CHEM 5013
Applied Chemical Principles
Chapter Thirteen
Professor Bensley
Alfred State College
Chapter Objectives




Define oxidation and reduction.
Write and balance half-reactions for
simple redox processes.
Describe the differences between galvanic
and electrolytic cells.
Use standard reduction potentials to
calculate cell potentials under standard
conditions.
Chapter Objectives


Use standard reduction potentials to
predict the spontaneous direction of a
redox reaction.
Calculate the amount of metal plated, the
amount of current needed, or the time
required for an electrolysis process.
Chapter Objectives



Distinguish between primary and
secondary batteries.
Describe the chemistry of some common
battery types and explain why each type
of battery is suitable for a particular
application.
Describe at least two common techniques
for preventing corrosion.
Oxidation / Reduction

Oxidation/Reduction Reaction (Redox):

Oxidation:

Reduction:

Oxidation Numbers:
Rule # Applies To
Statement
1
Elements
Oxidation number of an element is ALWAYS zero (0).
2
Compounds
The sum of the oxidation numbers of the atoms in a compound
is ALWAYS zero (0).
3
Monatomic
Ions
Oxidation number of a monatomic ion is ALWAYS equal to the
charge on the ion.
4
Polyatomic Ions The sum of the oxidations numbers of the atoms in a
polyatomic ion equals the charge on the ion.
5
Oxygen
The oxidation number of oxygen is -2 when it is in a compound
or a polyatomic ion.
6
Hydrogen
The oxidation number of hydrogen is +1 when it is in a
compound or a polyatomic ion.
7
Halogens
The oxidation number of halogen atoms is -1 when found in a
compound or polyatomic ion unless it is combined with
oxygen.
Oxidation Numbers
Assign oxidation numbers to Carbon in each of the
following compounds:
1. Carbon Monoxide
2. Carbon Dioxide
3. C6H12O6
4. Sodium Bicarbonate
Redox Half-Reactions



What happens when copper wire is placed
in a silver nitrate solution?
The solution’s blue color is indicative of what in
solution?
What are the crystals forming on the surface of
the copper wire?
Reducing and Oxidizing Agents

Reducing Agent:

Oxidizing Agent:
Oxidation
0
+1
+2
0
Cu(s) + 2Ag+ (aq)  Cu2+ (aq) + 2Ag (s)
Reducing
Agent
Oxidizing
Agent
Reduction
Building a Galvanic Cell

Galvanic cell:
Galvanic Cell Terminology

Salt Bridge:

Electrodes:

Anode:

Cathode:
Galvanic Cell Terminology

Cell notation:
anode | anode electrolyte || cathode electrolyte | cathode

EMF (electromotive force) or cell potential:

E0 (cell) =
Cell Potentials
1.
2.
Calculate the cell potential for the previous Copper/Silver
cell.
Calculate the cell potential for the following galvanic cell.
Fe(s) | Fe2+ (aq) (1.0M) || Cu
2+
(aq) | Cu (s)
Batteries

Battery:

Primary Cells:

Secondary Cells:
Primary Cells



Dry Cell Battery
1.5 Volts
Has a finite life
even when not
used since acidic
NH4Cl corrodes
can
Primary Cells



Alkaline Dry
Cell
1.5 Volts
Longer life than
dry cell but
more expensive
Primary Cells



Lithium – Iodine
Battery
High resistance,
low current
Used in
pacemakers and
is very reliable
(10 yrs)
Secondary Cells



Nickel-Cadmium
(NiCad)
Used in
calculators,
power tools,
shavers, etc.
Rechargeable
and light.
Secondary Cells



Lead Storage
Cell
Car battery rechargeable
Single cell is 2V,
6 cells in a row
so overall is
approx. 12 V
Electrochemistry Applications





Fuel Cell
Continuous supply of
fuel
Anode-hydrogen gas,
Cathode-oxygen gas
VERY efficient
Storage and transport
of Hydrogen is
limitation.
2 H2 (g) + O2 (g)  2 H2O (l)
Electrochemistry Applications

Corrosion – rust – forms only in the
presence of O2 and H2O.
Electrochemistry Applications

Galvanizing:

Cathodic Protection:
Electrolysis

Electrolysis:

Passive electrolysis:

Active electrolysis:
Active Electrolysis and
Electroplating


Electroplating:
Electrochemical reactions involved in the
plating of silver

Anode: Ag(s) + 2 CN- (aq) 
 Ag(CN)-2 (aq) + e-

Cathode: Ag(CN)- (aq) + e- 

Ag(s)
+
2
CN
(aq)
2
Electroplating

Current:

The unit of current, the ampere (A), is defined
as one coulomb per second:
1A=1C/s
Charge = current  time
Q= I  t
Current and Charge



Faraday’s constant: F = 96,485 C/mol
Use charge, Faraday’s Constant, and #
moles of electrons to determine mass of
metal plated on object.
In a copper plating experiment in which copper
metal is deposited from copper(II) ion solution,
the system is run for 2.6 hours at a current of
12.0 A. What mass of copper is deposited?