Chapter 4 The Major Classes of Chemical Reactions

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Chapter 4The Major Classes ofChemical Reactions

• Read/Study Chapter 4
• Memorize Tables 4.1, 4.2, 4.3
• Sample Problems 4.1 4.10
• Follow-up Problems 4.1 4.10
• End-of-Chapter Problems At least every 3rd
  problem.

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Displacement Reactions
Reaction Types
Single Displacement Single Replacement (Redox)
Formation Combination (Redox)
Decomposition (Sometimes Redox)
Other
Metathesis Double Displacement Double
Replacement (Never Redox)
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Important Definitions
Solution - A homogeneous mixture consisting of
one or more substances uniformly dispersed as
separate atoms, molecules, or ions in another
substance.
Solvent - The component of a solution that is
the dissolving medium. The solvent determines
the physical state of the solution (solid,
liquid, or gas).
Solute - The components of a solution that are
dis- solved by the medium.
Aqueous Solution - A solution wherein water is
the solvent.
When the solute and solvent are both in the same
physical state, the one in the largest quantity
is the solvent.
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Solubility - The amount of solute that will
dissolve in a given quantity of solvent at a
given temperature.
Saturated Solution - A solution that contains an
amount of solute that is equal to its solubility.
The dissolved solute in the solution is in
dynamic equilibrium with the un- dissolved solute
(the precipitate).
Dissolved solute Precipitate
The rate of dissolving is equal to the rate of
precipitation. This is a dynamic equilibrium.
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Unsaturated Solution - A solution that contains
an amount of solute that is less than its
solubility.
All of the solute present is dissolved in an
unsaturated solution.
Supersaturated Solution - A solution that
contains an amount of solute that is more than
its solubility. This is a metastable state
(without stability).
If disturbed in anyway, the excess solute will
precipitate out of solution and a saturated
solution will result.
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Practice Problem If 50 mL of water at 10oC
contain 0.000 0445 g of dissolved AgCl, (a) is
the solution unsaturated, saturated, or
supersaturated? (b) is the solution dilute or
concentrated?
The solubility of AgCl in water at 10oC is 0.000
089 g AgCl/100 mL H2O.
(0.000 0445 g AgCl/50 mL H2O)(100 mL H2O/100 mL
H2O)
0.000 089 g AgCl/100 mL H2O
(a) Saturated (b) Dilute
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Solutes in Aqueous Solutions
Non-Electrolytes
Electrolytes
Solutes that do NOT produce ions when dissolved.
Solutes that DO produce ions when dissolved.
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Properties of Solutes in Aqueous Solutions
Non-Electrolyte - A compound that does NOT
produce ions when melted or dissolved in water.
An aqueous solution of a non-electrolyte will NOT
conduct electricity.
Battery
C12H22O11 C12H22O11 C12H22O11
Smallest Particles are molecules!!
Sugar dissolved in water.
H2O
C12H22O11 (s) C12H22O11 (aq)
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Properties of Solutes in Aqueous Solutions
Electrolyte - A compound that produces ions
when melted or dissolved in water. An aqueous
solution of an electrolyte will conduct
electricity through the move- ment of the ions in
the solution.
Battery
Na Cl - Na Cl - Na Cl- Na Cl -
Smallest Particles are ions!!
Table salt dissolved in water.
NaCl (s) Na Cl -
H2O
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Properties of Solutes in Aqueous Solutions
Weak Electrolyte - A compound that produces only
a few ions when dissolved in water. An aqueous
solution of a weak electrolyte will conduct some
electricity through the movement of the ions in
the solution.
Battery
C2H3O2- H HC2H3O2 HC2H3O2
Smallest Particles are ions AND molecules!
Acetic Acid dissolved in water.
HC2H3O2 (aq) H C2H3O2 -
H2O
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Electrolytes
Acids Salts Bases
Strong Electrolyte - An electrolyte that is
completely or almost completely ionized
(dissociated) in solution.
H2O
NaCl (s) Na Cl -
H2O
HCl (g) H Cl -
Weak Electrolyte - An electrolyte that ionizes in
water only to a limited extent.
HC2H3O2 (aq) H C2H3O2 - NH3 (aq)
H2O (l) NH4 OH -
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Acid - Acids are substances that ionize in H2O to
form H3O - usually abbreviated simply as H.
(This concept was developed by Arrhenius)
HCl (aq) H2O (l) H3O Cl -
Strong Acid
HC2H3O2 (aq) H2O (l) H3O C2H3O2 -
Weak Acid
Or more simply...
HCl (aq) H Cl -
HC2H3O2 (aq) H C2H3O2 -
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Base - Bases are substances that produce OH -
in aqueous solution.
(According to Mr. Arrhenius)
NaOH (aq) Na OH -
Strong Base
NH3 (aq) H2O (l) NH4 OH -
Weak Base
CH3NH2 (aq) H2O (l) CH3NH3 OH -
Weak Base
Ba(OH)2 (aq) Ba2 2 OH -
Strong Base
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Salt - An ionic compound containing neither H
ion nor OH- ion. It is formed by the reaction of
an acid and a base in a reaction known as
acid-base neutralization.
Acid-Base Neutralization The reaction between
and acid and a base to produce water and a Salt.
NaOH (aq) HCl (aq) NaCl (aq) H2O (l)
KOH (aq) HC2H3O2 (aq) KC2H3O2
(aq) H2O (l)
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State whether the following compounds are acids,
bases, or salts KNO3 - Cr(OH)3 - NaOH -
H2SO4 - CaCl2 - NH3 - Ba3(PO4)2 -
Na2SO4 - HCl - H2O -
Salt
Base
Base
Acid
Salt
Base
Salt
Salt
Acid
Acid and Base
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Complete the following neutralization equations
HClO4 (aq) KOH (aq) H3PO4 (aq) Fe(OH)3
(s) HC2H3O2 (aq) LiOH (aq)
KClO4 (aq) H2O (l)
FePO4 (s) 3 H2O (l)
LiC2H3O2 (aq) H2O (l)
Ionic Solutions - Solutions containing ions
dispersed in some solvent.
H2O
NaCl (aq) Na Cl - KNO3 (aq) K
NO3 -
H2O
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NaCl (aq) KNO3 (aq) Na Cl -
K NO3 -
NaNO3 (aq) Na NO3 -
KCl (aq) K Cl -
NaNO3 (aq) KCl (aq) Na NO3 -
K Cl -
Even though we started with two different
salts, the two solutions have exactly the same
composi- tion. The cations and anions of the two
solutions are identical.
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Precipitation Reactions
Pb(NO3)2 (aq) 2 NaI (aq) ? 2 NaNO3 (aq)
2 PbI2 (s)
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Precipitation Reactions
Figure 4.6
2 NaI (aq) Pb(NO3)2 (aq) PbI2 (s) 2 NaNO3
(aq)
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Precipitation Reactions

• Form an insoluble compound
• Molecular Equation
• Complete Ionic Reaction
• Net Ionic Equation

Ca(NO3 )2 (aq) K2CO3 (aq) ? CaCO3 (s)
2 KNO3 (aq)
Ca2 2 NO3- 2 K CO32- ?
CaCO3 (s) 2 K 2 NO3-
Ca2 CO32- ? CaCO3 (s)
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Acid-Base Reactions

• Acid Base Salt Water

• Recall the acids and bases from previous
• discussion
• Molecular equation
• HBr (aq) KOH (aq) ? KBr (aq) H2O
  (l)
• Complete ionic equation

H Br - K OH- ? K Br -
H2O (l)
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Acid-Base Neutralization

• Net Ionic Equation

H OH- ? H2O (l)
Figure 4.8
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Acid-Base Neutralization

• Molecular Equation
• Ca(OH)2 (aq) 2 HNO3 (aq) ? Ca(NO3)2 (aq)
  2 H2O (l)
• Complete Ionic Equation
• Ca2 2 OH- 2 H 2 NO3- ?
• Ca 2 2 NO3- 2 H2O (l)
• Net Ionic Equation

2 H 2 OH- 2 H2O (l)
or H OH- H2O
(l)
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Displacement Reactions
Reaction Types
Single Displacement Single Replacement (Redox)
Formation Combination (Redox)
Decomposition (Sometimes Redox)
Other
Metathesis Double Displacement Double
Replacement (Never Redox)
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Displacement Reactions
Single Displacement
Double Displacement

• Occurs when ionic
• compounds switch
• partners and forms a
• precipitate
• gas or
• weak or non-
• electrolyte

Occurs when one element is more ACTIVE than
another
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Displacement Reactions (Single Replacement)
K Na Mg Al Zn Cr Fe Cd Co Ni Sn Pb H Cu Hg Ag Pt A
u
How would you know that Cu is more ACTIVE than Ag?
ACTIVITY SERIES
Look at the .
Figure 4-19 on Page 160.
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Displacement Reactions (Single Replacement)
K Na Mg Al Zn Cr Fe Cd Co Ni Sn Pb H Cu Hg Ag Pt A
u
Ag (s) CuCl2 (aq) ? ??
Molecular Equation
2 Ag (s) CuCl2 (aq) ? 2 AgCl (s) Cu (s)
Complete (total) Ionic Equation
2 Ag (s) Cu2 2 Cl - ? 2 AgCl (s) Cu
(s)
Net Ionic Equation
2 Ag (s) Cu2 2 Cl - ? 2 AgCl (s) Cu
(s)
Will this reaction actually occur??
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Displacement Reactions (Single Replacement)
K Na Mg Al Zn Cr Fe Cd Co Ni Sn Pb H Cu Hg Ag Pt A
u
Zn (s) FeCl3 (aq) ? ??
Molecular Equation
3 Zn (s) 2 FeCl3 (aq) ? 3 ZnCl2 (aq) 2
Fe (s)
Complete (total) Ionic Equation
3 Zn (s) 2 Fe3 6 Cl - ? 3 Zn2 6
Cl - 2 Fe (s)
Net Ionic Equation
3 Zn (s) 2 Fe3 ? 3 Zn2 2 Fe (s)
Will this reaction actually occur??
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Displacement Reactions (Single Replacement)
K Na Mg Al Zn Cr Fe Cd Co Ni Sn Pb H Cu Hg Ag Pt A
u
Au (s) HCl (aq) ? ??
Molecular Equation
2 Au (s) 6 HCl (aq) ? 2 AuCl3 (aq) 3 H2
(g)
Complete (total) Ionic Equation
2 Au (s) 6 H 6 Cl - ? 2 Au3 6 Cl -
(aq) 3 H2 (g)
Net Ionic Equation
2 Au (s) 6 H ? 2 Au3 3 H2 (g)
Will this reaction actually occur??
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Oxidation- Reduction - Redox It takes Two to
Tango
Oxidation-Reduction Reactions
1. Definition - A reaction in which electrons
are transferred from one substance to another
one substance is oxidized and the other substance
is reduced.
2. Applications - A. Batteries B.
Bleaches C. Photography D. Elements combining
to form compounds E. Elements replacing other
elements in compounds
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3. Oxidation - A Loss of electrons. 4.
Reduction - A Gain of electrons a REDUCTION in
positive charge. 5. Oxidizing Agent - The
substance in a Redox reaction that is REDUCED!
It Gains electrons. 6. Reducing Agent - The
substance in a Redox reaction that is OXIDIZED!
It Loses electrons. 7. Examples - 2 H2 (g)
O2 (g) 2 H2O (l)
Reducing Agent
Oxidizing Agent
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7. Examples (Continued) - Cu (s) 2 AgNO3
(aq) Cu(NO3)2 (aq) 2 Ag (s)
Reducing Agent
Oxidizing Agent
2 K (s) Br2 (l) 2 KBr
Reducing Agent
Oxidizing Agent
CH4 (g) 2 O2 (g) CO2 (g) 2 H2O (l)
Reducing Agent
Oxidizing Agent
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7. Examples (Continued) - CH4 (g) 2 O2 (g)
CO2 (g) H2O (l)
Reducing Agent
Oxidizing Agent
H2O2 (l) H2O2 (l) 2 H2O (l) O2 (g)
Oxidizing Agent
Reducing Agent
Autooxidation or Disproportionation Reaction
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Oxidation Number Rules
1. The oxidation number of an atom in a
free element (not in a compound!) is ALWAYS zero,
regardless of its formula.
0 0
2 H2 (g) O2 (g) --gt 2 H2O (l)
2. The oxidation number of a monatomic (simple)
ion is equal to the charge on the ion.
0 1 2 0
Fe (s) 2 Ag --gt Fe2 2 Ag (s)
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Oxidation Number Rules
3. The oxidation number of fluorine is ALWAYS -1
in its compounds.
0 0 1 -1
Cs (s) F2 (g) --gt CsF
4. The oxidation number of oxygen is -2 unless
it is bonded to fluorine or itself.
-1 2 -1 1 -1 -1 1 0 1 -2
F-O-F H-O-O-H O2 H2O
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Oxidation Number Rules
5. Hydrogen has an oxidation state of 1 in
all compounds except in metal hydrides where it
is -1.
0 0 -3 1
3 H2 (g) N2 (g) --gt 2 NH3 (g)
0 0 1 -1
H2 (g) 2 Cs (s) --gt 2 CsH (s)
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Oxidation Number Rules
6. The sum of the oxidation numbers in a
neutral compound is zero the sum for a
polyatomic ion is equal to the charge on the ion.
1 7 -2 -3 1 6 -2
NaClO4 NH4 SO42-

6 -1 3 -2 5 -2 4 -2
SF6 C2O42- NO3- N2O4
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Oxidation Number Rules

• 7. Some Common Oxidation Numbers
• Alkali Metals are ALWAYS 1
• Alkaline Earth Metals are ALWAYS 2
• Halogens are -1 except when bonded to
• oxygen or to a halogen above it in the
• Periodic Table

1 -1 0 5 -1 7 -2
NaCl Cl2 BrF5 BrO4-
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Oxidation Number Rules
8. For Main Group elements, the highest possible
oxidation state is equal to its group number.
1 -1 0 1 -2 3 -2 5 -2 7
-2
KBr Br2 BrO- BrO2-
BrO3- BrO4-
1 -2 0 4 -2 6 -2 4
-2 6 -2
Cu2S S8 SO2 SO3
SO32- SO42-
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Oxidation Number Rules
9. For Main Group elements, the lowest possible
oxidation state is equal to its group number
minus 8.
1 -1 1 -2 1 -3 1 -3
2 -2
NaCl Na2S K3P Li3N MgO
Practice Exercises
1 5 -1
KIF4 KSbF6 BrCl3 (NH4)3PO4
1 3 -1
3 -1
-3 1 5 -2
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Oxidation Numbers

• Assign oxidation numbers to each element in these
  compounds and polyatomic ions
• NaNO3
• 1 5 -2

K2Sn(OH)6 H3PO4 1 4 -2 1 You do
it!
1 5 -2

• SO32- HCO3- Cr2O72-

4 -2 14 -2 6 -2
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Balancing Redox Reactions
Oxidation Number Method Taught in your
textbook on pages 151 152. Half-Reaction
Method Based on recognizing that All redox
reactions have an oxidation half-reaction and a
reduction half-reaction. The two
half- reactions add up, of course, to a WHOLE
reaction! Both methods utilize the fact that not
only atoms have to be balanced in redox reactions
but also electrons have to be balanced.
(Pages 894 to 898)
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Half-Reaction Method

• In Acidic Solution
• Write the Net Ionic Equation (unbalanced).
• KMnO4 (aq) HCl (aq) ? MnCl2 (aq) Cl2
  (g)
• KCl (aq) H2O (l)
• K MnO4- H Cl- ? Mn2 2 Cl-
  Cl2 (g)
• K Cl- H2O (l)
• MnO4- H Cl- ? Mn2 Cl2 (g)
  H2O (l)

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Half-Reaction Method

• Find the substances that are oxidized and
  reduced.
• -1 0
• Cl- ? Cl2 (g) (Oxidization)
• 7 2
• MnO4- ? Mn2 (Reduction)
• Balance the half-reactions, first wrt atoms, then
  wrt charge.
• 2 Cl- ? Cl2 (g)
• 2 Cl- ? Cl2 (g) 2 e-

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Half-Reaction Method
MnO4- ? Mn2 4 H2O (l) 8 H MnO4-
? Mn2 4 H2O (l) 5 e- 8 H MnO4-
? Mn2 4 H2O (l)

• Multiply the half-reactions by numbers that
  equalize the numbers of electrons lost and
  gained.
• 2 Cl- ? Cl2 (g) 2 e-
• 5 e- 8 H MnO4- ? Mn2 4 H2O (l)

5
2
10 Cl- ? 5 Cl2 (g) 10
e- 10 e- 16 H 2 MnO4- ? 2 Mn2
8 H2O (l)
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Half-Reaction Method

1. Add the half-reactions together algebraically.

10 Cl- ? 5 Cl2 (g) 10 e- 10 e-
16 H 2 MnO4- ? 2 Mn2 8 H2O (l)
2 MnO4- 10 Cl- 16 H ? 5 Cl2 (g) 2
Mn2 8 H2O (l)

• Convert back to the molecular equation.
• 2 KMnO4 16 HCl ? 5 Cl2 (g) 2 MnCl2
  (aq)
• 8 H2O (l) 2 KCl (aq)

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Half-Reaction Method

• In Basic Solution
• Work it the same AS IF it was in acidic solution.

2 MnO4- 10 Cl- 16 H ? 5 Cl2 (g) 2
Mn2 8 H2O (l)

• Add enough OH- ions to both sides to balance the
• number of H ions in the equation.
• 16 OH- 2 MnO4- 10 Cl- 16 H ? 5 Cl2 (g)
• 2 Mn2 8 H2O (l) 16 OH-

8 H2O (l) 2 MnO4- 10 Cl- ? 5 Cl2 (g) 2
Mn2 16 OH-