Download Formal Charge

Oxidation Numbers and Balancing Redox Equations
Formal Charge. To determine whether an atom
within a covalently bonded structure bears a
formal charge, use this algorithm:
formal charge (q) = nv – np
(1)
where nV is the number of electrons present in the
valence shell of an isolated atom and np is the
number of electrons "possessed" by the atom
within the structure. Examples: A carbon atom
with four covalent bonds in any combination
(four single, one double and two single, etc.)
would possess 4 electrons (half of each covalent
bond) and would have nv = 4, thus q = 0. An
oxygen atom with three lone pairs of a electrons
and one covalent bond would have n = 7 (three
wholly owned pairs and half of one covalent
bond) and nv = 6; thus q = –1.
Oxidation number. To determine the oxidation
number of an atom within a covalently bonded
structure, use this algorithm:
oxidation number (N) = nen – nep + q (2)
where nen is the number of bonds to atoms more
electronegative than the atom in question, nep is
the number of bonds to atoms more
electropositive than the atom in question, and q is
the formal charge, if any, borne by the atom.
Bonds to atoms with equal electronegativity (i.e.,
bonds to another atom of the same element) need
not be considered. Only F is more electronegative
than O so NO is always –2 unless an F-O bond is
present. In commonly encountered organic
compounds, only B, Si, and P are more
electropositive than H, so NH is always +1 unless
B-H, Si-H, or P-H bonds are present. The sum of
all oxidation numbers within the structure will be
equal to the overall charge, if any. Examples: The
carbonyl carbon in acetone has NC = +2, each
methyl carbon has NC = –3, each H has NH = +1,
and the O has NO = –2. the carbon atom in
formaldehyde has NC = 0. The sulfur atom in
sulfate has NS = +6 and all of the oxygens have
NO = –2, two of them bearing formal negative
charges. In nitromethane (H3C-NO2), the carbon
atom has NC = –2, each O has NO = –2, each H
has NH = +l, and the nitrogen has NN = +3.
Balancing redox equations. There are several
systems. If you've forgotten or none are familiar,
try this one:
1. Determine which substances are oxidized and
which are reduced (i.e., assign oxidation numbers
to all atoms).
2. Select stoichiometric coefficients for the
oxidized and reduced species such that (i) the
number of electrons transferred to oxidants is
equal to the number of electrons transferred from
reductants and (ii) the coefficients are minimized.
3. Check to see whether the number of oxygen
atoms is equal on each side of the equation. If not,
add water either as a product or reactant as
required.
4. Check to see whether the number of hydrogen
atoms is equal on each side of the equation. If not,
add protons to either side as required.
5. The equation should now be balanced. Check
to see that the total net charge on the reactant side
is equal to that on the product side. Any
imbalance will indicate an error in one of the
previous steps. Check also to be sure that the
stoichiometric coefficients do not have a common
divisor.
Pauling electronegativities. Electronegativity
increases upward toward the right on the periodic
table, and decreases downward to the left.
Elements commonly found in organic compounds
are shown below. Remember that these values are
inherently approximate. Each element is assigned
a single-valued electronegativity even though the
electronegativity of an atom must depend on its
valence. Sulfur in SF6, for example, must in
reality have a different electronegativity than that
in H2S.