H2
Electrochemistry
This chapter investigates the new batteries that have been developed over recent years
and looks at the oxidation-reduction reactions employed by these batteries. Galvanic
and electrolytic cells are compared and the standard electrode potentials calculated.
Oxidation-reduction
reactions


Oxidation
Loss of electrons
Reduction
Gain of electrons
Displacement
reactions




Redox reactions
Electron transfer reactions
Reaction in which one metal displaces the ions of another
metal from solution.
The metal loses an electron and is oxidised
The metal ion gains an electron and is reduced
The more reactive metal is the one which displaces the
less reactive metal from a solution of its ions
Oxidation state
Is equal to the charge on a monatomic ion
Galvanic cell


A device that makes a chemical reaction occur in such a
way that it generates electricity.
Also known as a voltaic cell
Electrodes
Conductors of a cell that are connected to the external circuit
and are the sites for the oxidation and reduction reactions.
Electrolyte
A substance that conducts electricity when in solution or
molten.
Salt bridge
Allows migration of ions to occur between the electrodes so
the electric circuit is complete.
Galvanic cell
Produces electricity when
 The reaction at one electrode produces electrons that flow
from the electrode to the external circuit.
 The electrons flow through the metallic conductor of the
external circuit to the other electrode
 The reaction at the other electrode consumes these
electrons
 Ions migrate through the solutions and connecting salt
bridge to maintain electrical neutrality
Cell diagram
Is a shorthand way of representing cells
Cu│Cu2+║ Ag+│Ag
Anode



Electrode at which oxidation occurs
Negative terminal in a galvanic cell
Positive terminal in an electrolytic cell
Cathode



Electrode at which reduction occurs
Positive terminal for a galvanic cell
Negative terminal in an electrolytic cell
Leclanché cell






Also known as the dry cell
Zinc anode
Carbon rod as the cathode
Ammonium chloride (NH4Cl) paste as the electrolyte
Zn│Zn2+║ NH4Cl│MnO2, C
Oxidation reaction
Zn
Zn2+ + 2eReduction reaction
NH4+ + MnO2 + H2O + eMn(OH)3 + NH3

Other galvanic cells






Electrolysis
Process in which an electric current is used to bring about a
chemical reaction that does not occur spontaneously.
The electrode reactions depend on
 The nature of the ions present
 The concentration of the ions
 The nature of the electrodes
Electroplating
Electrolytic process of depositing a thin film of one metal on
the surface of an object made of another (generally cheaper)
metal
Lead-acid battery







Alkaline cell
Mercury cell
Lithium cell
Vanadium redox cell
Liquid junction voltaic cell
Silver oxide cell
Also known as the lead accumulator battery
Lead anode
Leadcathode
Sulfuric acid solution as the electrolyte
Pb│PbSO4║ PbO│Pb
Oxidation reaction
Pb + SO42PbSO4(s) + 2eReduction reaction
PbO2 + 4H+ + SO42- + ePbSO4(s) + 2H2O
Electromotive force



Standard electrode
potential



Calculating Eo for a
cell





EMF of a galvanic cell
The potential difference (voltage) across the electrodes of
the cell when a negligibly small current is being drawn.
The maximum voltage that the cell can deliver.
Eo
Is the potential of an electrode in the standard state
relative to the standard hydrogen electrode
Also assigned to the standard reduction half reactions
associated with the electrodes
Eo(cell)
Eo(red) + Eo(oxid)
When a standard reduction half equation is reversed the
sign of the Eo value is changed
Doubling the half equation does not change the Eo value
If Eo is positive the reaction occurs as written
If Eo is negative the reverse reaction occurs.