Chapter 5
Molecular View of Reactions
in Aqueous Solutions
Chemistry: The Molecular Nature
of Matter, 6E
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Reactions in Solution
 For reaction to occur
 Reactants needs to come into physical contact
 Happens best in gas or liquid phase
 Movement occurs
Solution
 Homogeneous mixture
 2 or more components mix freely
 Molecules or ions completely intermingled
 Contains at least 2 substances
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Definitions:
Solvent
 Medium that dissolves solutes




Component present in largest amount
Can be gas, liquid, or solid
Liquids most common
Aqueous solution—water is solvent
Solute
 Substance dissolved in solvent




Solution is named by solute
Can be gas—CO2 in soda
Liquid—Ethylene glycol in antifreeze
Solid—Sugar in syrup
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Iodine Molecules in Ethanol
Crystal of solute
placed in solvent
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Solute molecules dispersed
throughout solvent
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Solutions
 May be characterized using
Concentration
 Solute-to-solvent ratio
g solute
g solvent
or
g solute
g solution
 Percent Concentration
g solute
% concentration 
100 g solution
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Relative Concentration
Dilute solution
 Small solute to solvent ratio
Ex. Eyedrops
Concentrated
solution
 Large solute to solvent
ratio
Ex. Pickle brine
 Dilute solution contains less solute per unit
volume than more concentrated solution
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Concentration
Solubility
g soluteneeded to make saturatedsolution
Solubility 
100 g solvent
 Temperature dependent
Saturated solution
 Solution in which no more solute can be dissolved
at a given temperature
Unsaturated solution
 Solution containing less solute than maximum
amount
 Able to dissolve more solute
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Solubilities of Some Common
Substances
Substance
Solubility
Formula (g/100 g water)
Sodium chloride
NaCl
Sodium
hydroxide
NaOH
Calcium
carbonate
CaCO3
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35.7 at 0°C
39.1 at 100°C
42 at 0°C
347 at 100°C
0.0015 at 25°C
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Concentrations
Supersaturated Solutions
 Contains more solute than required for saturation
at a given temperature
 Formed by careful cooling of saturated solutions
 Unstable
 Crystallize out when add seed crystal – results in
formation of solid or precipitate (ppt.)
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Preciptates
Precipitate
 Solid product formed when reaction carried out in
solutions and one product has low solubility
 Insoluble product
 Separates out of solution
Precipitation reaction
 Reaction that produces precipitate
Pb(NO3)2(aq) + 2KI(aq)  PbI2(s) + 2KNO3(aq)
 1 mol Pb(NO3)2  2 mol KI
 0.100 mol Pb(NO3)2  0.200 mol KI
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Electrolytes in Aqueous Solution
 Ionic compounds conduct electricity
 Molecular compounds don’t conduct electricity
Why?
Bright
light
No
light
Ions
present
Molecular
CuSO4 & water
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Sugar & water
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Ionic Compounds (Salts) in Water
 H2O molecules arrange themselves around ions
& remove them from lattice.
Dissociation
 Break salts apart
into ions when
enter solution
Separated ions
 Hydrated
 Conduct electricity
 Note: Polyatomic ions
remain intact
 Ex. KIO3  K+ + IO3 NaCl(s)  Na+(aq) + Cl–(aq)
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Molecular Compounds In Water
 When molecules dissolve in water
 Solute particles are surrounded by water
 Molecules are not dissociated
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Electrical Conductivity
Electrolyte
 Solutes that yield electrically conducting solutions
 Separate into ions when enter into solution
Strong electrolyte
 Electrolyte that dissociates 100% in water
 Yields aqueous solution that conducts electricity
 Good electrical conduction
 Ionic compounds
 Strong acids and bases
Ex. NaBr, KNO3, HClO4, HCl
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Electrical Conductivity
Weak electrolyte
 Aqueous solution that weakly conducts electricity
due to low ionization
 Weak acids and bases
Ex. Acetic acid (HC2H3O2), ammonia (NH3)
Non-electrolyte
 Aqueous solution that doesn’t conduct electricity
 Molecules remain intact in solution
Ex. Sugar, alcohol
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Your Turn
How many ions form on the dissociation of
Na3PO4?
A. 1
B. 2
C. 3
D. 4
E. 8
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Your Turn
How many ions form on the dissociation of
Al2(SO4)3?
A. 2
B. 3
C. 5
D. 9
E. 14
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Equations for Dissociation
Reactions
 Ionic compound dissolves to form hydrated ions
 Hydrated = surrounded by water molecules
 In chemical equations, hydrated ions are
indicated by
 Symbol (aq) after each ions
 Ions are written separately
KBr(s)  K+(aq) + Br(aq)
Mg(HCO3)2(s)  Mg2+(aq) + 2HCO3(aq)
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Learning Check
Write the equations that illustrate the dissociation
of the following salts:
Na3PO4(aq) →
3 Na+(aq) + PO43(aq)
Al2(SO4)3(aq) →
CaCl2(aq) →
2 Al3+(aq) + 3 SO42(aq)
Ca2+(aq) + 2 Cl(aq)
Ca(MnO4)2(aq) →
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Ca2+(aq) + 2 MnO4(aq)
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Equations of Ionic Reactions
 Consider the reaction of Pb(NO3)2 with KI
Pb2+
NO3–
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K+
I–
PbI2(s)
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Equations of Ionic Reactions
 When two soluble ionic solutions are mixed,
sometimes an insoluble solid forms.
 Three types of equations used to describe
1. Molecular Equation
 Substances listed as complete formulas
2. Ionic Equation
 All soluble substances broken into ions
3. Net Ionic Equation
 Only lists ions that actually take part in
reaction
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Equations of Ionic Reactions
1. Molecular Equation
 Complete formulas for all reactants and products
 Formulas written with ions together
 Does not indicate presence of ions
 Gives identities of all compounds
 Good for planning experiments
Ex.
Pb(NO3)2(aq) + 2KI(aq)  PbI2(s) + 2KNO3(aq)
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Equations of Ionic Reactions
2. Ionic Equation
 Emphasizes the reaction between ions
 All strong electrolytes dissociate into ions
 Used to visualize what is actually occurring in
solution
 Insoluble solids written together as they don’t
dissociate to any appreciable extent
Ex.
Pb2+(aq) + 2NO3(aq) + 2K+(aq) + 2I(aq) 
PbI2(s) + 2K+(aq) + 2NO3(aq)
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Equations of Ionic Reactions
Spectator Ions
 Ions that don’t take part in reaction
 They hang around and watch
 K+ & NO3 in our example
3. Net Ionic Equation
 Eliminate all spectator ions
 Emphasizes the actual reaction
 Focus on chemical change that occurs
Ex. Pb2+(aq) + 2I(aq)  PbI2(s)
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Net Ionic Equations
 Many ways to make PbI2
1.Pb(NO3)2(aq) + 2KI(aq)  PbI2(s) + 2KNO3(aq)
2.Pb(C2H3O2)2(aq) + 2NH4I(aq)  PbI2(s) +
2NH4C2H3O2(aq)
 Different starting reagents
 Same net ionic equation
 Pb2+(aq) + 2I(aq)  PbI2(s)
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Converting Molecular Equations to
Ionic Equations
Strong electrolytes exist as dissociated ions in
solution
Strategy
1. Identify strong electrolytes
2. Use subscript coefficients to determine total
number of each type of ion
3. Separate ions in all strong electrolytes
4. Show states as recorded in molecular equations
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Learning Check: Convert Molecular
to Ionic Equations:
Write the correct ionic equation for each:
Pb(NO3)2(aq) + 2NH4IO3(aq) → Pb(IO3)2(s) + 2NH4NO3(aq)
Pb2+(aq) + 2NO3–(aq) + 2NH4+(aq) + 2IO3–(aq) →
Pb(IO3)2(s) + 2NH4+(aq) + 2NO3–(aq)
2NaCl (aq) + Hg2(NO3)2 (aq) → 2NaNO3 (aq) + Hg2Cl2 (s)
2Na+(aq) + 2Cl–(aq) + Hg22+(aq) + 2NO3–(aq) → 2Na+(aq)
+ 2NO3–(aq) + Hg2Cl2(s)
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Your Turn
Consider the following reaction :
Na2SO4(aq) + BaCl2(aq) → 2NaCl(aq) + BaSO4(s)
Which is the correct ionic equation?
A. 2Na+(aq) + SO42–(aq) + Ba2+(aq) + Cl22–(aq) → 2Na+(aq)
+ 2Cl–(aq) + BaSO4(s)
B. 2Na+(aq) + SO42–(aq) + Ba2+(aq) + 2Cl–(aq) → 2Na+(aq)
+ 2Cl–(aq) + BaSO4(s)
C. 2Na+(aq) + SO42–(aq) + Ba2+(aq) + Cl22–(aq) → 2Na+(aq)
+ 2Cl–(aq) + Ba2+(s) + SO42–(s)
D. Ba2+(aq) + SO42–(aq) → BaSO4(s)
E. Ba2+(aq) + SO42–(aq) → Ba2+(s) + SO42–(s)
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Converting Ionic Equations to Net
Ionic Equations
Strategy
1. Identify spectator ions
2. Eliminate from both sides
3. Rewrite equation using only ions that actually
react.
4. Show states as recorded in molecular and ionic
equations
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Learning Check: Convert Ionic
Equation to Net Ionic Equation
Write the correct net ionic equation for each.
Pb2+(aq) + 2NO3–(aq) + 2K+(aq) + 2IO3–(aq) →Pb(IO3)2(s)
+ 2K+(aq) + 2NO3–(aq)
Pb2+(aq) + 2IO3–(aq) → Pb(IO3)2(s)
2Na+(aq) + 2Cl–(aq) + Hg22+(aq) + 2NO3–(aq) → 2Na+(aq)
+ 2NO3–(aq) + Hg2Cl2(s)
2Cl–(aq) + Hg22+(aq) → Hg2Cl2(s)
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Your Turn
Consider the following molecular equation:
(NH4)2SO4(aq) + Ba(CH3CO2)2(aq) →
2NH4CH3CO2(aq) + BaSO4(s)
Which is the correct net ionic equation?
A. Ba2+(aq) + SO42–(aq) → BaSO4(s)
B. 2NH4+(aq) + 2CH3CO2–(aq) → 2NH4CH3CO2(s)
C. Ba2+(aq) + SO42–(aq) → BaSO4(aq)
D. 2NH4+(aq) + Ba2+(aq) + SO42–(aq) + 2CH3CO2–(aq)
→ 2NH4+(aq) + 2CH3CO2–(aq) + BaSO4(s)
E. 2NH4+(aq) + 2CH3CO2–(aq) → 2NH4CH3CO2(aq)
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Criteria for Balancing Ionic and
Net Ionic Equations
Material Balance
 There must be the same number of atoms of
each kind on both sides of the arrow
Electrical Balance
 The net electrical charge on the left must
equal the net electrical charge on the right
 Charge does not have to be zero
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Learning Check: Balancing Equations
for Mass & Charge
Balance Molecular Eqn. for mass
2Na3PO4(aq) + 3Pb(NO3)2(aq)  6NaNO3(aq) +
Pb3(PO4)2(s)
 Can keep polyatomic ions together when counting
Balance Ionic Eqn. for charge
6Na+(aq) + 2PO43(aq) + 3Pb2+(aq) + 6NO3(aq) 
6Na+(aq) + 6NO3(aq) + Pb3(PO4)2(s)
 Charge must add up to zero on both sides.
Net Ionic Eqn. Balanced for both mass & charge
3Pb2+(aq) + 2PO43(aq)  Pb3(PO4)2(s)
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Acids & Bases as Electrolytes
 Many common laboratory chemicals and
household products
Indicators
 Dye molecules that change color
in presence of acids or bases
Acids
 Turn blue litmus red
 Lemon juice, vinegar, H2SO4
Bases
 Turn red litmus blue
 Drano (lye, NaOH), ammonia (NH3)
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Neutralization Reaction
 Important reaction of acids and bases
 Acid reacts with base to form water and salt
(ionic compound).
Acid + base  salt + H2O
Ex. HCl(aq) + NaOH(aq)  NaCl(aq) + H2O
HBr(aq) + LiOH(aq)  LiBr(aq) + H2O
 1:1 mole ratio of acid:base gives neutral solution
Ionization reactions
 Ions form where none have been before
 Reactions of acids or bases with water
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Arrhenius
 Acid-base neutralization is
H+(aq) + OH–(aq)  H2O
 In solution, H+ attaches itself to H2O to form
H3O+ or hydronium ion in water
 H+ does not ever exist in aqueous solution
 When H3O+ reacts, it releases H+
 H+ is active ingredient
 Often use just H+ for simplicity

H2O
HCl(g ) 
 H (aq )
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
 Cl
(aq )
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Arrhenius Acid
 Substance that reacts with water to produce
the hydronium ion, H3O+

HCl(g) + H2O
Cl–(aq) + H3O+(aq)
Acid + H2O  Anion + H3O+
HA + H2O
 A– + H3O+
HC2H3O2(aq) + H2O  H3O+(aq) + C2H3O2−(aq)
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Acids Categorized by Number of H+s
Monoprotic Acids
 Furnish only one H+
HNO3(aq) + H2O  H3O+(aq) + NO3–(aq)
HC2H3O2(aq) + H2O  H3O+(aq) + C2H3O2–(aq)
Polyprotic acids
 Furnish more than one H+
Diprotic acids — furnish two H+
H2SO3(aq) + H2O  H3O+(aq) + HSO3–(aq)
HSO3–(aq) + H2O  H3O+(aq) + SO32–(aq)
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Acids Catagorized by Number of H+s
Polyprotic acids
 Triprotic acids — furnish three H+
–H+
–H+
–H+
H3PO4  H2PO4–  HPO42–  PO43–
 Stepwise equations
H3PO4(aq) + H2O  H3O+(aq) + H2PO4–(aq)
H2PO4–(aq) + H2O  H3O+(aq) + HPO42–(aq)
HPO42–(aq) + H2O  H3O+(aq) + PO43–(aq)
Net:
H3PO4(aq) + 3H2O  3H3O+(aq) + PO43–(aq)
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Acidic Anhydrides
Nonmetal Oxides
 Act as Acids
 React with water to form molecular acids that
contain hydrogen
SO3(g) + H2O  H2SO4(aq)
sulfuric acid
N2O5(g) + H2O  2HNO3(aq)
nitric acid
CO2(g) + H2O  H2CO3(aq)
carbonic acid
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Arrhenius Bases
 Ionic compounds that contain hydroxide ion,
OH–, or oxide ion, O2–.
or
 Molecular compounds that react with water to
give OH–.
1. Ionic compounds containing OH– or O2–
a. Metal Hydroxides
 Dissociate into metal & hydroxide ions
NaOH(s)  Na+(aq) + OH–(aq)
Mg(OH)2(s)  Mg2+(aq) + 2OH–(aq)
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Ionic Oxides
b. Basic Anhydrides
 Soluble metal oxides
 Undergo ionization (hydrolysis) reaction to
form hydroxide ions
 Oxide reacts with water to form metal hydroxide
CaO(s) + H2O  Ca(OH)2(aq)
O2–
H2 O
2OH–
 Then metal hydroxide dissociates in water
Ca(OH)2(aq)  Ca2+(aq) + 2OH–(aq)
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Strong vs. Weak Electrolyte
HCl(aq)
CH3COOH(aq)
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NH3(aq)
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Strong Acids
HClO4(aq)
HClO3(aq)
HCl(aq)
HBr(aq)
HI(aq)
HNO3(aq)
H2SO4(aq)
perchloric acid
chloric acid
hydrochloric acid
hydrobromic acid
hydroiodic acid
nitric acid
sulfuric acid
 Dissociate completely when dissolved in water
Ex. HBr(g) + H2O  H3O+(aq) + Br–(aq)
 Good electrical conduction
 Any acid not on this list, assume weak
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Arrhenius Bases
2. Molecular Bases
 Undergo ionization (hydrolysis) reaction to form
hydroxide ions
Base + H2O  BaseH+(aq) + OH–(aq)
B + H2O  BH+(aq) + OH–(aq)
NH3(aq) + H2O  NH4+(aq) + OH–(aq)
NH3
H2O
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
NH4+ OH–
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Strong Bases
 Bases that dissociate completely in water
 Soluble metal hydroxides
 KOH(aq)  K+(aq) + OH–(aq)
 Good electrical conductors
 Behave as (aq) ionic compounds
 Common strong bases are:
 Group IA metal hydroxides
 LiOH, NaOH, KOH, RbOH, CsOH
 Group IIA metal hydroxides
 Ca(OH)2, Sr(OH)2, Ba(OH)2
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Weak Acids
 Any acid other than 7 strong acids
 Only ionize partially (<100%)
Organic acids
HC2H3O2(aq) + H2O  H3O+(aq) + C2H3O2–(aq)
Acetic Acid
Molecule,
HC2H3O2
Only this H comes off as
H+
Acetate ion, C2H3O2–
Ex.
HCO2H(aq) + H2O  H3O+(aq) + HCO2–(aq)
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Why is Acetic Acid Weak?
H2O + C2H3O2–(aq)  HC2H3O2(aq) + H3O+(aq)
H3O+(aq) + C2H3O2–(aq)  HC2H3O2(aq) + H2O
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Dynamic Equilibrium
 2 opposing reactions occurring at same rate
 Also called Chemical equilibrium
Equilibrium
 Concentrations of substances present in solution do
not change with time
Dynamic
 Both opposing reactions occur continuously
 Represented by double arrow
HC2H3O2(aq) + H2O
H3O+(aq) + C2H3O2–(aq)
Forward reaction – Forms ions
Reverse reaction – Removes ions
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Weak Bases
 Molecular bases
 Do not dissociate
 Accept H+ from water inefficiently
 Accept H+ from acids preferentially
NH3(aq) + HCl(aq)  NH4Cl(aq)
Ex.
NH3(aq) + H2O  NH4+(aq) + OH(aq)
Or for general base
B(aq) + H2O  BH+(aq) + OH(aq)
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Equilibrium for Weak Base
Forward reaction
Reverse reaction
Net is dynamic equilibrium
NH3(aq) + H2O
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NH4+(aq) + OH(aq)
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Position of Equilibrium
 Extent of completion
 Depends on electrolyte
 Weak electrolyte
 Strong electrolyte
 Small % ionizes
 Large % ionizes
  dominant
  dominant
 Mostly reactants
 Mostly products
 Weak acids and bases  Strong acids & bases
 Lots of back reaction
 Little back reaction
 Write eqn. as
 Write eqn. as 
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Learning Check
 Write the ionization equation for each of the
following with water:
1. Weak base methylamine, CH3NH2.
CH3NH2(aq) + H2O
CH3NH3+(aq) + OH–(aq)
2. Weak acid nitrous acid, HNO2.
HNO2(aq) + H2O
H3O+(aq) + NO2–(aq)
3. Strong acid chloric acid, HClO3.
HClO3(aq) + H2O  H3O+(aq) + ClO3–(aq)
4. Strong base strontium hydroxide, Sr(OH)2.
Sr(OH)2(aq)  Sr2+(aq) + 2 OH–(aq)
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Your Turn
Which of the following is a weak acid?
A. HCl
B. HNO3
C. HClO4
D. HC2H3O2
E. H2SO4
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Your Turn
Which of the following is not a strong base?
A. NaOH
B. CH3NH2
C. Cs2O
D. Ba(OH)2
E. CaO
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Your Turn
Which of the following is not a product of the
reaction:
NH3(aq) +HCN(aq) ?
A.CN–(aq)
B.NH4+(aq)
C.NH3CN(s)
D.H2O
E.HCN
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Acid—Base Nomenclature
 System for naming acids and bases
Acids
 Hydrogen compounds of non-metals = binary
acids
 Hydrogen compounds of oxoanions = Oxoacids
 Naming acid salts
Bases
 Metal Hydroxides and oxides = ionic
 Molecular = molecular names
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Naming Acids
A. Binary Acids — hydrogen + nonmetal




Take molecular name
Drop –gen from H name
Merge hydro– with nonmetal name
Replace –ide with –ic acid
Name of Molecular
compound
Name of Aqueous
Binary Acid
HCl(g) hydrogen chloride HCl(aq) hydrochloric acid
H2S(g) hydrogen sulfide
Jespersen/Brady/Hyslop
H2S(aq) hydrosulfuric
acid
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Naming Acids
B. Oxo Acids
 Acids with hydrogen, oxygen and another nonmetal
element
 Most of the polyatomic ions in Table 3.5
 To name:
 Based on parent oxoanion name
 Take parent ion name
 Anion ends in –ate change to –ic (more O's)
 Anion ends in –ite change to–ous (less O's)
 End name with acid to indicate H+
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Oxoacids (Aqueous)
Named according to the anion suffix
 Anion ends in -ite, acid name is -ous acid
 Anion ends in -ate, acid name is -ic acid
Name of Parent
Oxoanion
nitrate
NO3
Name of Oxoacid
HNO3
nitric acid
SO42
sulfate
H2SO4
sulfuric acid
ClO2
chlorite
HClO2
chlorous acid
PO32
phosphite
H2PO3
phosphorous acid
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Learning Check: Name Each
Aqueous Acid
 HNO2
 nitrous acid
 HCN
 hydrocyanic acid
 HClO4
 perchloric acid
 HF
 hydrofluoric acid
 H2CO3
 carbonic acid
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Your Turn
Which of the following is the correct name for
HClO4 (aq)?
A. chloric acid
B. hydrochloric acid
C. perchloric acid
D. hypochlorous acid
E. chlorous acid
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Your Turn
Which of the following is the correct name for
H2SO3(aq)?
A. sulfuric acid
B. sulfurous acid
C. hydrosulfuric acid
D. hydrosulfurous acid
E. hydrogen sulfite acid
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Acid Salts
If polyprotic acids are neutralized stepwise
 Can halt neutralization before all H+’s are removed
 Must specify # of H's that remain on salt
Acid salt
 Ion containing H+ and anion
 Contains anion capable of furnishing additional
hydrogen ions
H2SO4(aq) + KOH(aq)  KHSO4(aq) + H2O(ℓ)
acid salt
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Naming Acid Salts—Polyprotic
 Must specify number of hydrogens still attached
to the anion
 Can be neutralized by additional base
Ex. Na2HPO4
sodium hydrogen phosphate
NaH2PO4
sodium dihydrogen phosphate
KHSO4
potassium hydrogen sulfate
 Some acid salts have common names
 NaHCO3
sodium hydrogen carbonate
or sodium bicarbonate
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C. Naming Bases
Oxides & Hydroxides
 Ionic compounds
 Named like ionic compounds
 Ca(OH)2
 Li2O
calcium hydroxide
lithium oxide
Molecular Bases
 Named like molecules




NH3
CH3NH2
(CH3)2NH
(CH3)3N
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ammonia
methylamine
dimethylamine
trimethylamine
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