Chapter 20
Electrochemistry
20.1 Introduction to Electrochemistry
Electrochemistry
• The branch of chemistry that deals with
electricity-related applications of oxidationreduction reactions.
• Electrochemical Cells:
A system of electrodes and electrolytes in which
either chemical reactions produce energy or an
electrical current produces chemical change
Components of Electrochemical Cells
Cu Electrode
Cathodewhere
reduction
takes place
Conducting
Wire
Electrolyte
Sol’n
ZnSO4
Electrode: conductor used
to establish electrical
contact with a nonmetallic
part of the circuit.
Electrolyte
Sol’n CuSO4
Zn Electrode
Anode- where
oxidation takes
place
Half-Cell: a
single
electrode
immersed in a
solution of its
ions
Cu Electrode
Cathode- written
as
Cu+2/Cu
Overall Cell
Written as:
anode | cathode
Zn | Cu
Zn Electrode
Anode- written as
Zn+2/Zn
Half-Cell: a
single electrode
immersed in a
solution of its
ions
Chapter 20
Electrochemistry
20.2 Voltaic Cells
Electrochemistry
Porous barrier which prevents
the spontaneous mixing of the
aqueous solutions in each
compartment, but allows the
movement of ions in both
directions to maintain electrical
neutrality
Voltaic /
Galvanic
Cell
• A chemical rxn that results in a
voltage due to a transfer of electrons
Rxns that
produce voltage
spontaneously
Batteries
Zn → Zn+2 + 2e-
2MnO2 + H2O + 2e- → Mn2O3 + 2OH -
• Two or more dry
voltaic cells
• Zinc-Carbon Battery
Batteries
Zn + 2OH → Zn(OH)2 +
-
2e-
2MnO2 + H2O + 2e- → Mn2O3 + 2OH-
• Alkaline Battery- no
carbon rod, smaller
Batteries• Mercury Battery- no
Zn + 2OH - → Zn(OH)2 + 2e-
HgO + H2O + 2e- → Hg + 2OH -
carbon rod, smallest
Fuel Cells
Cathode: O2 + 2H2O + 4e- → 4OH –
Anode: 2H2 + 4OH – → 4e- + 4H2O
Net: 2H2 + O2 → 2H2O
• A voltaic cell where
reactants are constantly
supplied and products are
removed.
Rxns that turn
chemical energy
into electrical
energy
12
Corrosion
Formation of Rust:
4Fe (s) + 3O2 (g) + xH2O → 2Fe2O3∙xH2O
Anode: Fe (s) → Fe+2 (aq) + 2eCathode: O2 (g) + 2H2O (l) + 4e- → 4OH –
Prevention of Corrosion
Galvanizing
Process by which iron or any
metal is coated with zinc.
Cathodic Protection
Since zinc is
more easily
oxidized, it is a
sacrificial
anode.
Electrode Potentials
• Reduction Potential: the tendency for the halfreaction to occur as a reduction half-reaction in an
electrochemical cell.
• Electrode Potential: the difference in potential
between an electrode and its solution
• Potential Difference (Voltage): a measure of the
energy required to move a certain electric charge
between the electrodes, measured in volts.
• Standard Electrode Potential (E°): a half-cell
measured relative to a potential of zero for the standard
hydrogen electrode (SHE)
Standard Electrode Potential, E°
• Positive E° means hydrogen is more willing to give up
its electron, so positive reduction potentials are favored.
Naturally occurring rxns have a positive value.
E° cell = E° cathode - E° anode
• Negative E° means the metal electrode is more willing
to give up its electron, this is not favored. These rxns
prefer oxidation over reduction.
Standard Electrode Potential, E°
• When a half-cell is multiplied by a constant (for
balancing) the E° value is NOT multiplied!
• When a rxn is reversed (flipped) the sign of the E° value
switches.
• In a voltaic cell, the half-rxn with the more negative
standard electrode potential is the anode, where
oxidation occurs.
Cell Potential
• The potential voltage a rxn can produce.
Cu2+ + 2e-  Cu
Eo = .34 V
Ag+ + e-  Ag
Eo = .80V
Because this is a spontaneous process:
(Ag+ + e-  Ag) x 2
Eo = .80V
Cu  Cu2+ + 2e-
Eo = -.34 V
Cu + 2Ag+ Cu2+ + 2Ag
Eo = .46 V
Reduction
potentials
Since both rxns are
reduction, one
must be oxidation,
flip it, positive
voltage must result
from spontaneous
rxns
Cell Potential
• The potential voltage a rxn can produce.
Na+ + e-  Na
Eo = -2.71 V
Cl2 + 2e-  2Cl-
Eo = 1.36 V
Because this is nonspontaneous process:
(Na+ + e-  Na) x 2
Eo = -2.71 V
2Cl-  Cl2 + 2e-
Eo = -1.36 V
2Na+ + 2Cl-  2Na + Cl2
Eo = -4.07 V
Nonspontaneous,
must end in
negative voltage.
Flip one to become
oxidation.
** Fuel Cell!
Chapter 20
Electrochemistry
20.3 Electrolytic Cells
Electrochemistry
• When electric voltage is used
to produce a redox reaction, it
is called electrolysis
Electrolytic Cell
Rxns that require an
energy source to react
Batteries
Discharge Cycle Rxn:
Pb + PbO2 + 2H2SO4 → 2PbSO4 + 2H2O
• Car Batteryrechargeable b/c
the alternator
reverses the ½
rxns and
regenerates the
reactants.
Electroplating
• An electrolytic process in which a
metal ion is reduced and a solid
metal is deposited on a surface
• Typically, an inactive metal is
able to be ionized and then
deposited on the surface of a more
active metal to prevent corrosion.
Anode
Silver ions are reduced at the cathode:
Ag+ + 1e- → Ag
Silver atoms are oxidized at the anode:
Ag → Ag + + 1e-
Voltaic vs. Electrolytic
• If the positive battery terminal is attached to the cathode
of a voltaic cell, and the negative terminal is attached to
the anode, the flow of electrons will change directions.
• Electrolytic cells need the electrodes attached to a
battery, where voltaic is its own source of electrical
power.
Voltaic = spontaneous
chemical energy → electrical energy
Electrolytic = non-spontaneous
electrical energy → chemical energy
Electrolysis
Anode: 6H2O → O2 + 4e- + 4H3O+
Cathode: 4H2O + 4e- → 2H2 + 4OH –
Using a current to generate
a redox reaction which
otherwise would have a
negative cell potential.
i.e. electroplating &
rechargeable batteries.