Chapt. 19 – Redox Reactions
19.1 Oxidation and Reduction
Section 19.1 Oxidation and Reduction
Oxidation and reduction are
complementary—as an atom is
oxidized, another atom is reduced.
• Describe the processes of oxidation and reduction.
• Identify oxidizing and reducing agents.
• Determine the oxidation number of an element in a
compound.
• Interpret redox reactions in terms of change in
oxidation state.
Section 19.1 Oxidation and Reduction
Key Concepts
• Oxidation-reduction reactions involve the transfer of
electrons from one atom to another.
• When an atom or ion is reduced, its oxidation
number is lowered. When an atom or ion is oxidized,
its oxidation number is raised.
Redox Reactions
In a redox reaction, both a reduction
and an oxidation process occurs
Redox Reactions - Oxidation
Loss of electrons
Chemical species becomes more
positively charged
Gain of oxygen
Loss of hydrogen
Redox Reactions – Reduction
Gain of electrons
Chemical species becomes more
negatively charged
Gain of hydrogen
Loss of oxygen
Redox Reactions
GER
LEO:
• Lose Electrons
Oxidation
LEO
LEO says GER
GER:
• Gain Electrons
Reduction
Oxidation and Reduction
Oxidation: loss of
electrons by substance
• Oxidation increases
positive charge
Reduction: gain of
electrons by substance
• Reduction reduces
positive charge
Oxidation of one substance always
accompanied by reduction of another
Synthesis of Sodium Chloride
Magnesium Oxidation
(Oxygen Reduction)
+

2+
2-
2+
2-

Redox Reactions
Element in one reactant oxidized while
element in another reactant reduced
Zn(s) + 2 H+(aq)  Zn2+(aq) + H2(g)
Which species oxidized? Zn(s)
Which species reduced? H+(aq)
Redox Reactions
Practical /everyday examples:
Corrosion of iron (rust formation)
Forest fire
combustion
Charcoal grill
Natural gas burning
Batteries
Production of Al metal from Al2O3 (alumina)
Metabolic processes
Redox Reactions & Reaction Types
Combustion – always redox rxn
• O2(g) + X  compound containing O
• Ox # of O from 0  value < 0
Single replacement – always redox rxn
• Element 1 + A  Element 2 + B
• El1 ox # from 0  value > 0
Double replacement – never redox rxn
• Ions swap positions but no electron xfer
Synthesis, Decomposition - varies
Redox Reactions
Metal oxid. by acids or H2O (single repl.)
Zn(s) + 2 H+(aq) 
Zn2+(aq) + H2(g)
2 Na(s) + 2 H2O(l)  2 NaOH(aq) + H2(g)
Metal-Metal Replacement (as net ionic)
Fe(s) +
+
2
Ni (aq)

+
2
Fe (aq)
+ Ni(s)
Combustion
C (s) + O2 (g)
 CO2 (g)
4 Fe (s) + 3 O2 (g)  2 Fe2O3 (s)
Oxidizing and Reducing Agents
Oxidizing agent: compound which causes
oxidation (loss of electrons); aka oxidant
• It itself is reduced
Reducing agent: compound which causes
reduction (gain of electrons); aka reductant
• It itself is oxidized
In complete chemical reaction, if oxidation
reaction occurs then reduction reaction
must also occur (electrons moved around,
never created or destroyed)
Common Oxidizing Agents
Elemental oxygen (O2(g))
Nitric acid
–
–
–
–
Oxyhalogens (ClO , ClO2 , ClO3 , ClO4 )
Permanganate ion (MnO4-)
High oxidation state metal ions (Mn+6)
Hydrogen peroxide (neutral or basic
solution)
Common Reducing Agents
Hydrogen gas
Hydrides
Carbon
Metals
Low oxidation state metal ions (Mn+2)
Hydrogen peroxide (acidic solution)
Redox Reactions
Electrons not explicitly shown in
chemical equations
Need to keep track of electrons to see if
redox reaction has occurred
If redox, want to know which atom has
lost electrons (been oxidized) and
which one has gained electrons (been
reduced)
Redox Reactions & Molecular
Compounds - Electronegativity
Relative tendency of element to oxidize
or reduce is related to a property called
electronegativity
We cover this property later in the
course – need not be concerned with it
at the present
If curious, see page 194 and discussion
on page 684
Practice
Problems 1 – 4, page 685
Redox Reactions – Oxidation Number
Way to keep track of electrons during
redox reaction
Hypothetical number assigned to
individual atom present in compound
using set of rules
Similar to rules used to form valid ionic
compounds but also apply to covalent
May be positive, negative, or zero
Determining Oxidation Numbers (p 686)
1. Ox. # of uncombined atoms = 0
2. Ox. # of monatomic ion = charge
3. Ox. # of more electronegative atom in
molecule or complex ion = same as
charge it would have if it were an ion
4. Fluorine always = -1 in compounds
5. Oxygen = -2 in covalent compounds
(except in peroxides where it = -1 and
when bonded to F where it is +2)
Determining Oxidation Numbers (p 641)
6. H = +1 in most compounds. Exceptions =
hydrides (bound to less electronegative
metal)
7. Group 1A, 2A, Al metals always have ox.
# = group # (charge taken in ionic
compound)
8. Sum ox. # = 0 in neutral compounds
9. Sum ox. # of atoms in polyatomic ion =
charge of ion
Determining Oxidation Numbers (p 641)
10. Can determine uncertain ox. #s by
using rules 8 & 9 + previous rules
Rules for Oxidation Numbers (#1)
ON always reported for individual
atoms/ions not groups of atoms/ions
For atom in elemental form, oxidation
number always zero
H2(g): oxidation # = 0 for each H atom
Cu(s): oxidation number = 0
Cl2(g): oxidation # = 0 for each Cl atom
Rules for Oxidation Numbers (#2)
Monoatomic ion, oxidation # = charge
on ion
K+ oxidation # = +1
Cl- oxidation # = -1
S2- oxidation # = -2
Rules for Oxidation Numbers (#6, #7)
Group 1A Metal Cations:
Always +1 (Na+, K+)
Group 2A Metal Cations:
Always +2 (Ca2+, Mg2+)
Hydrogen (H)
+1 when bonded to nonmetals
NH3, H2O
-1 when bonded to metals
 LiH, NaH, AlH3
Rules for Oxidation Numbers (#4, #5)
Oxygen (O)
-1 in peroxides (O22-)
H 2O 2
+ 2 when bonded to F (rare)
-2 in all other compounds
CaO, Fe2O3, Na2SO4
Fluorine (F)
always -1
Rules for Oxidation Numbers (#8, #9)
Sum of oxidation numbers of all atoms
in neutral compound is zero
• H2O
(+1 +1 -2 = 0)
Sum of oxidation numbers in
polyatomic ion equals charge on ion
• NO3-
(+5 -2 -2 -2 = -1)
Rules for Oxidation Numbers (#3)
Electronegativity rule – more
electronegative atom given same # as
ion
Will avoid giving problems that will use
this rule
Rules for Oxidation Numbers
For many compounds, can directly
apply rules to determine oxidation
number of all atoms except for one
Use “sum of charges” rules to
determine oxidation number of that last
element
Determining Oxidation Numbers
Determine oxidation state of all
elements in SO3
Which elements have specific rules?
O = -2
To find oxidation number of S:
S + 3(-2) = 0
so
S = +6
Determining Oxidation Numbers
Determine oxidation number of Mn and
O in MnO4
Which elements have specific rules?
O = -2
To find oxidation number of Mn:
Mn + 4(-2) = -1
so
Mn = +7
Oxidation Numbers
Determine oxidation number of each
element in Al(OH)3
Which elements have specific rules?
Ionic compound!
Al = its charge, O = -2, H = +1
To find oxidation number of Al:
Determine its charge
Al = +3
Assign oxidation numbers for all
elements in sodium bicarbonate
-2, -2, -2
+1
NaHCO3
+1
(zero charge on
molecule)
1 + 1 + C + 3*(-2) = 0
C = +4
Oxidation Numbers
Main group elements, when monatomic, usually adopt
a oxidation number equal to their preferred ionic charge
Oxidation Numbers of Elements
May have several different oxidation states
Max: A group number; Min: A group number -8
Metals: never negative
Oxidation Numbers
+1 +7 -2
KMnO4
+5 -2
+1 +6 -2
H2SO4
+6 -2
N2O5
SO42-
+1 -2
+3 -1
H2S
ClF3
Redox Reactions
2 Zn(s) + O2(g)  2 ZnO(s)
0
0
+2 -2
oxidation
numbers
Ox. # of Zn increases
oxidized
reducing agent
Ox. # of O decreases
reduced
oxidizing agent
Redox Reactions
Did redox reaction occur?
If so, oxidizing agent? Reducing agent?
2 H2 (g) + O2(g)
0
0

2 H2O (g)
+1 -2
Redox reaction occurred
H oxidized
H2 = reducing agent
O reduced
O2 = oxidizing agent
Redox Reactions
Did redox reaction occur?
Zn (s) + 2 H+ (aq)
0
+1

Zn2+ (aq) + H2(g)
+2
Zn (s) oxidized
Zn = reducing agent
H+ reduced
H+ = oxidizing agent
0
Redox Reaction Example
CH4 (g) + 2 O2 (g)  CO2 (g) + 2H2O (l)
C
H
O
Reactants
-4
+1
0
Products
+4
+1
-2
Which species is oxidized ?
(lost electrons/Ox state became more
positive)
Which species is reduced ?
(gained electrons/Ox state became more
negative)
2Na(s)+ 2H2O(l) NaOH(aq) + H2(g)
Chemical species which is oxidizing agent:
H2O
Chemical species which is reducing agent:
Na
Atom(s) being oxidized:
Na
Atom(s) being reduced:
H
Identify oxidation state of each non-oxygen atom
C2O42–(aq) + MnO4–(aq)  CO2(g) + Mn2+(aq)
+3
+7
+4
+2
Chemical species which is the oxidizing agent:
MnO4
–
Chemical species which is the reducing agent:
C2O42–
Atom(s) being oxidized:
C
Atom(s) being reduced:
Mn
Choosing a Reagent
Would oxidizing agent or reducing agent be
needed for following conversions?
+5
+4
ClO3- (aq)  ClO2 (g)
reducing agent
+6
-2
SO42- (aq)  S2- (g)
reducing agent
+2
+4
Mn2+ (aq)  MnO2 (g)
oxidizing agent
Practice
Problems 5 – 8, page 687
Problems 41 – 50, page 700