Section 7-2
Oxidation Numbers (aka Oxidation States)
• Are used to indicate and assign the general
distribution of electrons among the bonded
atoms in a molecular compound or ion.
• In an ionic compound, the charges of each
ion tell us about the electron distribution
• In a covalent compound, the atoms are
sharing electrons. But we can imagine that
the atoms inside the molecules are charged,
and we call these charges oxidation numbers.
Oxidation Numbers (aka Oxidation States)
• Essentially, oxidation numbers help us to
keep track of electrons in reactions.
• They are a book-keeping mechanism to help
us track electron distribution – it is helpful to
know what is happening with electrons
during reactions
They Are Useful
• In naming
compounds, in
writing formulas, in
balancing chemical
equations, and in
studying certain
types of chemical
reactions (Ch. 19).
Assigning Oxidation Numbers:
General Rule
• Shared electrons are assumed to belong to the
more electronegative atom in each bond.
Specific Rules For
Assigning Oxidation Numbers
I.
Ex:
Na
O2
S8
Atoms of a pure
element have an
oxidation number
of zero.
II. Electronegativity
• The more electronegative
element in a binary
compound is assigned the
number equal to the
negative charge it would
have as an anion, the less
electronegative one is
positive (as if it were a
cation).
III. Fluorine
• Assigned a value of
-1 in all compounds
because it is the most
electronegative
element.
IV. Oxygen
• Assigned a number of -2
in almost all compounds.
Exceptions:
In peroxides like H2O2
(when it is -1)
In compounds with the
halogens, like OF2 (when
it is +2).
V. Hydrogen
• +1 in all compounds
with elements that
are more
electronegative than
it is.
• It is -1 when it is
combined with
metals.
Algebraic Sums
VI. In a neutral
compound all
oxidation numbers
add up to zero.
VII. In a polyatomic
ion the sum is
equal to the charge
of the ion.
VIII. Not For Covalently Bonded Atoms
Only
• Monotomic ions
have an oxidation
number equal to the
charge of the ion:
Na+ = +1
Cl − = -1
Because of Rules I-VIII
• It is often possible to
assign oxidation
numbers when they
are not known.
Examples
• Assign oxidation numbers to each
atom in the following compounds or
ions:
• UF6 U=+6; F=-1
• H2SO4 H=+1; S=+6; O=-2
• ClO3- Cl=+5; O=-2
More Examples
• Assign oxidation numbers to each
atom in the following compounds or
ions:
• PCl3 P=+3; Cl=-1
• PCl5 P=+5; Cl=-1
• KH K=+1; H=-1
• P4O10 P=+5; O=-2
Using Oxidation Numbers For
Formulas and Names
• Tables on p. 205, p. 219, and in the appendix (p.
903), show that both metals and nonmetals can
have more than one oxidation number.
Fe = +2 or +3
Carbon
-4, +2, +4
Nitrogen -3, +3, +5
Phosphorus
-3, +3, +5
Sulfur
-2, +4, +6
Chlorine, Bromine, Iodine -1, +1, +3, +5, +7
The Stock System of Nomenclature
Iron (II) oxide
Iron (III) oxide
• In 7.1 we showed how
Roman numerals show
the charges of cations.
But in reality the
system was based on
oxidation numbers:
FeO = Fe2+ + O2−
Fe2O3 = 2Fe3+ + 3O2−
Used As Alternative To Prefix System
• The international
body that governs
nomenclature
endorses use of the
Stock system as
being more practical
for complicated
compounds.
Name the following compounds using
the Stock system
• PCl3
• PCl5
• N2O
phosphorus (III) chloride
phosphorus (V) chloride
nitrogen (I) oxide