Oxidation – Reduction
a.k.a.
REDOX
Textbook Sections:
4.4-4.6 and 20.1-20.2
Good website:
http://www.wfu.edu/~ylwong/redox/
• Oxidation – the loss of one or more
electrons by a substance (element,
compound, ion)
• Reduction – the gain of one or more
electrons by a substance (element,
compound, ion)
O. I. L. R. I. G.
Oxidation Is Loss, Reduction Is Gain
(of electrons)
L.E.O. the lion says G.E.R.
Loss of Electrons is Oxidation
Gain of Electrons is Reduction
A-2
A-1
A
A+1
Oxidation --------->
 A-1 + electron
 A + electron
 A+1 + electron
 A+2 + electron
<--------- Reduction
• Redox reaction – a process where electrons
are transferred from one substance to
another
• How can you tell when a redox reaction is
taking place?
1) Assign oxidation numbers to atoms in
substances
2) Compare oxidation numbers before and
after reaction to determine if atom has
lost or gained electrons
Rules For Assigning
Oxidation Numbers
Use the Rules in Order
1. The oxidation number of an atom in an element is 0.
Examples: Na, H2, Br2, S8, Ne
Ox. #
0
0
0
0
0
2. The oxidation number of a monatomic ion is the
same as its charge.
Examples: Na+1, Ca+2, Al+3, Cl-1, O-2
Ox #
+1
+2
+3
-1
-2
3. The sum of the oxidation
numbers of all atoms in a neutral
compound is zero.
4. The sum of the oxidation
numbers of all atoms in an ion is
equal to the charge on the ion.
5. In compounds, fluorine is always assigned
an oxidation number of -1 (the most
electronegative element in a compound
always has a negative oxidation number.)
6. Hydrogen’s oxidation number will be
- +1 when bonded to a nonmetal (HCl)
- -1 when bonded to a metal (NaH)
Examples:
NaH
CaH2
HCl
Na—H
H—Ca—H H—Cl
-1 +2 -1 +1 -1
+1 -1
H2S
H—S—H
+1 -2 +1
7.Oxygen usually has an oxidation number of
-2. Exceptions: in peroxides, oxygen will be
-1 and when combined with only F, it will be
+2 and in O2 it will be 0
Examples:
H2O
CaO
H2O2
H—O—H
Ca—O
H—O—O—H
+2 -2
+1 -1 -1 +1
+1 -2 +1
O2-2
OF2
[O—O]-2
F—O—F
-1 +2 -1
-1 -1
8. Halogens usually have an oxidation
number of -1.
Examples:
NaCl
MgI2
OCl2
HOBr
Na—Cl, I—Mg—I, Cl—O—Cl, H—O—Br
+1 -1 -1 +2 -1 +1 -2 +1 +1 -2 +1
** If none of the above rules help you get started…look for a atom
with a known charge and use that charge as its oxidation number
CdS:
Cd-S
+2 -2
Use algebra to determine oxidation numbers of
"difficult" atoms.
Example: H2SO4
H is +1 * 2 = +2
O is -2 * 4 = -8
2 + S + -8 = 0
S is +6
Example:
ClO4-1
O is -2 * 4 = -8
-8 + Cl = -1
Cl is +7
Example: NH4+1
H is +1* 4 = +4
4+N = 1
N is -3
FeSO4
O is -2 *4 = -8
Recognize this as an ionic compound – sulfate has a -2 charge
For sulfate: x + -8 = -2
S = +6
Then look at the compound as a whole
Fe + 6 + -8 = 0
Fe = +2
C3H8
H is +1 * 8 = 8
3C + 8 = 0
3C = -8
C = - 8/3
oxidation numbers do NOT have to be integers
Once oxidation numbers have been
assigned, compare them before and after the
reaction.
4 Fe (s) + 3 O2 (g)  2 Fe2O3 (s)
0
0
+3 -2
Fe is oxidized, going from 0 to +3
O is reduced, going from 0 to -2
Notice that a total of 12 electrons were lost
and 12 electrons were gained
2 Fe2O3(s) + 3 C(s)  4 Fe(s) + 3 CO2(g)
+3 -2
0
0
+4 -2
Fe is reduced, going from +3 to 0
C is oxidized, going from 0 to +4
O undergoes no change
12 electrons lost and 12 electrons gained
As seen in the above examples, oxidation
and reduction ALWAYS occur together
Reducing agent (reductant)
• causes reduction
• loses electrons
• undergoes oxidation
• oxidation number increases
Oxidizing agent (oxidant)
• causes oxidation
• gains electrons
• undergoes reduction
• oxidation number decreases
Assign oxidation numbers, indicate what is
oxidized and reduced, indicate what is the
oxidizing agent and reducing agent
Ca(s) + 2 H+1(aq)  Ca+2(aq) + H2(g)
0
+1
+2
0
Ca is oxidized – increasing from 0 to +2
H+1 is reduced – decreasing from +1 to 0
Ca is the reducing agent
H+1 is the oxidizing agent
2 Fe+2(aq) + Cl2(aq)2 Fe+3(aq) + 2 Cl-1(aq)
+2
0
+3
-1
Fe+2 is oxidized – increasing from +2 to +3
Cl2 is reduced – decreasing from 0 to -1
Fe+2 is the reducing agent
Cl2 is the oxidizing agent
• In general,
• metals and anions act as reducing agents
(are oxidized) and
• nonmetals and cations act a oxidizing
agents (are reduced).
• Periodic table – in general, metals on left of
table are more active, metals become less
active as you move to the right side of the
table
Predicting Products of Redox
Reactions
The simple ones you already know:
• Decomposition (except of acids, bases,
carbonates & hydrates)
• Composition (except of two oxides)
• Combustion
• Replacement
Replacement Reactions
Replacement Reactions are redox reactions.
General pattern:
A + BX  AX + B
Mg(s) + 2 HCl(aq)  MgCl2(aq) + H2(g)
The Mg is oxidized and the H+ is reduced.
Fe(s) + Ni(NO3)2(aq)  Fe(NO3)2(aq) + Ni(s)
• The net ionic equation shows the redox
chemistry well:
Fe(s) + Ni+2(aq)  Fe+2(aq) + Ni(s)
Fe is oxidized to Fe+2
Ni+2 is reduced to Ni.
• Always keep in mind that whenever one
substance is oxidized, some other substance
must be reduced.
The Activity Series
• The activity series is a list of metals in order
of decreasing ease of oxidation.
• The metals at the top of the activity series
are called active metals and are easily
oxidized (they WANT to be ions)
• The metals at the bottom of the activity
series are called noble metals and NOT
easily oxidized (they WANT to be atoms)
Oxidation of copper metal
by silver ions.
(Dime Lab)
• A metal in the activity series can ONLY be
oxidized by a metal ion below it (metal
doing the replacing has to be above what it
is replacing in the activity series)
• If we place Cu into a solution of Ag+ ions,
then Cu+2 ions can be formed because Cu is
above Ag in the activity series:
• Cu(s)+2AgNO3(aq)Cu(NO3)2(aq)+2Ag(s)
or
Cu(s) + 2Ag+(aq)  Cu+2(aq) + 2Ag(s)
• This is only part of the story – more later !
• Special Cases:
1. Hydrogen reacts with a hot metallic oxide
to produce the metal element and water.
Ex: H2 + MgO  Mg + H2O
2. A metal sulfide reacts with oxygen to
produce the metallic oxide and sulfur
dioxide.
Ex: 2MgS + 3O2  2MgO + 2SO2
3. Chlorine gas reacts with dilute sodium
hydroxide to produce sodium hypochlorite,
sodium chloride, and water.
Cl2 + 2NaOH  NaClO + NaCl + H2O
4. Copper reacts with concentrated sulfuric
acid to produce copper(II) sulfate, sulfur
dioxide, and water.
Cu + 2H2SO4  CuSO4 + SO2 + 2H2O
5. Copper reacts with dilute nitric acid to
produce copper(II) nitrate, nitrogen
monoxide, and water.
Cu + HNO3  Cu(NO3)2 + NO + H2O
6. Copper reacts with concentrated nitric
acid to produce copper(II) nitrate, nitrogen
dioxide, and water.
Cu + HNO3  Cu(NO3)2 + NO2+ H2O
Memorize… These are reduction reactions (oxidation number decreases)
Reactants (oxidizing agents)
Products
MnO4- in acidic solution
Mn2+
MnO2 in acidic solution
Mn2+
MnO4- in neutral or basic solution
MnO2(s)
MnO4- in very basic solution
MnO42-
Cr2O72- in acidic solution
Cr3+
HNO3, concentrated
NO2
HNO3, dilute
NO
H2SO4, hot, concentrated
SO2
Metallic ion (higher charge)
Metallous ion (lower charge)
Elemental Halogen
Halogen ion
Na2O2
NaOH
HClO4
Cl-
C2O42-
CO2
H2O2
H2O
Memorize… These are oxidation reactions (oxidation number increases)
Reactants (Reducing Agents)
Products
Halogen ions
Halogen element
Metal element
Metal ion
SO32- or SO2
SO42-
NO2-
NO3-
Halogen element, dilute basic solution
Hypo-halogen-ite ion
(Ex: ClO-, BrO-)
Halogen element, concentrated basic
solution
Halogen-ate ion
(Ex: ClO3-, BrO3-)
Metallous ion (lower charge)
Metallic ion (higher charge)
H2O2
O2