Chapter 15.
Oxidation-Reduction
Reactions
Oxidation-Reduction
Reactions
Oxidation was originally understood as reaction
of a substance with oxygen.
Combustion:
CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(g)
+ Energy!
Metabolism:
C6H12O6(s) + 6 O2(g)  6 CO2(g) + 6 H2O(g)
+ Energy!
Oxidation-Reduction
Reactions
Corrosion:
2 Mg(s) + O2(g)  2 MgO(g)
4 Fe(s) + 3 O2(g)  2 Fe2O3(s)
Rust!
Oxidation-Reduction
Reactions
Reduction was originally understood as the
loss of mass of metal ores in smelting.
2 Fe2O3(s) + 3 C(s)  4 Fe(s) + 3 CO2(g)
Oxidation is a process in which a substance
loses electrons.
Reduction is a process in which a substance
gains electrons.
4 Fe(s) + 3 O2(g)  2 Fe2O3(s)
4 (Fe0  Fe3+ + 3e1) Iron is oxidized
4 Fe0  4 Fe3+ + 12 e1 forms cation
3 (O2 + 4 e1  2 O2) Oxygen is reduced
3 O2 + 12 e1  6 O2
forms anions
Oxidation-Reduction
Reactions
A half-reaction is a chemical equation that
shows either the oxidation or reduction
part of an oxidation-reduction reaction.
Electrons appear as products in oxidations.
Zn(s)  Zn2+(aq) + 2 e1
Electrons appear as reactants in reductions.
I2(aq) + 2 e1  2 I1 (aq)
L E O the lion says G E R!
Lose Electrons Oxidation
Gain Electrons Reduction
Oxidation-Reduction
Reactions
Examples:
Which of these half-reactions show oxidations?
Which ones show reductions?
Zn(s)  Zn2+(aq) + 2 e1
Cu2+(aq) + 2 e1  Cu(s)
Ag(s)  Ag1+(aq) + e1
Cr2O72 + 14H1+ + 6 e1  2 Cr3+ + 7 H2O
2 F1(aq)  F2(g) + 2 e1
Oxidation-Reduction
Reactions
In ionic compounds, we can look at changes
in charge to figure out what's reduced and
what's oxidized.
What happens with molecular compounds?
CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(g)
Use Oxidation Numbers
Determining Oxidation Numbers
An oxidation number (or oxidation state)
is the charge that an atom appears to
have when the electrons in each bond in
which it is participating are assigned to the
more electronegative of the two atoms
participating in the bond.
Determining Oxidation Numbers
Draw Lewis Structure of molecule or ion.
Use Electronegativities (p 281, Fig 7.12) to
determine which atom "owns" electrons.
Lone pairs belong to the atom.
If two atoms of the same element are
bonded together, 1 electron to each.
ONAtom = VEAtom - TEAtom
Determining Oxidation Numbers
The oxidation number of a main-group element
can be anything from Column # (loses all e1)
to (Column #  8) (fills valence shell).
The sum of all oxidation numbers in a species
must equal its charge.
The oxidation number of an atom in its elemental
form is zero.
If a Lewis structure is complex, just work with the
atom of interest and the atoms bonded to it.
Determining Oxidation Numbers
Examples:
Determine the oxidation number of:
N in N2
H in H2
S in S8
C and H in CH4 N and O in N2O
S in H2SO4
Cl in ClO41
C in Urea, CH4N2O
Determining Oxidation Numbers
Polyatomic ions in which the central atom is a
transition metal:
Oxygen's ON is 2. Hydrogen's ON is +1.
The sum of all oxidation numbers in a
species must equal its charge.
What are the oxidation states of the metals?
MnO41
CrO42
Cr2O72
Vocabulary
A substance is oxidized if:
it loses electrons
it loses hydrogen atoms
it gains oxygen atoms
its charge or oxidation number
increases (becomes more positive).
Vocabulary
A substance is reduced if:
it gains electrons
it gains hydrogen atoms
it loses oxygen atoms
its charge or oxidation number
decreases (becomes less positive).
Vocabulary
An agent causes something to happen.
An oxidizing agent causes oxidation.
It does this by being reduced (gaining
electrons).
A reducing agent causes reduction.
It does this by being oxidized (losing
electrons).
Vocabulary
Typical Oxidizing Agents:
O2
ClO41
ClO31
NO31
O3
Oxygen
MnO41 Permanganate
Perchlorate CrO42 Chromate
Chlorate
Cr2O72 Dichromate
Nitrate
Cl2,
other halogens
Ozone
H 2O 2
Peroxides
Vocabulary
Typical Reducing Agents:
Hydrogen sources:
H2
Hydrogen
NaBH4 Sodium borohydride
LiAlH4 Lithium aluminum hydride
Active metals (1A and 2A metals, zinc)
Reduced Carbon (C, CO)
Examples:
Identify the substance being oxidized, the substance being reduced, the oxidizing agent,
and the reducing agent:
2 H2(g) + HCCH(g)  C2H6(g)
2 HCl(aq) + Zn(s)  ZnCl2(aq) + H2(g)
4 C3H6O(l) + NaBH4(s) + 4 H2O(l) 
4 C3H8O(l) + NaB(OH)4(aq)
Examples:
Identify the substance being oxidized, the substance being reduced, the oxidizing agent,
and the reducing agent:
2 Na(s) + Cl2(g)  2 NaCl(s)
C2H4O(l) + H2(g)  C2H6O(l)
2 KMnO4 + 5 H2O2 + 3 H2SO4 
2 MnSO4 + K2SO4 + 5 O2(g) + 8 H2O
Predicting Reactions
Who does what to whom?
Cu2+(aq) + 2 Ag(s)  Cu(s) + 2 Ag1+(aq)
Cu(s) + 2 Ag1+(aq)  Cu2+(aq) + 2 Ag(s)
Cu2+(aq) + Zn(s)  Cu(s) + Zn2+(aq)
Cu(s) + Zn2+(aq)  Cu2+(aq) + Zn(s)
2 Ag1+(aq) + Zn(s)  Zn2+(aq) + 2 Ag(s)
2 Ag(s) + Zn2+(aq)  2 Ag1+(aq) + Zn(s)
A Daniell Cell
Daniell Cells
Voltages from Daniell Cells provide information
about energy changes:
Cu(s) + 2 Ag1+(aq)  Cu2+(aq) + 2 Ag(s)
0.34 V
Cu2+(aq) + Zn(s)  Cu(s) + Zn2+(aq)
1.10 V
2 Ag1+(aq) + Zn(s)  Zn2+(aq) + 2 Ag(s)
1.56 V
An Energy Scale for Redox RXN’s
Standard Reduction Potentials are potentials
measured against a standard hydrogen electrode, with all solutions at 1.0 M and gas
pressures at 1.0 atm.
A Standard Hydrogen Electrode is an electrode in which the half-reaction is:
2 H1+ + 2 e1  H2(g)
Materials are at standard conditions, and the
half-cell potential is 0.00 V.
An Energy Scale for Redox RXN’s
From chemguide.co.uk
An Energy Scale for Redox RXN’s
Some Standard Reduction Potentials:
Reaction
E0, V
Li1+(aq) + e1  Li(s)
Al3+(aq) + 3 e1  Al(s)
Zn2+(aq) + 2 e1  Zn(s)
2 H 1+(aq) + 2 e1  H2(g)
Cu2+(aq) + 2 e1  Cu(s)
Ag1+(aq) + e1  Ag(s)
Cr2O7 + 14H1+ + 6 e1  2 Cr3+ + 7 H2O
F2(g) + 2 e1  2 F1(aq)
3.04
1.66
0.76
0.00
0.34
0.80
1.33
2.87
An Energy Scale for Redox RXN’s
Strongest oxidizing agents have highest
standard reduction potentials (E0's, in volts).
Strong oxidizing agents have low-energy
vacant orbitals.
Strongest reducing agents have lowest (often
negative) standard reduction potentials.
Strong reducing agents have electrons in
high-energy orbitals.
Predicting Reactions
Balancing equations for Redox Reactions
Use Table of Standard Reduction Potentials:
Find half-reaction for reduction,
write it as given.
Find half-reaction for oxidation,
write it as reverse of what is given
(electrons are products).
Predicting Reactions
Balance half-reactions so number of electrons
is equal for each.
Add half-reactions, cancelling out electrons
and any other species that appear on both
sides of equation.
Predicting Reactions
Will the Reaction be Spontaneous?
Calculate E0 for RXN:
E0 reduction
- E0 oxidation
E0 RXN
If E0 RXN is positive, the reaction is exothermic and will proceed spontaneously.
If E0 RXN is negative, the reaction is endothermic and will not proceed spontaneously.
Predicting Reactions
Balance the reactions.
Which of them are spontaneous ?
Cu(s) + Zn2+(aq)  Cu2+(aq) + Zn(s)
Zn(s) + Ag1+(aq)  Zn2+(aq) + Ag(s)
Cu(s) + H1+(aq)  Cu2+(aq) + H2(g)
Fe(s) + Al3+(aq)  Fe2+(aq) + Al(s)
Cu(s) + Cr2O72–(aq)  Cu2+(aq) + Cr3+(aq)
Batteries
Batteries are galvanic (a.k.a. voltaic) electrochemical cells in which a spontaneous reaction is used to convert chemical energy to
electrical energy.
The earliest batteries common galvanic cells
were Daniell Cells, from 1836.
Batteries
Batteries usually contain an electrolyte, which
is a solution that contains ions. Electric current can flow through the electrolyte.
The anode in a battery is the half-cell with the
lower (more negative) E0. Oxidation occurs
at the anode. Anions flow toward the anode.
The cathode in a battery is the half-cell with
higher (more positive) E0. Reduction occurs
at the cathode. Cations flow toward the
cathode.
A Daniell Cell
Batteries
Dry cell (ca. 1900, developed from Leclanché
cell, 1866. Alkaline cell, ca. 1950)
Batteries
Dry cell, acid form; Cell potential ~1.5 V
Anode half-reaction:
Zn(s)  Zn2+(aq) + 2 e1–
Cathode half-reaction:
2 MnO2(s) + 2 NH4Cl(aq) + 2 e1– 
Mn2O3(s) + 2 NH3(aq) + H2O(l) + 2 Cl1–(aq)
Batteries
Dry cell, alkaline form; Cell potential ~1.5 V
Anode half-reaction:
Zn(s) + 2 OH1–(aq)  ZnO(s) + H2O(l) + 2 e1–
Cathode half-reaction:
2 MnO2(s) + 2 H2O(l) + 2 e1– 
Mn(OH)2(s) + 2 OH1–(aq)
Batteries
“Button” battery, ca. 1940, HgO or Ag2O
Batteries
Button battery; Cell potential ~1.5 V
Anode half-reaction:
Zn(s) + 2 OH1–(aq)  ZnO(s) + H2O(l) + 2 e1–
Cathode half-reaction:
Ag2O(s) + 2 H2O(l) + 2 e1–  2 Ag(s) + 2 OH1–(aq)
HgO(s) + H2O(l) + 2 e1–  Hg22+(aq) + 2 OH1–(aq)
Batteries
Lead-Acid Battery (ca. 1859)
Batteries
Detail of Lead-Acid Battery: Six cells in series
produce about 12 volts.
Batteries
Lead-acid battery; Cell potential ~2.1 V
Anode half-reaction:
Pb(s) + SO42–(aq)  PbSO4(s) + 2 e1–
–0.36 V
Cathode half-reaction:
PbO2 (s) + 4 H1+(aq) + SO42–(aq) + 2 e1– 
PbSO4(s) + 2 H2O(l)
+1.69 V
Batteries
Nickel-Metal Hydride (Ni-MH) Battery;
cell potential ~1.4V
Anode half-reaction:
MH(s) + OH1–(aq)  M(s) + H2O(l) + e1–
Cathode half-reaction:
NiO(OH)(s) + H2O(l) + e1– 
Ni(OH)2(s) + OH1–(aq)
MH(s) is metal alloy, e.g. LaNi5, into which hydrogen
is absorbed.
Batteries
Lithium ion battery;
cell potential ~3.7V
Chemistry is complex
because water cannot be used in the
electrolyte.
Lithium embedded in
graphite is the anode.
Electrolytic Cells
An electrolytic cell is an electrochemical
cell in which a non-spontaneous chemical
change is caused to occur by application
of electrical energy.
Many important chemical processes are
carried out in electrochemical cells.
Electrolytic Cells
A chlor-alkali cell
Electrolytic Cells
A Hall-Heroult Cell for Aluminum
Electrolytic Cells
An electrorefining cell for purifying copper
Electrolytic Cells
An electroplating cell