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Oxidation-Reduction Reactions: Redox
The technology of metallurgy has allowed humanity to progress from the Stone Age,
through the Bronze age, and the Iron Age, to modern times. Very few metals (such as Au
& Ag) exist as pure elements; most exist in a variety of compounds mixed with other
substances in rocks called ores.
Oxidation
Reduction
Although the technological processes of refining vary from one metal to another, the
processes of refining involve a large volume of ore that is reduced to a smaller volume of
metal. From metallurgy, the term reduction came to be associated with producing metals
from their compounds. The production of iron, tin, and copper metals are typical
examples:
Fe2O3(s) + 3 CO(g)  2 Fe(s) + 3 CO2(g)
SnO2(s) + C(s)  Sn(s) + CO2(g)
CuS(s) + H2(g)  Cu(s) + H2S(g)
Oxidizing Agent
Another substance, called a reducing agent, causes or promotes the reduction of a metal
compound to an elemental metal.
1.
2.
Before metallurgy, humans discovered fire. The technology of fire has been crucial in the
development of human cultures, but only relatively recently (18th century) have we come
to realize the role of oxygen in burning. Understanding the connection of corrosion
(rusting, tarnishing, etc.) and burning is an even more recent development. Reactions of
substances with oxygen, whether they were the explosive combustion of gunpowder, the
burning of wood, or the slow corrosion of iron came to be called oxidation. It soon
became apparent that oxygen was not the only substance that could cause reactions with
characteristics of oxidation. The rapid reaction process we call burning may even take
place with gases other than oxygen. The term oxidation has been extended to include a
wide range of combustion and corrosion reactions, such as:
2 Mg(s) + O2(g)  2 MgO(s)
2 Al(s) + 3Cl2(g)  2 AlCl3(s)
Cu(s) + Br2(g)  CuBr2(s)
A substance that causes or promotes oxidation is called an oxidizing agent.
Theoretical Definitions of Reduction and Oxidation
Redox reactions involve a transfer of electrons. In a redox reaction, one substance loses
electrons and is oxidized, while another substance gains electrons and is reduced.
Oxidation and reduction are two halves of a reaction and one cannot happen without the
other. The number of electrons lost by one substance must equal the number of
electrons gained by the other. In order to track electron transfer, oxidation numbers
are used.
-
loss of electrons, and an increase in oxidation number
gain of electrons, and a decrease in oxidation number
LEO GER
Reducing Agent
causes oxidation by removing (gaining) electrons and is
reduced in the process
causes reduction by donating (losing) electrons and is oxidized
in the process
Rules for Assigning Oxidation Numbers
3.
4.
5.
6.
The oxidation number of any free element is zero.
The oxidation number of a monatomic ion (Na+, Ca2+, Cl-, etc.) is equal to the charge
on the ion.
The oxidation number of each hydrogen atom in a compound is 1+, except in metal
hydrides (NaH, LiH, etc.) where it is 1-.
The oxidation number of each oxygen atom in a compound is 2-, except in peroxides
(H2O2, Na2O2, etc.) where it is 1-.
For any neutral compound, the sum of the oxidation numbers of the atoms in the
compound must equal zero.
For a polyatomic ion, the sum of the oxidation numbers must equal the ionic charge
of the polyatomic ion.
Balancing Redox Reactions
Using the Half-Reaction Method
Step 1:
Step 2:
Step 3:
Step 4:
Write the equation in ionic form
Write the skeleton half-reactions for the oxidation and reduction processes
If necessary balance the atoms that are being oxidized/reduced and
adjust the number of electrons
Balance elements other than oxygen and hydrogen.
Balance oxygen and hydrogen.
In aqueous solutions, either H+ and H2O or OH- and H2O are available.
Chemistry 12 - Redox – page 1
Redox
1.
In the following, give the oxidation number for the indicated atoms.
a.
S in Na2SO3
b. Mn in KMnO4
c.
N in Ca(NO3)2
d. C in Na2CO3
g.
e.
N in NO2
h. S in Al2S3
f.
S in HSO4-
i.
S in H2S2O7
Mn in MnCl2
2. Some of the following unbalanced reactions are oxidation-reduction reactions, and some are not. In each
case is the reaction redox? If yes, name the element reduced, the element oxidized, the oxidizing agent,
and the reducing agent.
a.
BaCl2 + Na2SO4  NaCl + BaSO4
b. H2 + N2  NH3
c.
C + H2O  CO + H2
d. AgNO3 + FeCl3  AgCl + Fe(NO3)3
i.
HNO3 + H3PO3  NO + H3PO4 + H2O
j.
HNO3 + I2  HIO3 + NO2 + H2O
k.
Na2S + AgNO3  Ag2S + NaNO3
l.
H+ + NO3- + Fe2+  H2O + NO + Fe3+
e.
H2CO3  H2O + CO2
m. FeBr2 + Br2  FeBr3
f.
MgSO4 + Ca(OH)2  Mg(OH)2 + CaSO4
n.
S2O32- + I2  S4O62- + I-
g.
H2O2 + PbS  PbSO4 + H2O
o.
H2O2 + MnO4-  O2 + Mn2+
p.
Zn + MnO2 + NH4Cl ZnCl2 + Mn2O3 + NH3 + H2O
h. KCl + H2SO4  KHSO4 + HCl
3. Balance the following equations using half reactions. Identify the oxidation half reactions and the reduction
half reactions.
a.
Cr2O72- + H+ + I-  Cr3+ + I2 + H2O
b. As2O3 + H+ + NO3- + H2O  H3AsO4 + NO
c.
H3AsO4 + Zn  AsH3 + Zn2+
MnO4- + H2SO3 + H+  Mn2+ + HSO4- + H2O
g.
I2 + H2SO3 + H2O  I- + HSO4- + H+
h. HgS + Cl- + NO3-  HgCl42- + S + NO
d. MnO42- + H+  MnO4- + MnO2
e.
f.
MnO4- + SO2  Mn2+ + SO42- + H+
i.
NO2 + OH-  NO2- + NO3-
j.
S + H+ + NO3-  SO2 + NO + H2O
4. Balance the following equations after putting them in ionic form.
a.
Cu(s) + HNO3(aq)  Cu(NO3)2(aq) + NO(g) + H2O(l)
b. Fe(NO3)2(aq) + HNO3(aq)  Fe(NO3)3(aq) + NO(g) + H2O(l)
c.
Zn(s) + HNO3(aq)  Zn(NO3)2(aq) + NO2(g) + H2O(l)
d. Sb(s) + H2SO4(aq)  Sb2(SO4)3(aq) + SO2(g) + H2O(l)
e.
H2S(g) + H2SO3(aq)  S(s) + H2O(l)
f.
HCl(aq) + HNO3(aq)  HClO(aq) + NO(g)
h. KI(aq) + O2(g)  KI3(aq) + H2O(l)
i.
HNO3(aq) + H2SO4(aq) + Hg(l)  Hg2SO4(s) + NO(g) + H2O(l)
j.
CO(g) + I2O5(g)  CO2(g) + I2(g)
Chemistry 12 - Redox – page 2