CO2 + H2O
***Demo: Nitrocellulose***
11.3 Reactions in
Aqueous Solution
I. Net Ionic Equations
• A. Equation showing only particles that take
part in reaction
• B. Process: 1. Start with full equation
AgNO3 (aq) + Na2S (aq)  Ag2S (s) + NaNO3 (aq)
• C. 2. Separate ions in aqueous form
Ag+(aq) + NO3-(aq) + Na+ (aq) + S2-(aq) 
Ag2S (s) + Na+(aq) + NO3- (aq)
• D. 3. Cross off ions appearing on both sides
of reaction as aqueous
Ag+(aq) + NO3-(aq) + Na+(aq) + S2-(aq) 
Ag2S (s) + Na+(aq) + NO3- (aq)
• E. 4. Balance everything left to get…
Net Ionic equation:
2 Ag+ (aq) + S2-(aq)  Ag2S (s)
II. Try This
• Write the balanced net ionic equation for the following:
Step 1
Zn (s) + HCl (aq)  ZnCl2 (aq) + H2 (g)
Step 2
Zn (s) + H+ (aq) + Cl- (aq)  Zn2+ (aq) + Cl- (aq) + H2 (g)
Step 3
Zn (s) + 2 H+ (aq) 
Step 4
Zn2+ (aq) +
H2 (g)
III. Precipitates
• A. Solids formed when two aqueous mixtures form
insoluble (“non-dissolvable”) compound
• B. Precipitate formation based on ion solubility rules
Solubility Rules:
• Always Soluble: Alkali metals (1st column), NH4+, NO3-,
ClO3-, ClO4-, C2H3O2• Mostly Soluble: Cl-, Br-, I- (except Ag+, Pb2+, Hg22+)
F- (except Ca2+, Ba2+, Sr2+, Pb2+, Mg2+)
SO42- (except Ca2+, Ba2+, Sr2+, Pb2+)
• Mostly Insoluble: O2-, OH- (except w/ alkali metals, NH4+,
Ca2+, Sr2+, Ba2+ somewhat soluble)
CO32-, PO43-, S2-, SO32-, C2O42-, CrO42- (except w/ alkali
metals, NH4+)