Transcript Unit 6 – Redox Reactions and Electrochemistry

Unit 6 – Redox
Reactions and
Electrochemistry
Redox reactions involve a transfer of electron density
from one atom to another. Two common examples
include the formation of salt and the formation of
water:

2 Na + Cl2 2 NaCl & 2H2 + O2  2H2O
The transfer of electron density in the formation of
salt is more obvious since we form two ions,
however the transfer of electron density in making
water is more difficult to see. Redox reactions
always contain 2 parts that make this process
more obvious and shows the transfer of electron
density.
Redox - a reaction that is half
reduction and half oxidation
Reduction – a gain of electron density
(reactant to product)
Oxidation – a loss of electron density
(reactant to product)
A way to remember : LEO the lion says GER
Loss of Electron density is
GER
Oxidation, Gain of Electron
density is Reduction

If we break the first reaction into its
half reactions, oxidation and reduction,
we see the following:
Na  Na+ + e- (oxidation) and
Cl2 + 2e-  2Cl- (reduction)

Thus sodium is being oxidized and chlorine
is being reduced.
No substance is ever oxidized unless
another is reduced. The substance that
accepts electrons is called the
oxidizing agent (it is reduced).

The substance that gains electrons is
called the reducing agent (it is oxidized).
For the reaction making salt, the reducing agent
is sodium (recall that it is oxidized)
and the oxidizing agent is chlorine (recall that
it is reduced).
 Examples
of redox reactions in
everyday life:
any battery - this creates a flow of electrons
between the reducing agent (the one that is
being oxidized or losing electrons) and the
oxidizing agent (the that is being reduced or
gaining electrons).
Metabolism - these use oxygen to convert
carbohydrates and fats into CO2 and H2O
Bleach – oxidizes stains so that they are
easier to remove by detergents
Oxidation Numbers

In redox reactions there is always a transfer
of electron density. This change can be
monitored by observing the oxidation
numbers of the atoms and ions involved.
Oxidation numbers is the charge that an
atom would have if the electrons in each
bond belonged entirely to the more
electronegative element.
There are several rules that make the
assignment of oxidation numbers
easier for more complex situations that
just HCl. They are always used in
order and are as follows:
1. Any element that is not combined with
others to form a molecule is given the
value 0. (ex. Na , H2, Br2)
2. Any simple monatomic ion has an oxidation
number equal to its charge.
(ex. Cl- has an oxidation number of –1, Mg2+ has an
oxidation number of +2)
3. The sum of oxidation numbers for the all the atoms
in a formula must equal the overall charge of the
formula.

(ex. CaCl2 – Ca = +2, Cl = -1, overall +2 –1-1 = 0 )
4. In compounds the oxidation number of any group I metal is ALWAYS +1
5. In compounds the oxidation number of any group II metal is ALWAYS +2
6. In compounds the oxidation number of Aluminum is ALWAYS +3.
7. In binary compounds with metals, the oxidation number of a nonmetal is
equal to the charge it would have as a monatomic ion (ex. Na3N, Na = +1,
N = -3)
8.
In compounds, F is ALWAYS –1
9. In compounds, O is almost always –2
10. In compounds, H is almost always +1
(alternative: -1 hydride)
11. In polyatomic ions, once the oxidation number is
known it is to be considered the constant so long as
the polyatomic ion is unchanged.
(ex NH4+ has N –3, H +1)
‘FUN’ = Q 1-3 p. 653, p. 657 - 659, Q 12-16 p. 659