207
Chapter
12:
Oxidation
and
Reduction.
Oxidation-‐reduction
(redox)
reactions
At
different
times,
oxidation
and
reduction
(redox)
have
had
different,
but
complimentary,
definitions.
Compare
the
following
definitions:
Oxidation
is:
Reduction
is:
•Gaining
oxygen
•Losing
oxygen
•Losing
hydrogen
•Gaining
hydrogen
•Losing
electrons
(+
charge
increases)
•Gaining
electrons
(+
charge
decreases)
Oxidation
and
reduction
are
opposite
reactions.
They
are
also
paired
reactions:
in
order
for
one
to
occur,
the
other
must
also
occur
simultaneously.
While
the
first
two
definitions
of
oxidation-‐reduction
are
correct,
the
most
useful
definition
is
the
third
-‐
involving
the
gain
or
loss
of
electrons.
I
will
use
this
particular
definition
for
the
rest
of
our
discussion.
How
do
we
know
whether
or
not
an
element
has
gained
or
lost
electrons?
To
determine
which
elements
have
been
oxidized
or
reduced,
we
look
at
changes
in
the
oxidation
number
of
the
element.
English
scientist
Michael
Faraday
(1791
–
1867),
one
of
the
greatest
pioneers
in
electrochemistry
and
electromagnetism
of
the
19th
century,
developed
the
system
of
oxidation
numbers
to
follow
these
kinds
of
reactions.
The
oxidation
number
describes
the
oxidation
state
of
the
element
in
a
compound,
and
these
numbers
are
assigned
following
a
relatively
simple
set
of
rules.
Rules
for
assigning
oxidation
numbers.
1. The
oxidation
number
of
an
element
in
its
elemental
form
is
0
(zero).
2. The
oxidation
number
of
a
simple
ion
is
equal
to
the
charge
on
the
ion.
Both
the
size
and
the
polarity
of
the
charge
are
part
of
the
oxidation
number:
an
ion
can
have
a
+2
oxidation
number
or
a
-‐2
oxidation
number.
The
“+”
and
“-‐“
signs
are
just
as
important
as
the
number!
3. The
oxidation
numbers
of
group
1A
and
2A
(group
1
and
group
2)
elements
are
+1
and
+2
respectively.
4. In
compounds,
the
oxidation
number
of
hydrogen
is
almost
always
+1.
The
most
common
exception
occurs
when
hydrogen
combines
with
metals;
in
this
case
the
oxidation
number
of
hydrogen
is
typically
–1.
,
208
5. In
compounds,
the
oxidation
number
of
oxygen
is
almost
always
–2.
The
most
common
exception
is
in
peroxides,
when
the
oxidation
number
is
–1.
Peroxides
are
compounds
having
two
oxygen
atoms
bonded
together.
For
example,
hydrogen
peroxide
is
H-‐O-‐O-‐H.
In
hydrogen
peroxide,
each
oxygen
atom
has
a
-‐1
oxidation
number.
When
oxygen
is
bonded
to
fluorine,
as
in
hypofluorous
acid
(HOF),
the
oxidation
number
of
oxygen
is
0.
Oxygen-‐
fluorine
compounds
are
relatively
rare
and
not
too
terribly
important
for
our
studies.
6. In
compounds,
the
oxidation
number
of
fluorine
is
always
–1.
The
oxidation
number
of
other
halogens
(Cl,
Br,
I)
is
also
–1,
except
when
they
are
combined
with
oxygen.
The
oxidation
number
of
halides
(except
fluorine)
combined
with
oxygen
is
typically
positive.
For
example,
in
ClO-‐,
chlorine’s
oxidation
number
is
+1.
7. For
a
complex
ion,
the
sum
of
the
positive
and
negative
oxidation
numbers
of
all
elements
in
the
ion
equals
the
charge
on
the
ion.
8. For
an
electrically
neutral
compound,
the
sum
of
the
positive
and
negative
oxidation
numbers
of
all
elements
in
the
compound
equals
zero.
Identifying
redox
reactions.
Now
that
you
can
assign
proper
oxidation
numbers
to
the
elements
in
substances,
we
can
use
changes
in
the
oxidation
numbers
to
identify
oxidation
and
reduction
reactions.
Consider
the
following
chemical
equation:
Zn(s)
+
2HCl(aq)
→
Zn+2(aq)
+
2Cl-‐(aq)
+
H2(g)
The
oxidation
number
for
elemental
zinc
is
0.
The
oxidation
number
for
zinc
ion
is
+2.
The
oxidation
number
for
hydrogen
in
hydrogen
chloride
is
+1.
The
oxidation
number
for
elemental
hydrogen,
H2,
is
0.
The
oxidation
number
for
chlorine
in
hydrogen
chloride
is
–1.
The
oxidation
number
for
chloride
ion
is
–1.
Zinc
has
lost
two
electrons,
and
therefore
developed
a
+2
charge.
Zinc
has
been
oxidized
–
its
oxidation
number
has
become
more
positive
(from
0
to
+2).
Hydrogen
has
gained
an
electron,
and
its
positive
charge
has
decreased.
Hydrogen
has
been
reduced
–
its
oxidation
number
has
become
less
positive
(from
+1
to
0).
Chloride
has
not
changed
its
oxidation
state.
It
is
a
spectator
ion
in
this
reaction.
,
209
Figure
12.2
may
be
useful
in
deciding
if
an
element
has
been
oxidized
or
reduced.
If
an
elements
oxidation
number
increases
(moves
towards
the
right),
then
the
element
is
oxidized.
If
an
elements
oxidation
number
decreases
(moves
towards
the
left),
then
the
element
is
reduced.
NOTE:
an
element
doesn’t
have
to
become
positive
or
negative
for
oxidation
or
reduction
to
occur.
Instead,
the
element
has
to
become
more
positive
or
more
negative.
A
change
from
-‐3
to
-‐1
is
still
oxidation,
while
a
change
from
+3
to
+1
is
still
reduction.
Figure
12.2.
Graphic
description
of
oxidation
and
reduction.
Substances
that
cause
changes
in
the
oxidation
state
are
called
oxidizing
agents
or
reducing
agents.
An
oxidizing
agent
causes
oxidation
number
to
occur.
How
does
the
oxidizing
agent
cause
oxidation?
To
increase
the
positive
oxidation
number
of
an
element,
the
oxidizing
agent
must
take
one
or
more
electrons
from
the
element.
As
the
element
being
oxidized
loses
electron(s),
its
oxidation
number
becomes
more
positive.
However,
the
electrons
don’t
disappear!
The
oxidizing
agent
has
taken
these
electrons,
and
therefore
the
oxidizing
agent
becomes
more
negative
–
it
is
reduced!
In
our
reaction
above,
hydrogen
in
hydrogen
chloride
takes
an
electron
from
the
zinc
metal.
The
zinc
metal
becomes
more
positive;
it
is
oxidized.
By
taking
the
electron
from
the
zinc
metal,
the
hydrogen
in
hydrogen
chloride
becomes
less
positive;
it
is
reduced.
Hydrogen
chloride
is
the
oxidizing
agent,
because
it
contains
the
element
that
causes
oxidation
to
occur.
Similarly,
the
zinc
metal
donates
electrons
to
the
hydrogen
in
hydrogen
chloride,
causing
the
oxidation
state
of
hydrogen
to
decrease
from
+1
to
0.
By
