Electrochemistry “It is the study of the interchange of chemical and electrical energy”

Electrochemistry The generation of an electric current from a spontaneous chemical reaction The use of current to produce chemical change Applications of Electrochemistry in our daily life Batteries (car batteries – calculators-digital watches) Corrosion of iron Preparation of some industrial materials eg. Al, Cl 2, NaOH, …… In analytical chemistry e.g. the analysis of chemicals in blood to determine the development of a certain diseases

Oxidation Reduction Reactions (REDOX) Oxidation : Loss of electrons Reduction : Gain of electrons M M+M+ X-X- X e Reduced Gains electrons Oxidizing agent Oxidized Loses electrons Reducing agent

Example Mg →Mg e - (Oxidation Reaction) O 2 + 4e - →2O -- (Reduction Reaction) When these half equations are paired and electrons balanced (i) (Mg → Mg e- ) x 2 loses electrons – oxidized – reducing agent (ii) O2 + 4e-→2O -- gains electrons – reduced – oxidizing agent Adding (i) and (ii) 2Mg + O 2 +4e→2Mg ++ +4e + 2O -- 2Mg + O 2 → 2 MgO Redox Reaction A Redox reaction is a reaction in which electrons are transferred from a reducing to an oxidizing agent

When reducing and oxidizing agents are present in the same solution, electrons are directly transferred. If we separate the oxidizing agent from the reducing agent transfer of electrons through a wire current production working of a motor useful work But current flows for an instant and then stops because of charge buildup

To solve the problem the solution must be connected either by a salt bridge or a porous disk (allows flow of ions without mixing of solutions) Now electrons flow through the wire from reducing to oxidizing agent and ions flow from one compartment to the other to keep the net charge zero

Galvanic Cell A device in which chemical energy is changed into electrical energy

An Example of a Galvanic Cell Anode :Zn  Zn e - Cathode :Cu e -  Cu The voltmeter measures the cell potential or electromotive force (emf) of the cell) The unit is the volt Cell reaction :Zn + Cu ++  Zn ++ + Cu

What is Cell Potential? It is the “pull” or “driving force” on the electrons from reducing agent to oxidizing agent The reaction in a galvanic cell is always an oxidation- reduction reaction that can be broken into 2 half reactions. It would be convenient to assign a potential to each ½ reaction so that we can obtain the cell potential by summing the ½ cell potentials. E cell = E (anode, oxid.) + E (cathode, red.) E cell = E (anode, oxid.) – E (cathode, red) But although we can measure the total potential of a cell there is no way to measure the potentials of the individual electrode processes

Standard Hydrogen Electrode The potential of cathode reaction 2H + +2e - →H 2 = 0 v ∴ the potential of the anode reaction Zn→Zn ++ +2e - =0.76 v

Example Conside a galvanic cell based on the reaction: Fe +++ (aq)+Cu(s)→Cu ++ (aq) + Fe ++ (aq) What are the 2 half reactions? Give the balanced cell reaction and calculate E° of the cell. 2 [ Fe e→ Fe ++ ] E° cathode =0.77 Cu→Cu ++ +2e E° anode = _______________________________ Cu + 2Fe +++ → Cu Fe + E° cell =0.43 N.B.: The E° cell (cell potential) is always +ve for a galvanic cell E° cell = E° cathode - E° anode

Example Describe completely the galvanic cell based on the following ½ reactions under standard conditions: Ag + +e -  Ag E° cell =0.8 v Fe e -  Fe ++ E° cell = 0.77 v Since E° cell must be +ve, thus the ½ reactions are: Cathode (red.) Ag + + e - → Ag E°=0.8 v Anode (oxid) Fe ++ →Fe e - E°= v _____________________________________________ Cell reaction Ag + + Fe ++ → Ag + Fe +++ E° cell =0.03 v

Dependence of Cell Potential on Concentration Nernest Equation Cell potential at conc. C Standard cell potential No. of transferred electrons Ionic concentration (equilibrium constant)

Example If E° cell is 0.48 v for the galvanic cell based on the reaction 2Al(s) + 3 Mn ++ (aq)  2Al +++ (aq) + 3 Mn(s) What is the cell potential if [Mn ++ ] = 0.5 M and [Al +++ ]=1.5 M

Concentration Cells Since cell potential depends on concentration, we can construct galvanic cells where both compartments contain same component but at different concentrations Batteries A battery is a galvanic cell or a group of cells connected in series, where the potential of the individual cells add to give the total battery potential. Eg. Lead storage battery (in automobiles) Dry cell Mercury battery

Chemical Impact Soon you may reach for a compact disc in a record store, and as you touch it, the package will start playing one of the songs on the disc. Or you may stop to look at a product because the package begins to glow as you pass it in the stores. These effects could happen soon due to the invention of a flexible, super thin battery that can actually be printed on the package. The battery consists of 5 thin layers of zinc (anode) and manganese diioxide (cathode) and is only 0.5 mm thick. The battery can be printed onto paper with a regular printing press. This battery intends to bring light, sound and other special effects to packaging to entice potential customers. Within a year or two, you might see talking, singing or glowing packages on the shelves. Printed Batteries