Chem 150 Handout Electrochemistry
Voltaic Cells
In spontaneous oxidation-reduction (redox) reactions, electrons are transferred and
energy is released.
We can use that energy to do work if we make the electrons flow through an external
device. We call such a setup a voltaic cell. A typical cell looks like this:
Picture from Brown LeMay
Each compartment consisting of an electrode dipped in a salt solution is called a half cell. Two half cells
that are in electrical contact (both, the electrodes and the solutions need to be in electrical contact for
even one electron to flow from the anode to the cathode.
A salt bridge or a porous membrane is used to electrically connect the solutions in the two half cells
while a wire connects the two electrodes.
The oxidation occurs at the anode.
The reduction occurs at the cathode.
Water only spontaneously flows one way in a waterfall.
Likewise, electrons only spontaneously flow one way in a redox reaction—from higher to lower potential
energy.
Electromotive Force (emf): The potential difference between the anode and cathode in a cell is called
the electromotive force (emf).It is measured in volts (V) and is also called the cell potential and
designated Ecell.
Standard Cell Potentials: Because cell potential is based on the potential energy per unit of charge, it is
an intensive property and thus independent of the amount of substance
The cell potential at standard conditions can be found through this equation where the standard
potentials of each half-cell is measures against the standard hydrogen electrode:
E°cell = E°reduction(cathode) - E°reduction(anode)
Their values are referenced to a standard hydrogen electrode (SHE).
By definition, the reduction potential for hydrogen is 0 V.
2 H+ (aq, 1M) + 2 e−  H2 (g, 1 atm)
Picture, Brown Le May
Selected Standard Reduction Half-Potentials
Cell potentials under non standard conditions, Nernst Equation
Concentration cell
Lead acid battery (car battery)
Dry Cell (flash light battery)
The hydrogen fuel cell
Corrosion of iron (left) and corrosion of zinc on zinc plated steel (right)
Electrolysis
Voltaic cells are based on a spontaneous redox reaction but it is possible to use electricity to cause nonspontaneous redox reactions to occur. For example the electrolysis of molten sodium chloride to
produce sodium metal and chlorine gas. For example the electrolysis of molten sodium chloride to
produce sodium metal and chlorine gas.
Source Brown Le May
Class exercise: Can you think of a problem that could occur if a cell was set up in the way shown above?
What would happen if the electrolysis was carried out in water?
Electroplating is another instance where electrolysis is used to deposit a thin layer of one metal on
another to improve beauty or resistance to corrosion, for example the nickel-plating of steel.
History of Aluminum
At one time, more expensive than gold (in 1855, Al cost $100,000 per lb).
1885 – Washington Monument topped with world’s largest piece of Al (3 kg)
1886 – Price drops to $5 / lb.
Today, used in aircraft, beverage cans etc (cost $1 / lb.)
Production of Aluminum Hall-Héroult Process
Electrolysis at carbon electrodes (4.5 V; 150,000 A !!!) 15.7 kWh electricity needed per kg Al
In electrolysis reaction, cryolite (Na3AlF6) is used as a “solvent” and mixed with Al2O3.The melting point
of mixture is ~ 1000 °C (pure Al2O3 melts at 2054 °C)
Cathode / Reduction
Al3+ + 3 e-  Al (l)
Anode / Reduction
C (s) + 2 O2- (l)  CO2 (g) + 4 e- / times 3 to balance electron count
Net reaction:
4Al3+ (l) + 3C (s) + 6O2- (l)  4Al (l) + 3CO2 (g)
/ times 4
Large quantities of aluminum metal are produced each year at an enormous cost of electricity.
Recycling aluminum uses less than 10m% of the energy required to generate aluminum from raw
resources.