Introduction to OxidationReduction Reactions
Electron Transfer Reactions
Types of Chemical
Reactions

There are four types of chemical reactions:





Acid/Base
Precipitation/Solubility
Complex Formation/Complex Dissociation
Oxidation/Reduction
Any chemical reaction consists of one (or more)
of these basic categories.
Oxidation/Reduction
Reactions

Acid/Base reactions


involve a donation /acceptance of protons
Precipitation/ Solubility reactions

involve a donation/ acceptance of negative
charge
what is being donated and accepted in a
redox reaction?
Oxidation/Reduction
Reactions


Electrons!
Consider the reaction taking place in a
disposable battery:
2Zn + 3MnO2  Mn3O4 + 2ZnO
How can you tell that electrons are being
donated and accepted? Which species is
donating electron( s) and which is accepting
electron (s)?
REDOX REACTIONS
Redox reactions are characterized by
ELECTRON TRANSFER between an
electron donor and electron acceptor.
REDOX REACTIONS
Transfer leads to—
1.
increase in oxidation number of
some element = OXIDATION
2. decrease in oxidation number of
some element = REDUCTION
Electron Transfer in Redox
Reactions

Oxidation



Reduction



Loss of electrons
Gain in oxygen
Gain of electrons
Loss of oxygen
“LEO the lion goes Ger”
Example


The reaction of a metal and non-metal
All the electrons must be accounted for!
Mg
+
S
→
Mg
2+ +
2S
Oxidation-Reduction

Oxidation means an increase in oxidation state - lose
electrons.

Reduction means a decrease in oxidation state - gain
electrons.

The substance that is oxidized is called the reducing
agent.

The substance that is reduced is called the oxidizing
agent.
Assigning Oxidation States

An Oxidation-reduction reaction
involves the transfer of electrons.

You should memorize these rules
Rules for Oxidation States

The charge the atom would have in a molecule
(or an ionic compound) if electrons were
completely transferred.
 The oxidation state of elements in their
standard states is zero.
 Example: Na, Be, K, Pb, H2, O2, P4 = 0
Assigning Oxidation States
 Oxidation state for monatomic ions are the same as
their charge.
 Example: Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2
 Oxygen is assigned an oxidation state of -2 in its
covalent compounds except as a peroxide.
Rules for Oxidation States
4. The oxidation number of hydrogen is +1
except when it is bonded to metals in binary
compounds. In these cases, its oxidation
number is –1.
5. Group IA metals are +1, IIA metals are +2 and
fluorine is always –1.
6. The sum of the oxidation numbers of all the
atoms in a molecule or ion is equal to the
charge on the molecule or ion.
Practice in Oxidation States

Assign the oxidation states to each
element in the following.

K2SO4
NO3H2SO4
Fe2O3
Fe3O4




Identify the

Oxidizing agent

Reducing agent

Substance oxidized

Substance reduced

On the worksheet
Types of Oxidation-Reduction Reactions
Combination Reaction
A+B
C
0
+4 -2
0
S + O2
SO2
Decomposition Reaction
C
+1 +5 -2
2KClO3
A+B
+1 -1
0
2KCl + 3O2
Types of Oxidation-Reduction Reactions
Displacement Reaction a.k.a Single Replacement
A + BC
0
+1
+2
Sr + 2H2O
+4
0
TiCl4 + 2Mg
0
AC + B
-1
Cl2 + 2KBr
0
Sr(OH)2 + H2 Hydrogen Displacement
0
+2
Ti + 2MgCl2
-1
Metal Displacement
0
2KCl + Br2
Halogen Displacement
The Activity Series for Metals
Hydrogen Displacement Reaction
M + BC
AC + B
M is metal
BC is acid or H2O
B is H2
Ca + 2H2O
Ca(OH)2 + H2
Copper Demonstration


Copper Pennies reacting with nitric
acid.
Can you figure out the equation?
Types of Oxidation-Reduction Reactions
Disproportionation Reaction
Element is simultaneously oxidized and reduced.
0
Cl2 + 2OHChlorine Chemistry
+1
-1
ClO- + Cl- + H2O
Classify the following reactions.
Ca2+ + CO32NH3 + H+
Zn + 2HCl
Ca + F2
CaCO3
NH4+
ZnCl2 + H2
CaF2
Half-Reactions

All redox reactions can be thought of as
happening in two halves.

One produces electrons - Oxidation half.

The other requires electrons - Reduction
half.
Half-Reactions

Write the half reactions for the following.

Na + Cl2 → Na+ + Cl-

SO3- + H+ + MnO4- → SO4- + H2O + Mn+2
Balancing Redox Equations

In aqueous solutions the key is the
number of electrons produced must be
the same as those required.

For reactions in acidic solution an 8 step
procedure.
Balancing Redox Equations

Write separate half reactions

For each half reaction balance all
reactants except H and O

Balance O using H2O
Acidic Solution

Balance H using H+

Balance charge using e-
Acidic Solution

Multiply equations to make electrons
equal

Add equations and cancel identical
species

Check that charges and elements are
balanced.
Practice

Balance the following reactions:

Sn 2+ (aq) + 2Fe 3+ → Sn 4+ (aq) + 2Fe 2+

MnO4- (aq) + C2O4-2 (aq) → Mn2+
(aq) + CO2 (g)
Practice



The following reactions occur in
aqueous solution. Balance them
Cr(OH)3 + OCl- + OH- 
CrO4-2 + Cl- + H2O
MnO4- + Fe+2 Mn+2 + Fe+3
Now for a tough one

Fe(CN)6-4 + MnO4- Mn+2 + Fe+3 + CO2 + NO3-
Basic Solution

Do everything you would with acid, but add
one more step.

Add enough OH- to both sides to neutralize
the H+

CrI3 + Cl2  CrO4- + IO4- + Cl-

CN- + MnO4- → CNO- + MnO2
Redox Titrations

Same as any other titration.

the permanganate ion is used often
because it is its own indicator. MnO4- is
purple, Mn+2 is colorless. When reaction
solution remains clear, MnO4- is gone.

Chromate ion is also useful, but color
change, orangish yellow to green, is