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REDOX
electrochemistry
Redox reactions involve the transfer of the
electron.
• Spontaneous redox reactions can
transfer energy
– Electrons (electricity)
– Heat
• Non-spontaneous redox reactions can
be made to happen with electricity.
Trends in Oxidation
and Reduction
Active metals:
Lose electrons easily
Are easily oxidized
Are strong reducing agents
Active nonmetals:
Gain electrons easily
Are easily reduced
Are strong oxidizing agents
Oxidation
- is defined as the loss of an electron
•
(LEO - loss of electron; oxidation)
- when oxidation happens, charge becomes
more positive
• Fe → Fe2+ + 2e4 Fe (s) + 3 O2 (aq) →
2 Fe2O3 (s)
Reduction
• is defined as the gain of an electron
• (GER- gain of electrons – reduction
• - when reduction happens, charge becomes
more negative
•
Fe2+ + 2e- → Fe
• - term is derived from the observation that metal
oxides lose mass when refined to produce the
pure metal:
•
2 Fe2 O3 (s) → 4 Fe (s) + 3 O2 (aq)
•
•
N2O → NO
Has the nitrogen gained or lost
electrons?
• N2O →
NO
+1
+2
Nitrogen has lost one electron and is
therefore being oxidized.
0
1
0
1
2 Na  Cl 2  2 Na Cl
1
0
Na  Na  e

Each sodium atom loses one electron:
Lose Electrons = Oxidation
0

1
Cl  e  Cl
Each chlorine atom gains one electron:
Gain Electrons = Reduction
• These pairs of reactions are called halfreactions. For example, in the reaction
•
Cu2+ (aq) + Fe (s) → Cu (s) + Fe2+ (aq)
•
two half-reactions occur; iron is oxidized
while the copper (II) ion is reduced. Seen
another way, the copper (II) ion causes the
oxidation of the iron; it is the oxidizing agent.
Likewise, the iron is the reducing agent for the
copper (II) ion.
There are five basic rules for the
determination of oxidation number:
• Rule 1: The oxidation number for any atom in
its elementary state is 0.
• Rule 2: The oxidation number for any simple
ion is the change on the ion.
a. The oxidation number of alkali metals in
compounds is 1+
(Li1+ , Na1+ , K1+ , Rb1+ , Cs1+ , Fr1+ ).
b. The oxidation number of alkaline-earth
metals in compounds is 2+
(Mg2+ , Ca2+ , and Ba2+ ).
Find the oxidation number for
magnesium and chlorine
Mg  Cl 2  Mg Cl2
0
0
2
1
Mg  Cl 2  Mg Cl
• Rule 3: The oxidation number for oxygen
usually is 2-. In peroxides, it is 1-.
1
2
H2O
water
1
1
H 2 O2
peroxide
• Rule 4: The oxidation number for
hydrogen is 1+ in all its compounds except
in metallic hydrides like NaH or BaH2 ,
where it is 1-.
• Rule 5: All other oxidation numbers are
assigned so that the sum of oxidation
numbers equals the net charge on the
molecule or polyatomic ion.
Find the oxidation numbers
• H2S
Ca(OH)2
1
2
H2 S
2(+1) + (-2) = 0
H
O
2
2 1
Ca(O H ) 2
(+2) + 2(-2) + 2(+1) = 0
Ca
O
H
Find the oxidation numbers
? 2
N O3

? 2
S O4
2
X + 3(-2) = -1
N
O
X + 4(-2) = -2
S
O
 X = +5
 X = +6
Find the oxidation number of Cr:
• Ex: Cr2O722X + 7(-2) = -2
Cr
O
 X = +6
Write the oxidation numbers for
each atom:
•
•
•
•
•
•
S8
SO2
S2O32SO42MgSO4
H2SO4
0
S = +4 and O = -2
S = +2 and O = -2
S = +6 and O = -2
Mg = +2, S = +6 and O = -2
H = +1, S = +6 and O = -2
Find the oxidation numbers
•
•
•
•
•
•
P4
P2O5
PCl5
H3PO4
PO43Na3PO4
0
P = +5 and O = -2
P = +5 and Cl = -1
H = +1, P = +5 and O = -2
P = +5 and O = -2
Na = +1, P = +5 and O = -2
A summary of terminology for oxidationreduction (redox) reactions
e-
X
transfer
Y
or shift of
electrons
X loses electron(s)
Y gains electron(s)
X is oxidized
Y is reduced
X is the reducing agent
Y is the oxidizing agent
X increases its
oxidation number
Y decreases its
oxidation number
Not all reactions are redox
1 5 2
1
1
1
1
1 5 2
Ag N O3 (aq)  Na Cl (aq)  Ag Cl ( s)  Na N O3 (aq)
1 2 1
1
6 2
1
6 2
1
2
2 Na O H (aq)  H 2 S O 4 (aq)   Na 2 S O 4 (aq)  H 2 O(l )
Reactions in which there has been no change in
oxidation number are not redox reactions.
• Just like the Bronsted-Lowry theory of
acids, oxidation and reduction also
happens in pairs; a species cannot donate
an electron unless another species gains
the electron. Since oxidation and
reduction always happen together, these
are most often called redox reactions.
Reducing and oxidizing agents
The substance reduced is the oxidizing agent
The substance oxidized is the reducing agent
1
0
Na  Na  e

Sodium is oxidized – it is the reducing agent
0

1
Cl  e  Cl
Chlorine is reduced – it is the oxidizing agent
FIND THE OXIDATION
NUMBERS:
•
2H2(g) + O2(g) → 2H2O(g)
0
0
+1
-2
• 2H2(g) + O2(g) → 2H2O(g)
The O.N. of H increases; it is oxidized; it is the reducing agent.
The O.N. of O decreases; it is reduced; it is the oxidizing agent.
Find the oxidation numbers:
2Al(s)+ 3H2SO4(aq)
0
+1
+6 -2
2Al(s) + 3H2SO4(aq)
Al2(SO4)3(aq) + 3H2(g)
+3 +6 -2
0
Al2(SO4)3(aq) +3H2(g)
The O.N. of Al increases; it is oxidized; it is the reducing agent.
The O.N. of H decreases; it is reduced; it is the oxidizing agent.
Find the oxidation numbers
• PbO(s) + CO(g) → Pb(s) + CO2(g)
+2 -2
+2 -2
0
+4 -2
• PbO(s) + CO(g) → Pb(s) + CO2(g)
The O.N. of C increases; it is oxidized; it is the reducing agent.
The O.N. of Pb decreases; it is reduced; it is the oxidizing agent.