Oxidation-Reduction Reactions
A.K.A Redox
Reduction-Oxidation Reactions
• Reactions that involve a change in charge of
atoms as they add or remove electrons in a
chemical reaction.
• Oxidation- Loss of electrons
(or gain in positive charge)
• Reduction- Gain of electrons
(or “reduction” of positive charge)
Oxidation State
• The formal oxidation state is the hypothetical
charge that an atom would have if all bonds to
atoms of different elements were 100% ionic.
• Oxidation- Results in an increase in its
oxidation state (more positive)
• Reduction- Results in a decrease (reduction) in
its oxidation state (more negative)
Simple Redox 1
• 2Fe + O2  2FeO
0
0
0
0
+2 -2
+2 -2
• Fe + Fe + O=O  FeO + FeO
• Iron was oxidized (it lost two electrons and
had a gain of positive charge from 0 to +2.)
• Oxygen was reduced (it gained two electrons
and became more negative from 0 to -2).
Why Call It Oxidation?
• Because Oxygen is very good at removing
electrons from other elements with its high
electron affinity and electronegativity.
• In addition, oxygen is far more common than
other highly electronegative elements like
fluorine and chlorine.
• Other elements can oxidize metals, but oxygen
is the most common!
Simple Redox 2
• Fe + CuSO4 → FeSO4 + Cu
• Fe0 + Cu+2 + SO4-2  Fe+2 + SO4-2 + Cu0
• The electrons simply transfer from the iron to the
copper.
2 e−
• Fe0  Fe+2
 Cu+2  Cu0
• In many ways, the Sulfate is irrelevant to this reaction.
Simple Redox 2
• Fe + CuSO4 → FeSO4 + Cu
• Fe0 + Cu+2 + SO4-2  Fe+2 + SO4-2 + Cu0
• Which one was oxidized and which one was
reduced?
2 e−
Cu+2  Cu0
• Fe0  Fe+2
• Iron was the one oxidized and Copper was
reduced?
Oxidizers
• Oxidizers are chemicals that have a strong
likelihood of causing oxidation of an atom.
• In other words, they accept electrons.
• Oxidizers include elements with high
electronegativity like F, O, Cl, and Br.
• Oxidizers also include those with high
oxidation states that want to gain electrons:
Mn(+7)O4-, Cr(+6)O3, Cr(+6)2O7-2
Reducers
• Elements that tend to give away electrons and
thus reduce other elements (while they
themselves being oxidized).
• Alkali and Alkaline Earth Metals are the best
examples.
• Any transition metal will give away electrons.
Is This A Redox Reaction?
• CuSO4 + 2NaOH  Cu(OH)2 + Na2SO4
• Polyatomic ions never change their charge, so
we can just ignore them.
• Cu+2 + Na+  Cu+2 + Na+
NO!!!
• Notice that Cu and Na never changed their
charges, so this is not a Redox reaction.
Formal Oxidation State
• In covalent bonds, the one with the higher
electronegativity gets a negative charge.
• So, in the case of CH4, since C is more
electronegative than H, then carbon has a
formal oxidation number of -4.
• Each hydrogen gets a +1.
Is Combustion a Redox Reaction?
-4 +1
0
+4 -2
+1 -2
• CH4 + O2  CO2 + H2O
• Even though electrons do not formally change
hands, due to these being covalent bonds, the
highly electronegative oxygens do cause a
change in oxidation state.
Formal Oxidation State
• In covalent bonds, the one with the higher
electronegativity gets a negative charge.
• So, in the case of C2H6, since C is more
electronegative than H, then each carbon has
a formal oxidation number of -3.
• However, the C-C bond is exactly equal, so
neither can claim that electron.
Determine the Oxidation State of Each
Element in this Reaction
+2 -2
+1 -2
+4 -2
0
• CO + H2O  CO2 + H2