Chemistry Reference
In-Depth Information
Lesson 6-4: Oxidation and Reduction Reactions
Historically, the term oxidation was used to describe the process of a
substance combining with oxygen. For example, the carbon in the reaction
that follows would be said to “oxidize” as it combined with oxygen to form
carbon dioxide.
C
(s)
+ O
2(g)
CO
2(g)
carbon + oxygen yields carbon dioxide
Today, this reaction would still be considered an example of oxidation,
but the actual definition of oxidation has been expanded to include more
reactions, including many in which oxygen plays no part. Oxidation is when
a substance loses electrons or appears to lose electrons, as its oxidation
number appears to increase algebraically. If we look at our reaction again,
we can see that the oxidation numbers of both oxygen and carbon change.
How do we assign oxidation numbers to the substances in the reaction?
You should start by reviewing Lesson 5-1, where we first went over the
rules for assigning oxidation numbers. For convenience, I have provided
the rules here.
Rules for Assigning Oxidation Numbers
1.
In a free element (e.g., Ba or I
2
) the oxidation number for each
atom is 0.
2.
With the exception of hydrogen, all of the elements in group 1
show an oxidation number of +1 in compounds.
3.
Hydrogen has an oxidation number of +1 in all compounds
except for hydrides (e.g., NaH), where it shows an oxidation
number of -1.
4.
The elements found in group 2, the alkaline earth metals, show
an oxidation number of +2 in all compounds.
5.
When any of the elements in column 17, the halogens, show a
negative oxidation state, it will be -1.
6.
Oxygen shows an oxidation number of -2 in all compounds,
except for the peroxides, where it is -1.
7.
In all neutral compounds, the algebraic sum of the oxidation
numbers is 0.
8.
In polyatomic ions, the algebraic sum of the oxidation numbers
is equal to the charge on the polyatomic ion.
