1. electrochemical cells
Electrochemistry
Redox reactions
electrochemical cells
galvanic:
spontaneous chemical rxns to produce electrical energy
(DG = -nFE, negative)
electrolytic:
utilisation of energy (ex: applied V) to force a chemical rxn to take place (DG +)

2. Cells reduction at the cathode oxidation at the anode galvanic cells
half-rxns
galvanic cells
line notation
|interface between two phases.
||salt bridge
Cd(s) | CdCl2(aq, M) || AgNO3(aq, M) | Ag(s)

3. Standard reduction potentials
To predict the reactivity of oxidants or reductants we need to measure the
potential of each half-rxn.
 impossible!!....for every oxidation rxn we have a reduction reaction
Define a standard half-cell of potential = 0V
against which all other half-cell reduction potentials are measured. Each component in these standard cells having unit concentration.
By convention: Standard (or Normal) Hydrogen Electrode is used
Pt(s) | H2(g) | H+(a) || Ag+(a) | Ag(s)
|_______________|
NHE
H+(aq) + e-<=>H2(g)E0=0V

4. Electrochemical series
half-rxns
oxidant reductant E0 (V)
stronger oxidant
F2(g) + 2e- <=> 2F
Ce4+ + e- <=> Ce
Ag+ + e- <=> Ag(s)
Fe3+ + e- <=> Fe
O2 + 2H+ + 2e- <=> H2O
Cu2+ + 2e- <=> Cu(s)
2H+ + 2e- <=> H2(g)
Cd2+ + 2e- <=> Cd(s)
Zn2+ + 2e- <=> Zn(s)
K+ + e- <=> K(s)
Li+ +e- <=> Li(s)
stronger reducer

5. When Concentrations are not unity?
Nernst equation
for a half-rxn
aOx + ne- <=> bRed
E = E0 – RT/nF ln{[Red]b/[Ox]a}
R= gas constant T= temperature in Kelvin
n= no. of electrons in half-rxn F= Faraday constant
Converting ln to log10 (x 2,303) and at 25oC (298.15K)
E = E0 – /n log{[Red]b/[Ox]a}
The cell potential
E = E+ - E-
(E+ and E- calculated using the Nernst equation)

6. Applications of Galvanic Cells
Potentiometry and Ion Selective Electrodes
the measure of the cell potential to yield chemical information (conc., activity, charge ....)
A difference in the activity of an ion on either side of a selective membrane results in a thermodynamic potential difference being created across that membrane

7. The glass pH electrode

8. Corrosion Fe2+ +2e Fe E0=-0.44V 2H+ + 2e H2 E0=0V
2H2O + O2 =4e4OH- E0=1.23V
Iron is oxidized in water or humid conditions to give rust.
Inhibit this by coating with another material (Zn for example that forms a protective oxide on the iron), or by providing a sacrificial anode (b).

9. Batteries-providing electricity from chemistry
The Lead Acid Storage Battery was developed in the late 1800's and has remained the most common and durable of the battery technologies (in vehicles).

10. Lead-acid batteries
When the battery is used as a voltage supply, electrons flow from the Pb metal to the Pb(IV)oxide. The reactions aren't quite the reverse of the formation reactions, because now the sulfate ions in the solution begin to play a role. The two reactions are: PbO2 + 4H+ + 2e + SO4-2  PbSO4 + 2H2O Pb + SO4-2  PbSO4 + 2e The overall reaction if we combine the hydrogen ions and the sulfate: PbO2 + Pb + 2H2SO4  2 PbSO4 + 2 H2O Lead sulfate is fairly insoluble so that as soon as Pb(II) ions are formed by either reaction, the ions immediately precipitate as lead sulfate. The beauty is that this lead sulfate stays attached to the grids so that it is there for recharging of the battery.

11. Other batteries Primary battery-non rechargeable Longer shelf-life
Offer higher efficiencies compared to burning fuels

12. Images of batteries
Leclanche Alkaline Fuel Cell

13. Electrolysis Use Faraday’s Laws to evaluate the number of moles of
a substance oxidised or reduced by passage of charge
(current over a given period of time = I.t) through an electrode
Faraday: Q (charge) = nF
N=number of moles of electrons
F=constant of Coulomb/mole
Example (try it): What current is needed to deposit 0.500g of chromium from a solution containing Cr3+ over a one hour period (MW for Cr=52)? (Ans=0.77A)

14. Applications of Electrolytic Cells
Aluminium refining: The major ore of aluminium is bauxite, Al2O3.
Anhydrous Al2O3 melts at over 2000°C. This is too high to permit its use as a molten medium for electrolytic formation of free aluminium. The electrolytic process commercially used to produce aluminium is known as the Hall process, named after its inventor, Charles M. Hall. Al2O3 is dissolved in molten cryolite, Na3AlF6, which has a melting point of 1012oC and is an effective conductor of electric current. Graphite rods are employed as anodes and are consumed in the electrolysis process. The cell electrolytic reaction is:
2Al2O3 + 3C  4Al(l) + 3CO2(g)

15. Electrolysis of brine
Chlorine and sodium hydroxide are both manufactured by electrolysis of brine (aqueous sodium chloride) using inert electrodes. Chlorine is evolved at the anode, Cl- 1/2Cl2 + e
Hydrogen is evolved at the cathode: H+ + e 1/2H2
The removal of chloride ions and hydrogen ions leaves sodium ions and hydroxide ions in solution.
Chlorine is used to disinfect municipal water supplies and water in swimming pools. It is used to manufacture household bleaches and disinfectants. It is used to manufacture plastics (e.g. PVC), pesticides, anaesthetics, CFCs etc.
Sodium hydroxide is used in the manufacture of synthetic fibres, soaps and detergents.

16. Electroplating
In all aspects of our lives we are surrounded by products with electroplated surfaces. Whether we are looking at a silver-plated watch through gold-plated glasses, watching television, using the washing machine, getting into a car or boarding a plane: electroplating plays an important part in all of these situations. The objective is to prevent corrosion and wear,produce hardness and conductivity, and give products an attractive appearance.
The principle: thin metallic layers with specific properties are deposited on base materials including steel, brass, aluminium, plastic and die-cast parts.
Silver electroplating was the first large scale use of electrolysis for coating base metal objects with a higher value decorative finish.