Summary Galvanic and Electrolytic Cells

GALVANIC AND ELECTROLYTIC CELLS


GALVANIC CELLS
 Convert chemical energy into electrical energy
 Self-sustaining electrode reactions



HALF CELLS
A half-cell consists of an electrode and an electrolyte
Electrolyte – a substance that can conduct electricity by forming free ions
when molten or in solution
Anode – the electrode where oxidation takes place
Cathode – the electrode where reduction takes place


Gaseous half-cells:
The hydrogen half cell
 H2 gas bubbled over Pt electrode through an electrolyte containing H +
ions
The chlorine half cell
 Cl2 gas bubbled over Pt electrode through an electrolyte containing Cl-
ions
Electrons flow from one half cell to the other
When a voltmeter is attached it will measure a potential difference which will
be the emf of the galvanic cell
The position of the substance of the electrodes must be compared using the
redox table

,  Whichever substance is higher on the right of the table is a stronger
reducing agent i.e. it will be oxidised
 Anode  decrease in mass cathode  increase in mass



FUNCTIONS OF A SALT BRIDGE
1) Complete the circuit
2) It maintains half-cell neutrality
Half-cell neutrality:
When cations enter or leave solution they create an excess of either +/- charge
which prevents more cations from entering or leaving solution. The salt bridge
provides either cations or anions in order to balance out the excess charges in
the solutions.



STANDARD CONDITIONS FOR A GALVANIC CELL
 Concentration of electrolyte = 1 mol.dm-3
 Temperature = 25oC
 Gas electrode pressure = 1 atm (100 kPa)
In a galvanic cell electrons flow from -  +



STANDARD CELL NOTATION
Anode on left
Cathode on right