Lesson Objectives
The student will:
 assign the correct oxidation number to any element in a compound or ion.

identify the substance being oxidized, the substance being reduced, the oxidizing agent, and the
reducing agent in an oxidation-reduction equation.
Vocabulary




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oxidation
oxidation number
oxidizing agent
redox reaction
reducing agent
reduction
Introduction
Many important chemical reactions involve the exchange of one or more electrons. The stoichiometric
calculations that chemists make on chemical reactions require a balanced equation. The “inspection”
method for balancing equations works well and quickly for many reactions, but when a more complex
electron exchange is involved, a new method for equation balancing is needed.
Definition of Oxidation and Reduction
After Lavoisier, a substance was said to be oxidized when it reacted with oxygen. A reaction with
oxygen was called oxidation. Today, the words “oxidized” and “oxidation” are still used for those
situations, but these words have now acquired a much broader, second meaning. The broader sense
of oxidation is defined as losing electrons. When a substance reacts with oxygen, it almost always
loses electrons to the oxygen, so we are simply extending the term oxidation to mean losing
electrons, whether it is to oxygen or to any other substance.
The other half of this process, the gaining of electrons, also needs a name. When an atom or an ion
gains electrons, the positive charge on the particle is decreased. For example, if a neutral sulfur atom
(charge of 0) gains two electrons, its charge becomes
, and if an
ion gains an electron, its
charge changes from
to
. In both cases, the charge on the particle is reduced by the gain of
electrons. Reduction means the gain of electrons. In chemical systems, oxidation and reduction must
occur simultaneously, so the number of electrons during oxidation must be the same as the number
of electrons gained during reduction. In oxidation-reduction reactions, electrons are transferred from
one substance to another. Here’s an example of an oxidation–reduction reaction:
In this reaction, the silver ions are gaining electrons to become silver atoms. Therefore, the silver ions
are being reduced. The copper atoms are losing electrons to become copper ions and are being
oxidized. Whenever a chemical reaction involves electrons being transferred from one substance to
another, the reaction is an oxidation–reduction reaction. An oxidation-reduction reaction can also be
referred to as a redox reaction for short.
Oxidizing and Reducing Agents
When a substance is oxidized, it loses electrons. In chemical reactions, this requires that another
substance take on those electrons and be reduced. Therefore, when a substance undergoes
oxidation, it causes another substance to be reduced. The substance that caused another substance
to be reduced is called a reducing agent. As you can see, the substance undergoing oxidation and
the reducing agent are the same substance.
Similarly, when a substance gains electrons, it is reduced. By gaining electrons, it is causing some
other substance to give up those electrons. Therefore, by undergoing reduction, the substance is
causing another substance to be oxidized and is called an oxidizing agent. Again, the substance
undergoing reduction and the oxidizing agent are the same substance.
In the oxidation-reduction reaction above, silver ions are being reduced and are the oxidizing agent.
Similarly, copper atoms are being oxidized and are the reducing agent. These substances are always
on the reactant side of the equation.
Oxidation Numbers
In order to balance oxidation and reduction reaction equations, it is necessary to have a bookkeeping
system to keep track of the transferred electrons. The bookkeeping system chemists use to keep
track of electrons in oxidation-reduction reactions is called oxidation numbers. The assignment of
oxidation numbers to all the atoms or ions in a reaction follows a set of rules. For the most part, these
rules will have the oxidation number of a particle be the number of electrons the atom has gained or
lost from its elemental state. For example, for a
ion, the calcium ion has clearly lost two
electrons from its elemental form, so its oxidation number is
. Similarly, it is clear that a fluoride
ion,
, has gained one electron from its elemental state and has an oxidation number of
.
The first rule for assigning oxidation numbers is for substances in their elemental form. Substances in
elemental form have oxidation numbers of zero. It is clear that substances in elemental form have
neither gained nor lost any electrons from their elemental state (Table below).
Examples of Oxidation Numbers for Substances in Elemental Form
Substance
Oxidation Number
The second rule for assigning oxidation numbers is for monatomic ions. For monatomic ions, the
oxidation number is the same as the charge on the ion. Again, it should be apparent that the charge
on the ion is an indication of how many electrons have been gained or lost. Table below shows
examples of oxidation numbers for monatomic ions.
Examples of Oxidation Numbers for Monatomic Ions
Substance
Oxidation Number
The third rule is for the atoms of family IA, the alkali metals, in compounds. Alkali metals always lose
their single valence electron when they combine. Therefore, for IA metals, the oxidation number in
compounds is
(see Table below for examples).
Examples of Oxidation Numbers for Alkali Metals in Compounds
Substance
Oxidation Number for the Alkali Metal
The fourth rule is for the atoms of family IIA, the alkali earth metals, in compounds. Alkali earth metals
always lose both of their valence electrons when they combine chemically, so for IIA metals, the
oxidation number in compounds is
(see Table below for examples).
Examples of Oxidation Numbers for Alkali Earth Metals in Compounds
Substance
Oxidation Number for the Alkali Earth Metal
The fifth rule concerns hydrogen atoms when they are in compounds. In the great majority of
compounds that hydrogen forms, it either completely or at least partially loses that electron. In
compounds where hydrogen is the more electropositive atom, the oxidation number for hydrogen
is
(see Table below for examples).
Examples of Oxidation Numbers for Hydrogen in Compounds
Substance
Oxidation Number for Hydrogen
There is, however, an exception to this rule for hydrogen. It is possible for hydrogen to form
compounds with some metals that are even more electropositive than hydrogen. In these cases,
hydrogen becomes an electron acceptor instead of an electron donor. Active metals lose or partially
lose their valence electrons to hydrogen. Since hydrogen is acting as the more electronegative
element in these compounds, the compounds are named hydrides (see examples in Table below). In
hydrides, the oxidation number of hydrogen is
.
Examples of Oxidation Numbers for Hydrogen in Hydride Compounds
Substance
(lithium hydride)
(sodium hydride)
(magnesium hydride)
Oxidation Number for Hydrogen
The sixth rule is about the oxidation number of oxygen in compounds. Oxygen is a very
electronegative element that draws two electrons completely or partially from a bonding element in
almost all of its compounds. Therefore, the oxidation number for oxygen in compounds is almost
always
(see examples in Table below).
Examples of Oxidation Numbers for Oxygen in Compounds
Substance
Oxidation Number for Oxygen
Like hydrogen, there is an exception to the rule for oxygen. In a group of compounds named
“peroxides” (hydrogen peroxide
, sodium peroxide
, etc.), each of the oxygen atoms
shares a bond with the other oxygen atom. Therefore, the oxygen atoms only accept one electron
from the other element, so oxygen has an oxidation number of
in peroxides (see examples
in Table below).
Examples of Oxidation Numbers for Oxygen in Peroxides
Substance
Oxidation Number for Oxygen
Some other elements exhibit only one possible gain, loss, or sharing of electrons when they form
compounds. Aluminum, for example, always loses or partially loses three electrons when it forms
compounds, so its oxidation number in compounds is
. Zinc always loses or partially loses two
electrons when it combines, so its oxidation number in compounds is
. The halogens (family VIIA)
always gain one electron in binary compounds and would have an oxidation number of
in these
compounds. Some of the halogen atoms form compounds where there are three elements, and one
of them is oxygen (
, for example). In these compounds, the halogen atom is almost never
.
There will be quite a few atoms whose oxidation number may be different in different compounds. For
these, you will have to calculate their oxidation numbers. To allow you make such calculations, there
is a general rule that the sum of the oxidation numbers of all the atoms in a compound must be zero,
and the sum of all the oxidation numbers of the atoms in a polyatomic ion must equal the charge on
the ion.
Example:
What is the oxidation number of sulfur in
Solution:
?
We have three pieces of information that will allow us to calculate the oxidation number of sulfur. We
know that the sum of the oxidation numbers of all the atoms will equal zero.
We also know the oxidation number of sodium from rule is
and the oxidation number of oxygen
from rule is
. We simply plug these values into the equation and solve for the oxidation number
of sulfur.
so
Example:
What is the oxidation number of chromium in
Solution:
?
so
Example:
What is the oxidation number of nitrogen in the nitrate ion,
Solution:
?
so
Example:
What is the oxidation number of phosphorus in the phosphate ion,
Solution:
so
?
Example:
What is the oxidation number of iron in
Solution:
?
so
Example:
What is the oxidation number of iron in
Solution:
?
This example was chosen specifically to make a point.
so
In this case, we get an oxidation number that is not a whole number. Since this number supposedly
represents the number of electrons gained or lost from the elemental state, we should feel
uncomfortable about an oxidation number of . We know that an atom did not lose a fraction of an
electron. The reason that the oxidation number appears to be fractional is because some of the iron
atoms in
lost 3 electrons, and some lost 2. Another way to represent this molecule is as
FeO•Fe2O3. In this case, it is easy to see that the Fe attached to one O is in a +2 oxidation state,
whereas the
in the
is in the +3 oxidation state.
Review Questions
1. Indicate the oxidation numbers for each of the following atoms.
a.
b.
c.
in
d.
in
e.
in
f. arsenic in
g. chlorine in
h. sulfur in
2. In the following reaction, identify the element that is being oxidized and the element that is being
reduced:
.