Chapter 18—Oxidation-Reduction Reactions
I.
Price
Section 18.1—Oxidation Reduction Reactions
A. A very common type of reaction where electrons are
transferred from one atom (or group of atoms) to
another.
B. When a metal and a nonmetal react the metal gives
some electrons to the nonmetal and they form an ionic
compound
C. The metal is “oxidized”
D. The nonmetal is “reduced”
E. Oxidation=a loss of electrons
F. Reduction=a gain of electrons
II.
Section 18.2—Oxidation States
A. Electron Transfer is more difficult to see when the
redox reaction involves 2 or more non metals
B. To make this easier we give “imaginary” charges to the
atoms in covalent compounds
C. Oxygen is a part of all redox reactions involving
nonmetals
D. Oxygen = -2 (the peroxide exception!)
E. Hydrogen = +1
F. Oxidation states depend on the electronegativity of
the atoms involved
1. The more electronegative element always
controls the electrons
2. The most electronegative elements are
F>O>N>Cl
III. Section 18.3—Redox Reactions Between Nonmetals
A. Let’s examine the combustion of methane (CH4)
CH4 + 2O2  CO2 + 2H2O
CH4 + 2O2  CO2 + 2H2O
B. Pure elements have an oxidation number of 0 (zero)
C. We know oxygen is always a -2 and hydrogen is always
a +1
D. We use these facts to deduce the oxidation numbers
of the rest of the elements in the reaction
CH4 + 2O2  CO2 + 2H2O
C=-4 H=+1
oxygen=0  C=+4 oxygen=-2
H=+1 O=-2
E. Notice that the oxidation number of the carbon
increases from -4 to +4…carbon appears to have lost
8 electrons. The methane has been oxidized
F. Notice that the oxidation number of the oxygen has
decreased from 0 to -2. The oxygen has been
reduced
G. Oxidation=an increase in oxidation number
H. Reduction=a reduction in oxidation number
Rules for Assigning Oxidation Numbers
1.
2.
3.
4.
5.
The oxidation number for all uncombined elements is zero.
The oxidation number of a single ion is the same as its charge
Hydrogen has an oxidation number of +1
Oxygen has an oxidation number of -2 (In peroxides oxygen is a -1)
In binary compounds without hydrogen or oxygen the atom with the
greater ectronegativity is assigned the negative oxidation number based
on its charge
6. The overall charge of a compound MUST BE ZERO
7. For polyatomic ions, the sum of the oxidation numbers must equal the
overall charge
I. “Oxidizing Agents” and “Reducing Agents”
2Na + Cl2  2NaCl
1. It’s easy to see that the sodium is oxidized and the
chlorine is reduced in this simple metal/nonmetal
redox reaction
2. Sodium loses electrons, chlorine gains electrons
3. We say that sodium is the “reducing agent” and
that chlorine is the “oxidizing agent”
IV. Section 18.4--Balancing Redox Reactions
A. Until now we have been able to balance redox
reactions just like we would balance any other type of
reaction
B. Unfortunately, some redox reactions (those involving
ions that take place in aqueous solutions) are so
complex that we separate the reaction into two halfreactions
C. Half-reactions are equations that have electrons as
reactants or products.
EMAMPLE:
Ce4+ + Sn2+  Ce3+ + Sn4+
This reaction can be separated into a half-reaction involving the substance being reduced
e- + Ce4+  Ce3+
and a half-reaction involving the substance being oxidized
Sn2+  Sn4+ + 2e-
The key to balancing these half-reactions is that the number of
electrons lost must equal the number of electrons gained
It is easy to see that we need to double everything in the first half-reaction
2e- + 2Ce4+  2Ce3+
Now add the second half-reaction
Sn2+  Sn4+ + 2e2e- + 2Ce4+ + Sn2+  2Ce3+ + Sn4+ + 2eFinally, cancel out the 2 electrons on each side
2e- + 2Ce4+ + Sn2+  2Ce3+ + Sn4+ + 2e-
For the final, balanced reaction
2Ce4+ + Sn2+  2Ce3+ + Sn4+
D. It turns out that most redox reactions happen in
solutions that are acidic. This makes things a bit
more complicated
Rules for Balancing Reactions by the Half-Reaction Method
1. Identify and write the equations for the oxidation and
reduction half-reactions
2. For each half-reaction
a. balance all of the elements except hydrogen and
oxygen
b. Balance oxygen using H2O
c. Balance Hydrogen using H+ ions
d. Balance the charge using electrons
3. If necessary, multiply one or both half-reactions to equalize
the number of electrons transferred
4. Add the half-reactions and cancel out anything that appears
on both sides
5. Check to be sure the elements and charges balance
V.
Section 18.5—Electrochemistry: An Introduction
A. Definition: Electrochemistry is the study of the
interchange of chemical and electrical energy.
B. Electrochemistry involves 2 types of processes
1. The production of an electrical current from an
oxidation-reduction reaction
2. The use of an electrical current to produce a
chemical change
C. Question: How do we harness the electron transfer
of redox reactions in order to do work for us?
D. Answer: We have to separate the oxidizing agent
(the electron acceptor) from the reducing agent (the
electron donator) and force the electrons to travel
through a wire in order to complete the transfer.
E. Consider the following typical electrochemical cell
However, when we construct the above apparatus no flow of electrons
occurs. Why? The problem is that if the electrons flowed from the right to
the left the left compartment would build up a negative charge and the right
would build up a net positive charge—this is no good!
We must connect the two solutions so that ions can be transferred in order
to cancel out the charge. To do this we need a salt bridge or a semipermeable membrane that will allow ions to move but will not let the
solutions to mix extensively. This type of a battery is called a galvanic
cell.
F. Definitions:
1. Electrode: Another name for the terminals of
a battery (one negative, one positive)
2. anode: The electrode where oxidation occurs
3. cathode: The electrode where reduction
occurs
4. electrolysis: using electrical energy to produce
a chemical change
VI. Section 18.6—Batteries
A. 2 important types of batteries have been around for a
long time:
1. Lead storage battery (like the type in your car)
2.Dry cell batteries (like Energizers and Duracells)
Both of these types of reactions are powered by
complex redox reactions.
VII. Section 18.7 Corrosion
A. Corrosion (what we call rust) is simply a redox
reaction where the metal is slowly reacting with the
oxygen in the atmosphere.
B. Many metals develop an oxide coating that helps
prevent further corrosion of the interior atoms.
C. Many methods to retard or prevent corrosion have
been developed: plating, alloying and cathodic
protection are mentioned in this section.
VIII.Section 18.8—Electrolysis
A. Many batteries can be recharged by forcing the
electrical current to go backwards and deposit
electrons onto the anode
B. This is an example of electrolysis. When we force an
electrical current through water we get hydrogen gas
and oxygen gas
Electrical current + 2H2O  2H2 + O2
C. Purifying metal ores and electroplating one metal with
another are important examples of electrolysis