Oxidation and Reduction
Reactions
Oxidation (Read only)
Original definition:
When substances combined with oxygen.
Ex:
All combustion (burning) reactions
CH4(g) + 2O2(g)
CO2(g) + 2H2O(l)
All “rusting” reactions
4Fe(s) + 3O2(g)
2Fe2O3(s)
Reduction (Read Only)
Original Definition:
Reaction where a substance “gave up” oxygen.
Called “reductions” because they produced
products that were “reduced” in mass because
gas escaped.
Ex:
2Fe2O3(l) + 3C(s)
4Fe(l) + 3CO2(g)
Oxidation/Reduction
Deals with movement of ELECTRONS
during a chemical reaction.
(Oxygen doesn’t have to be present)
Electron Transfer Reactions
Oxidation: LOSS of one or more electrons.
Reduction: GAIN of one or more electrons
Electron Transfer Reactions
Oxidation & reduction always occur together.
Electrons travel from what is oxidized towards
what is reduced.
One atom loses e-, the other gains e-
Redox Reactions:
ALWAYS involve changes in charge
A competition for electrons between atoms!
Remember!!
Or…Remember
Conservation of “Charge”
Total electrons lost = Total electrons gained
Oxidizing/Reducing Agents
Oxidizing Agent:
substance reduced
 Gains electrons
Reducing Agent:
substance oxidized
 Loses electrons
The “Agent” is the “opposite”
Assigning Oxidation Numbers
Practice Problems
http://www.usca.edu/chemistry/genchem/oxn
umb.htm
Animation of Oxidation and Reduction
http://www.ausetute.com.au/redox.html
Identify What is Changing in
Charge
What is oxidized and reduced?
What are the oxidizing and reducing agents?
Ex:
3Br2 + 2AlI3
2AlBr3 + 3I2
0
+3 -1
3Br2 + 2AlI3
+3 -1
0
2AlBr3 + 3I2
Br2 is reduced and is the oxidizing agent
I-1 is oxidized and is the reducing agent
What is oxidized and reduced?
What are the oxidizing and reducing agents?
Mg + CuSO4
2K + Br2
Cu + 2AgNO3
MgSO4 + Cu
2KBr
Cu(NO3)2 + 2Ag
NOTE:
Atoms in a polyatomic ion DO NOT change in charge!
0
+2
+2
Mg + CuSO4
0
MgSO4 + Cu
Mg oxidized (reducing agent)
Cu+2 reduced (oxidizing agent)
0
0
+1 -1
2K + Br2
2KBr
K oxidized (reducing agent)
Br2 reduced (oxidizing agent)
0
+1
Cu + 2AgNO3
Cu oxidized (reducing agent)
Ag+1 reduced (oxidizing agent)
+2
0
Cu(NO3)2 + 2Ag
Redox or Not Redox
(that is the question…)
Redox Reactions: must have atoms changing in
charge.
Not all reactions are redox.
Easy way to spot a redox reaction!!!
Look for elements entering and leaving
compounds.
Is it Redox?
Look for Changes in Charge!
Are elements entering and leaving compounds?
Synthesis:
Ex:
2H2 + O2
2H2O
Decomposition:
Ex:
2KClO3
2KCl + 3O2
Is it Redox?
Synthesis: YES
0
0
+1 -2
Ex: 2H2 + O2
2H2O
Decomposition: YES
+1 +5 -2
Ex: 2KClO3
+1 -1
0
2KCl + 3O2
Is it Redox?
Combustion:
CH4 + 2O2
CO2 + 2H20
Single Replacement:
Zn + CuCl2
ZnCl2 + Cu
Is it Redox?
Combustion: YES
-4 +1
0
CH4 + 2O2
+4 -2
CO2 + 2H20
Single Replacement:
0
+2 -1
Zn + CuCl2
+1 -2
+2 -1
YES
0
ZnCl2 + Cu
Is it Redox?
Double Replacement:
AgNO3 + LiCl
AgCl + LiNO3
Is it Redox?
Double Replacement: NO!!!!
Ions switch partners, but don’t change in charge
+1 +5 -2
+1 -1
AgNO3 + LiCl
+1 -1
+1 +5 -2
AgCl + LiNO3
Remember charges of atoms inside polyatomic ions
do not change!
Writing Half Reactions
Redox Reactions are composed of two parts
or half reactions.
Half Reactions Show:
Element being oxidized or reduced.
Change in charge
# of electrons being lost or gained
Writing Half Reactions
0
0
+1 -1
2Na + F2
Oxidation:
or
2NaF
Na
2Na
Na+1 + 1e2Na+1 + 2e-
Note: e- are “lost” (on the right of arrow)
Reduction:
or
F + 1eF2 + 2e-
F-1
2F-1
Note: e- are “gained” (on the left of arrow)
Ox’s Have Tails!!

Oxidation Half reactions always have
“tails” of electrons
Na
Na+1 + 1e-
0
+2 -1
+2
Zn + CuCl2
-1
0
ZnCl2 + Cu
Zn+2 + 2e-
Ox:
Zn
Red:
Cu+2 + 2e-
Cu
Balancing Simple Redox Rxns
Must be:
Balanced for Mass
ATOMS balance
Balanced for Charge
Total e- Lost = Total e- Gained
Balancing Harder
Redox Reactions
(Honors)
Oxidation Number Method
(Balancing in Acid Solution)
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Find ox #’s and use brackets to connect elements
changing in charge.
Balance atoms changing in charge
Find total e- involved in each change
If necessary balance e- by multiplication
Balance all other atoms except H and O
Balance oxygen by adding H2O to side deficient
Balance hydrogen by adding H+1 to side deficient
Check for balance with respect to atoms and charge.
Half Reaction Method
(Ion/Electron Method)
(In acid solution)
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Separate equation into two “basic” half reactions
Balance all atoms except H and O
Balance oxygen by adding H2O
Balance hydrogen by adding H+1
Balance charge by adding electrons to more
positive side
If necessary balance e- by multiplication
Add together half reactions and simplify
Check for balance of atoms and charge
Applications of Redox Reactions
Corrosion of Metals
the metal gets oxidized
forming metal oxides on
the surface
Prevention: Use paint, oil,
plating or attach to negative
terminal of a battery.
Gold doesn’t
rust…Why?
Photograph Development involves oxidation and
reduction of silver atoms and ions
Bleach acts on
stains by oxidizing
them, getting reduced
in the process
Explosives form
neutral gases like N2
from compounds!
Reactivity of Metals
Reference Table J
Metals Higher on Table J
are more ‘active”
It is easier for more “active”
metals to be oxidized
or lose electrons.
Copper replaces
silver!
Cu0(s) + AgNO3(aq)
Ag0(s) + CuNO3(aq)
Ag0(s) + CuNO3(aq)
wouldn’t happen!!!
Reactivity of Nonmetals
Reference Table J
Nonmetals higher on Table J
are more “active”
It is easier for more “active”
nonmetals to “gain” electrons
and be reduced.
Electrochemical Cells (Batteries)
Chemical reaction that produces electricity.
Called “voltaic cells” as they produce voltage
This happens SPONTANEOUSLY.
Moving Electrons = Electricity
Electrons given off by oxidized substance
travel towards substance being reduced.
Traveling electrons move
through “external circuit”
where they do work.
How do the Electrons Move?
Batteries often contain 2 metals.
Start with Table J
Electrons travel from the more “Active metal”
toward the less active metal.
Metal above = oxidized
Ion on Metal below = reduced
Electrons flow
“Down Table J”
e-
From metal
above to ion of
metal below
Parts of a Simple Battery
(Voltaic Cell)
Made of Two “Half Cells” containing:
2 Metal Electrodes
2 Solutions of Ions
External Wire
Salt Bridge
Electrons need to flow in a “circuit” that
is connected.
External Wire:
allows electrons to flow
between metal electrodes
Salt Bridge: allows ions
to flow between solutions
Zn/Zn+2//Cu+2/Cu
What is Ox/Red?
See Table J
Metal above is oxidized
Zn
Ion of metal below reduced
Cu+2
Which way do electrons
flow in the external wire?
See Table J
Electrons flow “Down” the
table from what is oxidized
towards what is reduced.
from Zn to Cu
e-
Which electrode is negative?
Which electrode is positive?
Electrons flow from negative to
positive electrode.
eNegative electrode: Zn
Positive electrode: Cu
Which electrode is the
anode and cathode?
Anode: metal electrode
where oxidation occurs
Zn
Cathode: metal electrode
where reduction occurs
Cu
Remember
AN OX
RED CAT
Anode is where oxidation happens
Cathode is where reduction happens
What are the Half Reactions?
What is the Net Equation?
Ox:
Zn0
Zn+2 + 2e-
e-
Red:
Cu+2 + 2e-
Cu0
Net: (add ½ reactions)
Zn0 + Cu+2
Zn+2 + Cu0
Make sure final net equation is
balanced for electrons and atoms!
Which electrode gains/loses weight?
Look at half reactions!!
Which one forms solid metal?
Which one forms dissolved ions?
Ox:
Zn0
Zn+2 + 2eRed:
Cu+2 + 2eCu0
Zinc electrode loses mass
Copper electrode gains mass
Which way to do the ions in the salt
bridge “migrate” or move?
Remember:
“The negative ions complete the circuit”
(The ions actually end up moving towards
the solution of opposite charge that forms.)
Follow the ions

http://www.mhhe.com/physsci/chemistry/e
ssentialchemistry/flash/galvan5.swf
Dead Battery
Voltage = 0
Means the reaction in the battery has
reached EQUILIBRIUM.
You try it…
Mg/Mg+2//Al+3/Al
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Draw and label Battery
What is oxidized/reduced?
What are the half reactions and net(balanced)?
What is the neg/pos electrode?
What is the anode/cathode?
Which way do e- flow in wire?
Which way do -/+ ions flow in salt bridge?
Which electrode gains/loses mass?
Finding Voltage of a Battery
(Honors)
Use Voltage Table
Find your half reactions and record voltage
Note:
All ½ reactions shown are reductions.
For oxidation, reverse the sign of the voltage
Nerntz Equation (Honors)
Find voltage of a battery when the conc. of dissolved
ions is not 1 Molar (as on “standard voltage” table)
Ecell = E0 –
0.0592 log [product ion]x
n
[reactant ion]y
n = total # of moles electrons being transferred
The concentration of dissolved ions can affect voltage.
Greater concentration of reactant ions (see net)
increases the overall voltage.
Electrolytic Cells &
Electrolysis Reactions
Uses electricity to “split” or “lyse” a
compound into it’s neutral elements
An outside electrical source provides electrons
to force a non-spontaneous redox reaction to
occur.
Electrolysis Set Up
Single Cell filled with electrolyte with +/- ions
Attach battery to two electrodes.
Electrodes are made of an inert substance
(like platinum or graphite) that conducts.
Electrodes don’t chemically change like in a battery,
they just provide current
Role of the Battery
Pulls electrons off one electrode
Making it POSITIVE
Adds electrons onto one electrode
Making it NEGATIVE
Which Way do the Ions Move?
To electrode of
opposite charge
What is Oxidized/Reduced?
At neg. electrode
electrons are gained by ion
(reduction at CATHODE)
At positive electrode
electrons are lost by ion
(oxidation at ANODE)
Remember
AN OX
RED CAT
Anode is where oxidation happens
Cathode is where reduction happens
Half Reactions & Net Equation
Rxn at Anode: (Ox)
ClCl + 1eOr more correctly
2ClCl2 + 2e-
DIATOMIC!!!!!!
Rxn at Cathode: (Red)
Na+ + 1eNa
(Multiply by 2 to balance electrons)
NET:
2Na+ + 2Cl-
2Na + Cl2
Determining Voltage Needed
Use the Voltage Table to determine the total
voltage needed to run the Electrolytic cell.
Total voltage should be a NEGATIVE number
Electrolysis of Molten NaCl (l)
Electrolysis of PbCl2(l)
What is oxidized?
What is reduced?
What are the ox/red
half reactions?
Negative
Electrode
Positive
Electrode
What is the net
equation?
Electrolysis of PbCl2(l)
Oxidized: ClHalf Reactions
Ox:
Cl-1
2Cl-1
Reduced: Pb+2
Cl + 1eCl2 + 2e-
Red:
Pb+2 + 2e-
Pb
Net:
Pb+2 + 2Cl-1
Pb + Cl2
Electrolysis of NaCl(aq)
Electrolysis of NaCl (aq)
Electrolysis of Water
At Positive Electrode:
Ox: O-2
O + 2ebut there is a diatomic!
2O-2
O2 + 4eAt Negative Electrode
Red: H+1 + 1eH
but there is a diatomic!
2H+1 + 2eH2
Net: 2H2O
2H2 + O2
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Lemon Battery Demo
http://youtu.be/AY9qcDCFeVI
Electrolysis of Copper Sulfate
http://youtu.be/xBz9HJ32Ouw
Electrolysis of Water/ Silver Nitrate and Cu reaction
http://youtu.be/Bcfp8VtcrSA
Electrolysis of Water (Animation)
http://youtu.be/2t13S-KpGeE
Electrolysis of Water (Simple)
http://youtu.be/HQ9Fhd7P_HA
Electroplating
Electrolysis reaction
used to coat a
substance with a thin
layer of metal.
Often coating is a less
reactive metal that is
not easily oxidized or
corroded.
Electroplating

Negative Electrode
Is the OBJECT TO BE
PLATED
 so the positive metal ions
would go towards it and be
REDUCED.
 It is the CATHODE

Red: Ag+ + 1e-
Ag0
Electroplating

Positive Electrode
Made of plating metal
 It dissolves into solution as
metal strip gets OXIDIZED.
 It is the ANODE
 This replenishes the ions for
plating.

Ox: Ag0
Ag+ + 1e-
Electroplating Problems
(Honors)
Coulomb = measure of electrical charge
1 mole e- = 96,500 coulombs
# coulombs = # amps x seconds
Electroplating Problems
(Honors)
Reduction:
Happens on object to be plated
Look at Reduction half reaction
Look at mole relationships
between electrons and metal atoms.
Ex: Ag+ + 1e-
Ag0
Electroplating Problems
(Honors)

You can now answer questions regarding
the amount of a substance in moles or
grams that can be electroplated over a
certain amount of time.
Electroplating Problems
(Honors)
If 10 amps are run through a CuSO4 solution for
5 minutes, calculate the grams of Cu that will plate
onto the spoon.
We Know:
1 mole e- = 96,500 coulombs
# coulombs = # amps x seconds
Red:
Cu+2 + 2e-
1 mole Cu = 63.5 grams
Cu0
So….Let’s start here
# coulombs
= 10 amps x 300 seconds
= 3000 coulombs
3000 coul. x 1 mole e- x 1 mole Cu x 63.5g Cu = .987 grams
96,500 coul 2 mole e1 mole Cu
Mole ratio from
Reduction half reaction
You Try One
How long will it take to deposit 20 grams of
silver from a solution of AgCl onto a copper
tray if a current of 5 amps is used?
Answer = 3, 574 sec
or 59.5 minutes or about 1 hour
You Try One
How many amps are needed to deposit .504g.
of Iron in 40 minutes by passing a current
through a solution of Iron II Sulfate?
Answer: .72 amps