John E. McMurry • Robert C. Fay
C H E M I S T R Y
Chapter 17
Electrochemistry
Galvanic Cells
2 CuSO4 (aq) + Zn(s) ------>ZnSO4 (aq) + Cu(s)
As this spontaneous reaction proceeds, the Zn(s) dissolves, Cu precipitates. This a
redox reaction:
Net ionic equation: Cu2+ (aq) + Zn (s) ----> Cu (s) + Zn2+ (aq)
This can be split into two half reactions:
Cu2+ (aq) + 2e-  Cu (s) reduction (gain of electrons)
Zn (s)  Zn2+ (aq) +2e- oxidation (loss of electrons
Electrochemistry: The area of chemistry concerned with the
interconversion of chemical and electrical energy
Galvanic (Voltaic) Cell: A spontaneous chemical reaction which
generates an electric current
Electrolytic Cell: An electric current which drives a nonspontaneous
reaction
Electrochemical Cell Components




Two conductors (anode and
cathode)
Electrolytes solution:
solution that each electrode
is emerse in it
External circuit: provide a
pathway for electron to
move from one electrode to
another
Salt Bridge: provide
neutrality
Zn(s) + Cu2+(aq)
Zn2+(aq) + Cu(s)
Galvanic Cells
•
Anode:
• The electrode where
oxidation occurs.
• The electrode where
electrons are produced.
• Is what anions migrate
toward.
• Has a negative sign.
Anode (-)
Cathode (+)
•Cathode:
•The electrode where
reduction occurs.
•The electrode where
electrons are
consumed.
•Is what cations
migrate toward.
•Has a positive sign
Galvanic Cells
•
Salt Bridge: a U-shaped tube that contains a gel permeated with a
solution of an inert electrolytes
• Maintains electrical neutrality by a flow of ions
• Anions flow through the salt bridge from the cathode to anode
compartment
• Cations migrate through salt bridge from the anode to cathode
compartment
Why do negative ions (anions) move toward
the negative electrode (anode)?
Shorthand Notation for Galvanic
Cells or Voltaic Cell
Salt bridge
Anode half-cell
Cathode half-cell
Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)
Electron flow
Phase boundary
Phase boundary
Shorthand Notation for Galvanic
Cells or Voltaic Cell
How do you write the line-notation for a galvanic cell
that has an aqueous or gaseous component instead of a
solid metal?
Cell involving gas
Additional vertical line due to
presence of addition phase
List the gas immediately adjacent
to the appropriate electrode
Detailed notation includes ion
concentrations and gas pressure
Pt(s)|Fe2+(aq), Fe3+(aq)||Ag+(aq)|Ag(s)
Example

Consider the reactions below
◦ Write the two half reactions
◦ Identify the oxidation and reduction half
◦ Identify the anode and cathode
◦ Give short hand notation for a galvanic cell that employs the
overall reaction
Fe (s) + Sn2+ (aq)  Fe2+ (aq) + Sn (s)
Example

Given the following shorthand notation, sketch out the
galvanic cell
Pb (s) | Pb2+ (aq) || Br2 (l) | Br- (aq) | Pt (s)
Cell Potentials and Free-Energy
Changes for Cell Reactions
Electromotive Force (emf): The force or electrical potential that
pushes the negatively charged electrons away from the anode (
electrode) and pulls them toward the cathode (+ electrode).
It is also called the cell potential (E) or the cell voltage.
The standard hydrogen electrode (S.H.E.) has been chosen to be the
reference electrode.
2H+(aq,
1 M) +
H2(g, 1 atm)
2e
H2(g, 1 atm)
2H+(aq, 1 M) + 2e
E°ox = 0 V
E°red = 0 V
Standard Reduction Potentials

Eocell is the standard cell potential when both products and
reactants are at their standard states:
◦ Solutes at 1.0 M
◦ Gases at 1.0 atm
◦ Solids and liquids in pure form
◦ Temp = 25.0oC
E°cell = E°ox + E°red
Standard Reduction Potentials
Anode half-reaction:
H2(g)
Cu2+(aq) + 2e
Cathode half-reaction:
H2(g) + Cu2+(aq)
Overall cell reaction:
2H+(aq) + 2e
Cu(s)
2H+(aq) + Cu(s)
E°cell = E°ox + E°red
0.34 V = 0 V + E°red
A standard reduction potential can be defined:
Cu2+(aq) + 2e
Cu(s)
E° = 0.34 V
Standard Reduction Potentials

Spotaniety of the reaction can be determined by the positive
Eocell value

The cell reaction is spontaneous when the half reaction with the
more positive Eo value is cathode

Note: Eocell is an intensive property; the value is independent of
how much substance is used in the reaction
Ag+(aq) + e-  Ag(s)
2 Ag+(aq) + 2e-  2 Ag(s)
Eored = 0.80 V
Eored = 0.80V
Standard Reduction Potentials
Examples

Of the two standard reduction half reactions below, write the
net equation and determine which would be the anode and
which would be the cathode of a galvanic cell. Calculate Eocell
a.
b.
Cd2+(aq) + 2e-  Cd(s)
Eored = -0.40 V
Ag+(aq) + e-  Ag(s)
Eored = 0.80 V
Fe2+(aq) + 2e-  Fe(s)
Eored = -0.44 V
Al3+(aq) + 3e-  Al(s)
Eored = -1.66 V
Cell Potentials and Free-Energy
Changes for Cell Reactions
faraday or Faraday constant
The electric charge on 1 mol of electrons and is equal to 96,500 C/mol e
DG = nFE
Free-energy change
or
DG° = nFE°
Cell potential
Number of moles of electrons transferred in the
reaction
Cell Potentials and Free-Energy
Changes for Cell Reactions
The standard cell potential at 25 °C is 1.10 V for the reaction:
Zn(s) + Cu2+(aq)
Zn2+(aq) + Cu(s)
Calculate the standard free-energy change for this reaction at 25 °C.
Is the reaction spontanous at this condition?
Examples

Calculate the cell potential at standard state (Eocell) for the
following reaction. Then write the half reactions
I2(s) + 2 Br-(aq)  2I-(aq) + Br2(l)
ΔGo = 1.1 x 105J
Standard Cell Potentials and
Equilibrium constants
Nernst Equation: describe the relationship between Ecell and the concentration
of species involved in the cell reaction
0.0592 V
E = E° 
log Q
n
At Equilibrium E = 0
E° =
0.0592 V
n
log K
in volts, at 25 oC
Standard Cell Potentials and
Equilibrium Constants
Example

What is the value of Eo for a redox reaction involving the
transfer of 2 mol electrons if its equilibrium constant is
1.8 x 10-5?
Example

Calculate the concentration of cadmium ion in the galvanic cell below
Cd(s)|Cd2+(aq)(?M)||Ni2+(aq)(0.100M)|Ni(s)
Ecell = 0.30V
Electrolysis and Electrolytic Cells


Anode: where oxidation takes place
◦ Anions are oxidized at this electrode
◦ labeled positive to reflect anions attraction to anode
Cathode: where reduction takes places
◦ Cations are reduced at this electrode
◦ Labeled negative to reflect the cations attraction to cathode
Electrolysis and Electrolytic Cells
Electrolysis: The process of using an electric current to bring about
chemical change.
Electrolysis and Electrolytic Cells
•
Electrolysis: The process of using
an electric current to bring about
chemical change.
•
Process occurring in galvanic
cell and electrolytic cells are the
reverse of each other
•
In an electrolytic cell, two
inert electrodes are dipped into an
aqueous solution
Predicting the Products of
Electrolysis

– The cations (+) are attracted to the cathode (-) and the

anions (-) are attracted to the anode (+)

Electrolysis of molten salts – used for industrial isolation of the most
active elements (Na, Li, Mg, Al, …; F2, Cl2, Br2, …)

– The cation is reduced at the cathode

– The anion is oxidized at the anode

Example: Isolation of Na and Cl2 by electrolysis of molten NaCl
Na+(l) + e- → Na(l) (×2)
cathode, reduction
2Cl-(l) → Cl2(g) + 2e-
anode, oxidation
2Na+(l) + 2Cl-(l) → 2Na(l) + Cl2(g) net-ionic equation
Electrolysis of Molten Salts

Write the half-reactions for the electrolysis of the following
molten compounds
KCl(s)

MgO(s)

Electrolysis of mixed molten salts
– The cation with higher Eo value (the stronger oxidizing agent) is reduced
at the cathode
– The anion with lower Eo value (the stronger reducing agent) is oxidized at
the anode
Example: Predict the products of the electrolysis of a molten mixture of NaCl
and AlF3

→ Possible cathode half-reactions (reduction)
1) Reduction of Na+ and 2) Reduction of Al3+

Possible anode half-reactions (oxidation)
1) Oxidation of F- and 2) Oxidation of Cl-