Oxidation and Reduction
Reactions and Electrochemistry
“The Ubiquitous Electron”

Redox and Iron in your Body
Types of Reactions
1.
2.
Ions or molecules react w/ no apparent
change in electronic structure (ex.
Double displacement)
Ions or atoms undergo changes of
electronic structure, the way e- transfer
or the way atoms share e- changes.
Oxidation- Reduction Reaction
Definition:



Chemical change that occurs when
electrons are transferred between
reactants
All oxidation reactions are accompanied
by reduction reactions
Important: in the corrosion of metals,
sources of energy, life processes
Oxidation

Part of the redox rxn in which electrons
are removed or apparently removed from
an atom (loss of electrons  atom gets
more positively charged)

Movie
Reduction

Part of the redox rxn in which electrons
are added or apparently added to an atom
(gain of electrons atoms get more
negatively charged)

Movie
OIL RIG


Oxidation Is Losing
Reduction Is Gaining
LEO the lion goes GER


Loss of Electrons in Oxidation
Gain of Electrons in Reduction
Ionization or Solvation =
the process of
surrounding solute
particles with solvent
particles to form a
solution
 Video
 “like dissolves like”
Net Ionic Equations




When reactions take place in water
chemists write the equation in ionic form
(particles ionize – break into their ions in
water)
Chemists only write down the ions that
take part in the reaction
Spectator ions- ions that aren’t involved
in the reaction (chemists don’t write
these)
Makes rxn easier to balance
Cu + NO3-1  Cu+2 + NO

Show chemistry connections video: 7:36
minutes into video, found in redox folder
Rules for Assigning Oxidation
Numbers:

Use oxidation numbers (charges on
atoms) to determine which atom
underwent reduction and which atom
underwent oxidation
Rules:
1.
The oxidation number for any free
element is 0 (zero). Also any diatomic
molecule is 0 (zero)
H2, O2, I2, Cl2, F2, N2, Br2
Fe = 0 charge
O2 = 0 charge
2.
The oxidation number of any
monoatomic ion is equal to the charge
written on the ion.
Na +1 = +1
Cl-1 = -1
3.
Oxidation number of hydrogen in most
of its compounds is +1 (except for LiH
then H is –1)
+1
Ex. HCl
4.
Oxidation # of oxygen in most of its
compounds is –2.(except peroxides= -1)
-2
Ex. H2O
-1
Ex. H2O2
5.
Sum of the oxidation numbers of all of
the atoms must equal the apparent
charge of that particle.
Ex. H2SO4
-zero charge
+1 ? -2
H2SO4
+2 +6-8=0
S= +6
Ex. NO3 –1
? + -2(3) = -1
+5 + (-6) = -1
N= +5
6.
Group 1 +1
Group 2  +2
Aluminum & Boron  +3
Group 17  -1
Ex. KMnO4
K= +1
Mn = +7
O = -2

Page 174 #67, 69

Identifying redox, chemistry connections
11:29 minutes in
Identifying Redox Reactions


First, figure out the oxidation numbers of
all elements in the reaction
If oxidation number changes as you move
from reactants to products it is REDOX.
This is REDOX, Mg- loss e(oxidation), H –gained e-(reduction)
This is NOT REDOX

P 618 in modern chem- #2, 15
Oxidizing & Reducing Agents



Think of these agents as “causers” of
redox rxns
Look at reactants
Some substances are better oxidizing or
reducing agents

Reducing Agents: substance that donates
the electron (contains the atoms that are
oxidized- or loss the e-)
• Causes the reduction to occur

Oxidizing Agent: substance that gains
the e- (contains the atoms that are
reduced or gains e-)
• Causes oxidation to occur
Ex.
0
0
+3
-2
4Al + 3O2  2Al2O3
Al- lost e- , oxidized
-reducing agent
O- gained e-, reduced
-O2 is the oxidizing agent
Balancing Redox Reactions
--Half Reaction Method


Half Reaction: equation that shows just
the oxidation or reduction part of the rxn.
In balancing we balance each of the half
rxns first, then add them together &
reduce
Steps:
1.
2.
3.
4.
Place oxidation #’s on everything after
it is in the net ionic form.
ID the oxidation ½ rxn and the
reduction ½ rxn
Write out the ½ rxns.
Balance the atoms by placing
coefficients in front of the atoms 
except for H and O
Ex. Cl2  Cl-1 become Cl2  2Cl-
5. Place the # of electrons lost on the
product side of oxidation ½ rxn, place #
of electrons gained on reactant side of
reduction ½ rxn
6. To balance hydrogens and oxygens:
Acidic soln: add H+ & H2O
Basic soln: add OH- & H2O
7.
8.
Balance the charges (# e- lost must
equal # e- gained) by using a least
common multiple ( multiply the whole
½ rxn)
Add two ½ rxns together and reduce if
necessary.

Chemistry connections- balancing with
blood alcohol tests (21:00-26:00)
Electrochemistry
Movie


Because redox reactions involve electron
transfer, the release or absorption of energy
can occur in the form of electrical energy
rather than heat
Electrochemistry is the branch of chemistry
that deals w/ electricity related applications
of redox reactions
Electrochemical Process



Conversion of chemical to electrical
energy
Ex. Flashlight batteries, biological
systems, electroplating
If the substance that is oxidized is
separated from the substance that is
reduced you get an energy transfer of
electrical energy instead of heat


Electrons can be transferred from one side
to the other through a connecting wire
Electric current moves in a circuit (while
the electrons are being balanced by the
movement of ions in solution)
Part of a Cell:

Electrodes:
–
Conductor in a circuit that carries electrons
from one substance to another
– Anodes: electrode where oxidation occurs,
anions (-) are attracted to this when they are
oxidized by losing electrons (the positive
electrode)
– Cathode: electrode where reduction occurs,
cations (+) are attracted to this when they are
reduced by gaining electrons (negative
electrode)

Salt Bridge:
–
Porous partition that separates the 2 half
reactions
– Contains a conducting solution that allows the
passage of ions from one compartment to the
other w/ out mixing the solutions in the half
reactions

Half Cell:
–
Part of the voltaic cell in which either oxidation
or reduction occurs
– The two half cells together make a complete
electrochemical cell

Ex. Oxidation half cell
–
Zn  Zn+2 + 2 e• (zinc rod in zinc sulfate)

Reduction half cell
–
Cu+2 + 2e-  Cu
• (copper rod in copper sulfate)

Complete Cell Notation
Anode electrode |anode solution || cathode solution |cathode electrode
(the double line || represents the salt bridge)
Ex. Zn (s) | Zn +2 (aq) || Cu+2 (aq) | Cu (s)
e-
e-
e-
e-
Anode-positive
electrode, oxid.
occurs
Zn
rod
ZnSO4
Salt
bridge
Cu
rod
CuSO4
Zn(s) |ZnSO4(aq)||CuSO4 (aq)| Cu (s)
Cathodeneg.
electrode,
red. occurs

Fuel Cell
Type of Cells

Dry Cell: voltaic cell in which the electrolyte
(conducting solution) is a paste
–
–
–
–
Generates direct current by converting chemical to
electrical energy by a spontaneous redox reaction
Also called galvanic cells or voltaic cells
Ex. Batteries (zinc-carbon, alkaline, mercury)
Ex. Flashlight battery (zinc-carbon)
• Zinc container (anode) filled w/ a moist paste (salt paste) made
of MnO2, ZnCl2, NH4Cl and water w/ a graphite rod (cathode)
embedded into it
–
Alkaline batteries (do not have a carbon rod
cathode which allows them to be smaller- uses
a graphite/ MnO2 mix)
– Mercury (cathode is HgO/carbon mix)
–
Lead storage batteries
• Group of cells that are connected together
• Can be recharged (use in a car)
• Ex. 12 V battery- 6 voltaic cells connected together
–
–
–
–
Each cell contains 2 lead electrodes or grids
Anode- grid packed w/ spongy lead
Cathode – grid packed w/ PbO2
Immersed in 5M H2SO4
• Recharging occurs whenever the car is running
• Doesn’t last forever- byproduct PbSO4 falls from
electrodes and collects on bottom (loses too much
lead)
–
Fuel Cells
• A voltaic cell in which the reactants are being
continuously supplied and the product are being
continuously removed
• A fuel substance undergoes oxidation, from which
electrical energy is obtained continuously
• No recharging, no pollution
• Ex. H-O cell: submarines, military vehicles, Apollo
Electrical Potential



In a voltaic cell, the oxidizing agent at the
cathode pulls the electrons through the wire
away from the reducing agent at the anode
The “pull” on the electrons is called the
electric potential
Electrical potential is measured in volts (V)
Electrode potential:



The potential difference measure across the
complete voltaic cell is easily measured
It equals the sum of the electrode potentials for
each of the two half-reactions
The individual electrode potential for a halfreaction cannot be measured directly, but it can be
measured by connecting to a standard half-cell as
a reference (we use a Hydrogen electrode that is in
a 1.0M acidic solution at 1 atm and 25 C)
Standard Reduction Potentials
(p. 796-book)



Electrode potentials are always written as
reductions
The more negative the voltage oxidation
(stronger reducing agent)
The more positive the voltage reduction
(stronger oxidizing agent)
Standard Cell Potential (E° cell)

Use this formula:
E°cell = E°reduction - E°oxidation
or
E°cell = E°cathode - E°anode

A spontaneous reaction will have positive
value for E° cell
Zn (s) | Zn +2 (aq) || Cu+2 (aq) | Cu (s)

Oxidation: Zn+2 + 2 e-  Zn
– E°Zn +2 = -.76 V

Reduction: Cu +2 + 2e-  Cu
– E°Cu+2 = .34V
E°cell = E°reduction - E°oxidation
=.34V - (-.76V)
=1.10V
Cu+2 + Zn  Cu + Zn+2

Zn (s) | Zn +2 (aq) || Fe+2 (aq) | Fe (s)
(anode)
(cathode)

Oxidation: Zn+2 + 2 e-  Zn
– E°Zn +2 = -.76 V

Reduction: Fe +2 + 2e-  Fe
– E°Fe+2 = -.44V
E°cell = E°reduction - E°oxidation
=-.44V - (-.76V)
=.32V
Fe+2 + Zn  Fe + Zn+2

Practice





Mn| Mn +2 || Br2 | BrH2C2O4| CO2 || MnO4-1 | Mn+2
Ni | Ni +2 || Hg2+2 | Hg
Cu | Cu+2 || Ag+1 | Ag
Pb| Pb +2 || Cl2 | Cl-

Mn| Mn +2 || Br2 | Br–
–

Ecell= 1.07-(-1.18)= 2.25 V
Br2 +Mn  Mn+2 + 2Br-
H2C2O4| CO2 || MnO4-1 | Mn+2
–
E cell= 1.51- (-.49) = 2.00 V
– 2 MnO4- +6 H+ + 5 H2C2O4  2Mn+2 + 8H2O + 10 CO2

Ni | Ni +2 || Hg2+2 | Hg
–
1.04 V
– Ni + Hg2 +2  Ni +2 + 2Hg

Cu | Cu+2 || Ag+1 | Ag
–
.46 V
– Cu + 2 Ag+  Cu+2 + 2 Ag

Pb| Pb +2 || Cl2 | Cl–
–
`1.49V
Pb +Cl2  Pb+2 + 2Cl-



Video
How its made nails
Corrision Pics
Redox/ Electrochemistry Quest
(anode song)
Redox
–
–
–
–
–
oxidation #’s
ID if redox or not
Oxidizing or reducing agent (strengths)
Balancing- set up ½ rxns
Balancing oxygens/hydrogens
• Acids (add H+ and H20)
• Bases (add OH- and H20)

Electrochem
–
What is an electrochemical cell
– Example
– Parts: anode, cathode, salt bridge, what each
part does)
– Standard cell potential (getting the voltage and
write equation)