Chemistry, The Central Science, 10th edition
Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten
Unit 4 (Chapter 4):
Aqueous Reactions &
Solution Stoichiometry
John D. Bookstaver
St. Charles Community College
St. Peters, MO
 2006, Prentice Hall, Inc.
Salts:
Ionic Solids: (metal-nonmetal)
dissociate (dissolve) by separation into ions
Electrolytes:
ions in solution that
conduct electricity
Non
Weak
C11H22O11
CH3OH
H2O
NO
ions
only
molecules
Strong
CH3COOH
HNO2
NH3
SOME
ions
NaOH
HNO3
KCl
ALL
ions
partially
completely
HW
ionize p.159 #33 dissociate
Electrolytes: Strong, Weak, or Non?
(ions conduct electricity)
metal-nonmetal
Compound
Ionic
nonmetals (Covalent)
Molecular
Acid
(H____)
Not Acid
STRONG
KBr
STRONG
WEAK
CaI2
NON
(6)
(& NH3)
FeCl3
C11H22O11
NaOH HCl, HBr, HI CH3COOH C2H5OH
HNO3
HNO2
Ca(OH)2
H2O
HF
(strong bases) H2SO4
HClO4
HW p.157-159 #1,2,4,5,38
Precipitation Reactions
Double Replacement:
2 (aq) +
(aq)  2
(precipitate)
(aq) +
precipitate:
insoluble ionic compound
(as predicted by solubility rules)
Pb2+I–
( )
Solubility Rules
ALWAYS Soluble Ions:
+, K+, etc.
Na
*
+
NH
4
*
–
NO
3
*
HCO3–
C2H3O2– (CH3COO–)
ClO3–, ClO4–
group I (alkali metals)
ammonium
nitrate
bicarbonate
acetate (ethanoate)
chlorate, perchlorate
WS Solubility & NIE’s #1
Common Precipitates form with:
Ag+, Pb2+, Hg2+
(halogens)
OH–
(hydroxide)
CO32–
(carbonate)
examples
AgCl, PbI2
Cu(OH)2
CaCO3
H
H
H
–
O
O
Cl
+
H
Cl
H
H
HCl + H2O  H3O+ + Cl–
 Acid: proton (H+) donor
 Base: proton (H+) acceptor
NH3 + H2O  NH4+ + OH–
H
H
H
N
H
O
H
H
N
H
H
–
+
H
O
H
Strength of Acids and Bases
STRONG (complete dissociation)
HA(aq) + H2O(l)
H3O+(aq) + A–(aq)
B(aq) + H2O(l)
BH+(aq) + OH–(aq)
WEAK (partial dissociation)
HA(aq) + H2O(l)
H3O+(aq) + A–(aq)
B(aq) + H2O(l)
BH+(aq) + OH–(aq)
Only 6 strong acids:
Strong Acids:
• Nitric (HNO3)
HI + H2O  H3O+ + I– • Sulfuric (H SO )
2
4
proton (H+) donors
• Hydrochloric (HCl)
• Hydrobromic (HBr)
• Hydroiodic (HI)
• Perchloric (HClO4)
The strong bases are
Strong Bases:
soluble hydroxides
–) of…
(OH
–
+
OH + H3O  H2O + H2O
• Group 1 (Li,Na,K)
+
proton (H ) acceptors
• CBS (Ca, Ba, Sr)
Mg(OH)2 not as soluble
ase
QUIZ!!!
(at the bell)
Electrolytes: Strong, Weak, or Non?
metal-nonmetal
Compound
nonmetals (Covalent)
Ionic
STRONG
Molecular
Acid
(H____)
STRONG
(6)
WEAK
(& NH3)
Not Acid
NON
Acid-Base Neutralization Reactions
strong
acid
(H+A–)
strong
base
(M+OH–)
ionic
compound
(M+A–)
ACID + BASE
SALT + WATER
+
Cl
H
Na
O
water
H 2O
(HOH)
H
HCl(aq) + NaOH(aq)
Na
–
Cl
H
O
H
NaCl(aq) + H2O(l)
HW p.158 #40a
Molecular Equation
The molecular equation lists the reactants
and products in their molecular form.
AgNO3(aq) + KCl(aq)  AgCl(s) + KNO3(aq)
Ionic Equation
• In the ionic equation all strong electrolytes
(strong acids, strong bases, and soluble ionic salts)
are dissociated into their ions.
• This more accurately reflects the species that are
found in the reaction mixture.
AgNO3(aq) + KCl(aq)  AgCl(s) + KNO3(aq)
Ag+(aq) + NO3–(aq) + K+(aq) + Cl–(aq) 
AgCl(s) + K+(aq) + NO3–(aq)
Net Ionic Equation
• Cross out Spectator Ions that do not change
(same state & same charge) from the left side of
the equation to the right.
• The only species left are those things that react
(change) during the course of the reaction.
Ag+(aq) + NO3–(aq) + K+(aq) + Cl–(aq) 
AgCl(s) + K+(aq) + NO3–(aq)
NIE:
Ag+(aq) + Cl–(aq)  AgCl(s)
Balanced Net Ionic Equations
comp – diss – cross – net – bal
1. Write a complete molecular equation.
2. Dissociate all strong electrolytes (aq) .
(solubility rules)
3. Cross out spectators
(same charge & state)
4. Write the net ionic equation with the
species that remain and balance it.
Balanced Net Ionic Equations
comp – diss – cross – net – bal
+
2–
2+
–
+
–
1) (NH4)2SO4 + Ba(NO3)2 → BaSO4 + NH4NO3
Ba2+ + SO42– → BaSO4
+
–
2+
–
+ –
2) NaOH + MgBr2 → NaBr + Mg(OH)2
Mg2+ + 2 OH– → Mg(OH)2(s)
HW p.158 #21
Neutralization Reactions
When a Strong Acid reacts with a Strong Base,
the net ionic equation is…
H+ + OH–  H2O
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
H+ + Cl– + Na+ + OH–  Na+ + Cl– + H2O
Neutralization Reactions
When a Weak reacts with a Strong, the
net ionic equation is…
HX + OH–  X– + H2O
HF(aq) + KOH(aq)  KF(aq) + H2O(l)
HF + Na+ + OH–  Na+ + F– + H2O
HW p.159 #40 (finish)
Balanced Net Ionic Equations
comp – diss – cross – net – bal
+
2–
2+
–
+
–
(NH4)2SO4 + Ba(NO3)2 → BaSO4 + NH4NO3
Ba2+ + SO42– → BaSO4(s)
+
–
+ –
HF(aq) + KOH(aq)  KF(aq) + H2O(l)
HF + OH–  F– + H2O
WS Solubility & NIE’s #2
Gas-Forming Reactions
H2 Demo
(M0)
(H+)
Single Rep: Metal + Acid
+
2–
Ex: Zn(s) + H2SO4(aq)
NIE: Zn(s) + 2 H+(aq)
(M+)
(gas)
Metal Ion + H2
2+ 2–
ZnSO4(aq) + H2(g)
Zn2+(aq) + H2(g)
(gas)
H2O(l) + CO2(g)
(CO32–)
CO2 Demo (H+)
Double Rep: Acid + Carbonate
Salt + H2CO3(aq)
(or Bicarbonate)
(decomposes
HW p. 159 #43
(HCO3–)
immediately)
Ex: HCl(aq) + CaCO3(s)
CaCl2(aq) + H2O(l) + CO2(g)
NIE: 2 H+(aq) + CaCO3(s)
Ca2+(aq) + H2O(l) + CO2(g)
CH3COOH + NaHCO3  CH3COONa + H2O + CO2
Oxidation-Reduction Reactions
(REDOX)
video clip
(One
cannot
occur
without
the other)
LEO
says
GER
Oxidation Numbers
Is it a redox reaction? To find out…
1) assign oxidation numbers* (or oxidation states)
to each element in a reaction.
2) check if any oxidation states changed
(↓ reduced , ↑ oxidized)
*oxidation numbers of elements describe
electrons that would be lost or gained IF the
compound was 100% ionic.
*charges of ions show electrons transferred
IN an ionic compound
Assigning Oxidation Numbers
1. All elements are 0. (all compounds are 0)
2. Monatomic ion is its charge. (Ex. Na+ ion)
3. Most nonmetals tend to be negative, but
some are positive in certain compounds or
ions. (Ex. SO3)




O is −2 always, but in peroxide ion is −1 (O2–2).
H is +1 with nonmetals, −1 with a metals.
F is always −1.
other halogens are −1, but can be positive,
like in oxyanions.
Ex. ClO3– or NO3– or SO42–
Oxidation Numbers
• The sum of the ox. #’s in a neutral
compound is 0.ex: NaCl = 0
• The sum of the ox. #’s in a polyatomic
ion is the charge on the ion. Ex: PO4=-3
Calculate the oxidation number of each:
Sulfur in… SO3
Chromium in… K2Cr2O7
Nitrogen in… NH4+
Cobalt in… [CoCl6]3–
Classifying REDOX Reactions
All rxns (but…NOT double replacement)
Synthesis
A + B → AB
(0
0 → +/–)
Decomposition
2→1
AB → A + B
(+/– → 0 0)
1→2
Single Replacement
Combustion
AB + C → A + CB
(+/– 0 → 0
+/–)
CxHy + O2 → CO2 + H2O
(–/+
0 → +/–
+/–)
Single Replacement (REDOX)
silver ions
oxidize
copper metal
Cu(s) + 2 Ag+(aq)  Cu2+(aq) + 2 Ag(s)
+(aq)
Cu2+(aq) + 2 Ag(s) 
Cu
(s)
+
2
Ag
X
Cannot
displace H+
from acid to
make H2(g)
increasing ease of oxidation
Activity
Series
of
Metals
Writing REDOX Reactions
Write an overall equation for the
synthesis of calcium sulfide from its
elements. Then write the RED and OX halfequations to identify the REDOX process.
0
0
+2 –2
Ca + S  CaS
OX:
Ca0  Ca+2 + 2 e–
RED: 2 e– + S0  S–2
Ca + S  CaS
Writing REDOX Reactions
Write an overall equation for the
decomposition of aluminum oxide into its
elements. Then write the RED and OX halfequations to identify the REDOX process.
+3 –2
0
0
Al2O3  Al + O2
3 ( 2 O–2  O20 + 4 e– )
OX:
6 O–2  3 O20 + 12 e–
RED: 4 ( 3 e– + Al+3  Al0 )
12 e– + 4 Al+3  4 Al0
2 Al2O3  4 Al + 3 O2
Writing REDOX Reactions
Write the net ionic equation for the reaction
of solid zinc in a solution of hydrochloric acid.
comp – diss – cross – net – bal
0
+1 –1
+2 –1
0
Mg(s) + HCl(aq)  MgCl2(aq) + H2(g)
Mg + 2 H+  Mg+2 + H2
Classify the reaction in two ways.
Single-Replacement and Redox
ox
Mg + 2 H+  Mg2+ + H2(g)
red
What is red & what is ox?
WS 5c #1-2
Solutions:
+
• homogeneous mixtures.
( same
_____ throughout)
• solvent is present in
greatest abundance.
• solute dissolved
in/by solvent
Molarity
• Molarity (M) is a measure of the
concentration of a solution.
moles of solute (mol)
Molarity (M) =
liters of solution (L)
units:
mol/L
or
–1
∙
mol L
What is the molarity of a solution with 29.2 g
of sodium chloride in 250. mL of water?
29.2 g NaCl x 1 mol NaCl = 0.500 mol NaCl = 2.00 M
58.44 g NaCl
NaCl
0.250 L
Preparing a Solution
WS #1-2 Conc. Calc’s
1-mass solute
2-add solvent, swirl to dissolve
3-add
solvent
to
mark
HW p.160 #60, 67
WS Concentration & Dilutions
#1
5.00 g NaHCO3 x
1 mol NaHCO3 x
1 L NaHCO3
=
84.01 g NaHCO3 0.100 mol NaHCO3
0.595 L
NaHCO3
#2
1.20
mol
CuSO4
159.62
g
CuSO4
x
0.275 L CuSO4 x
1 L CuSO4
1 mol CuSO4
=
52.7 g
CuSO4
Dilution M1V1 = M2V2
1-calc M1V1=M2V2
2-pipet V1 from concentrated
3-fill to mark w DI
WS #3-4 Dilutions
Dilution M1V1 = M2V2
What volume of water must be added to
prepare 2.0 L of 3.0 M CuSO4 from a
stock solution of concentration 8.0 M ?
WS Concentration & Dilutions
M1V1 = M2V2
#3
M1V1 = M2V2
(12.0 M)V1 = (1.25 M)(500 mL)
V1 = 52.1 mL (0.0521 L)
#4
M1V1 = M2V2
(2.50 M)V1 = (0.200 M)(250 mL)
V1 = 20.0 mL (0.0200 L)
Solution Stoichiometry
Rxn:
gA
A(aq) + 2 B(aq)  C + 2 D
molar
mass A
g A
1 mol A
mol A
L of A
mol A
1L
mol-to-mol
ratio
gB
molarity
A (M)
HW p. 161 #81
g B
1 mol B
mol B
1L
molar
mass B
molarity
B (M)
mol B
L of B