Galvanic Cell

Electrochemistry
Applications of Redox

Review
• Oxidation reduction reactions involve a
transfer of electrons.
• OIL- RIG
• Oxidation Involves Loss
• Reduction Involves Gain
• LEO-GER
• Lose Electrons Oxidation
• Gain Electrons Reduction

Applications
• Moving electrons is electric current.
• 8H + +MnO 4 - + 5Fe +2 +5e -
Mn +2 + 5Fe +3 +4H 2 O
• Helps to break the reactions into half
reactions.
• 8H + +MnO 4 - +5e - Mn +2 +4H 2 O
• 5(Fe +2 Fe +3 + e - )
• In the same mixture it happens without
doing useful work, but if separate

• Connected this way the reaction starts
• Stops immediately because charge builds
up.
H +
MnO 4
-
Fe +2

Galvanic Cell
Salt
Bridge
allows
current
to flow
H +
MnO 4
-
Fe +2

• Electricity travels in a complete circuit
H +
e -
MnO
-
Fe +2
4

•Instead of a salt bridge
Porous
Disk
H +
MnO 4
-
Fe +2

e - e - e - e -
Anode
e -
Reducing
Agent
Cathode
e -
Oxidizing
Agent

Cell Potential
• Oxidizing agent pulls the electron.
• Reducing agent pushes the electron.
• The push or pull (“driving force”) is called
the cell potential E cell
• Also called the electromotive force (emf)
• Unit is the volt(V)
• = 1 joule of work/coulomb of charge
• Measured with a voltmeter

0.76
H 2 in
Anode
Zn +2 H +
SO
-2 Cl -
4
1 M ZnSO 4
Cathode
1 M HCl

Standard Hydrogen Electrode
• This is the reference
all other oxidations
are compared to
•Eº = 0
• º indicates standard
states of 25ºC,
1 atm, 1 M
solutions.
H +
Cl -
1 M HCl
H 2 in

Cell Potential
• Zn(s) + Cu +2 (aq) Zn +2 (aq) + Cu(s)
• The total cell potential is the sum of the
potential at each electrode.
• Eº cell = Eº Zn Zn +2 + Eº Cu +2 Cu
• We can look up reduction potentials in a
table.
• One of the reactions must be reversed,
so change it sign.

Cell Potential
• Determine the cell potential for a galvanic
cell based on the redox reaction.
• Cu(s) + Fe +3 (aq) Cu +2 (aq) + Fe +2 (aq)
• Fe +3 (aq) + e - Fe +2 (aq)
Eº = 0.77 V
• Cu +2 (aq)+2e - Cu(s)
• Cu(s) Cu +2 (aq)+2e -
Eº = 0.34 V
Eº = -0.34 V
• 2Fe +3 (aq) + 2e - 2Fe +2 (aq) Eº = 0.77 V

Reduction potential
• More negative Eº
– more easily electron is added
– More easily reduced
– Better oxidizing agent
• More positive Eº
– more easily electron is lost
– More easily oxidized
– Better reducing agent

Line Notation
• solidAqueousAqueoussolid
• Anode on the leftCathode on the right
• Single line different phases.
• Double line porous disk or salt bridge.
• If all the substances on one side are
aqueous, a platinum electrode is
indicated.

•For the last reaction
•Cu(s)Cu +2 (aq)Fe +2 (aq),Fe +3 (aq)Pt(s)
Cu 2+ Fe +2

In a galvanic cell, the electrode that
acts as a source of electrons to the
solution is called the __________;
the chemical change that occurs at
this electrode is called________.
a. cathode, oxidation
b. anode, reduction
c. anode, oxidation
d. cathode, reduction

Under standard conditions, which of
the following is the net reaction that
occurs in the cell?
Cd|Cd 2+ || Cu 2+ |Cu
a. Cu 2+ + Cd → Cu + Cd 2+
b. Cu + Cd → Cu 2+ + Cd 2+
c. Cu 2+ + Cd 2+ → Cu + Cd
d. Cu + Cd 2+ → Cd + Cu 2+

Galvanic Cell
• The reaction always runs
spontaneously in the direction that
produced a positive cell potential.
• Four things for a complete description.
1) Cell Potential
2) Direction of flow
3) Designation of anode and cathode
4) Nature of all the componentselectrodes
and ions

Practice
• Completely describe the galvanic cell
based on the following half-reactions
under standard conditions.
• MnO 4 - + 8 H + +5e - Mn +2 + 4H 2 O
Eº=1.51 V
• Fe +3 +3e - Fe(s)
Eº=0.036V

Batteries are Galvanic Cells
• Car batteries are lead storage batteries.
• Pb +PbO 2 +H 2 SO 4 PbSO 4 (s) +H 2 O

Batteries are Galvanic Cells
• Dry Cell
Zn + NH 4 + +MnO 2
Zn +2 + NH 3 + H 2 O + Mn 2 O 3

Batteries are Galvanic Cells
• Alkaline
Zn +MnO 2 ZnO+ Mn 2 O 3 (in base)

Batteries are Galvanic Cells
• NiCad
• NiO 2 + Cd + 2H 2 O Cd(OH) 2 +Ni(OH) 2

Corrosion
• Rusting - spontaneous oxidation.
• Most structural metals have reduction
potentials that are less positive than O 2 .
• Fe Fe +2 +2e -
• O 2 + 2H 2 O + 4e - 4OH -
• Fe +2 + O 2 + H 2 O Fe 2 O 3 + H +
• Reactions happens in two places.
Eº= 0.44 V
Eº= 0.40 V

Salt speeds up process by increasing
conductivity
Water
Fe 2+
Rust
e -
Iron Dissolves- O
Fe Fe +2 2 + 2H 2 O +4e - 4OH -
Fe 2+ + O 2 + 2H 2 O Fe 2 O 3 + 8 H +

Preventing Corrosion
• Coating to keep out air and water.
• Galvanizing - Putting on a zinc coat
• Has a lower reduction potential, so it is
more easily oxidized.
• Alloying with metals that form oxide
coats.
• Cathodic Protection - Attaching large
pieces of an active metal like magnesium
that get oxidized instead.

Electrolysis
• Running a galvanic cell backwards.
• Put a voltage bigger than the potential
and reverse the direction of the redox
reaction.
• Used for electroplating.

1.10
e - e -
Zn
1.0 M
Zn +2
1.0 M
Cu +2
Cu
Anode
Cathode

e - A battery e -
>1.10V
Zn
1.0 M
Zn +2
Cathode
1.0 M
Cu +2
Anode
Cu

A way to remember
• An Ox – anode is where oxidation occurs
• Red Cat – Reduction occurs at cathode
• Galvanic cell- spontaneous- anode is
negative
• Electrolytic cell- voltage applied to make
anode positive