ElectrochemiElectrochemistrystry

“It is the study of the interchange of chemical and

electrical energy”

Electrochemistry

The generation of an electric currentfrom 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, Cl2, NaOH, ……

•In analytical chemistry e.g. the analysis of chemicals in blood to determine

the development of a certain diseases

Oxidation Reduction Oxidation Reduction Reactions (REDOX)Reactions (REDOX) Oxidation : Loss of electrons Reduction : Gain of electrons

M

M+

X-

X

eReduced

Gains electrons

Oxidizing agent

OxidizedLoses electronsReducing agent

ExampleExample

Mg →Mg++ + 2e- (Oxidation Reaction)

O2 + 4e-→2O-- (Reduction Reaction)

• When these half equations are paired and electrons balanced

(i) (Mg → Mg+++ 2e- ) x 2 loses electrons – oxidized – reducing agent

(ii) O2 + 4e-→2O-- gains electrons – reduced – oxidizing agent

• Adding (i) and (ii)

2Mg + O2+4e→2Mg++ +4e + 2O--

2Mg + O2 → 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 CellGalvanic CellA device in which chemical energy is changed into electrical energy

An Example of a Galvanic CellAn Example of a Galvanic Cell

Anode :Zn Zn++ + 2e-

Cathode :Cu++ + 2e- Cu

The voltmeter measures the cellpotential 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.

Ecell = E(anode, oxid.)+ E(cathode, red.)

Ecell = 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-→H2 = 0 v

∴ the potential of the anode reaction Zn→Zn+++2e-=0.76 v

++ZnZn2H+Hcell E+E=E→→

++ZnZncell E=E

→

ExampleExample

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= -0.34_______________________________Cu + 2Fe+++ → Cu++ + 2Fe+ E°cell=0.43

N.B.: The E°cell (cell potential) is always +ve for a galvanic cell

E°cell= E°cathode - E°anode

ExampleExample

Describe completely the galvanic cell based on the following ½ reactions under standard conditions:

Ag++e- Ag E°cell =0.8 vFe+++ + 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°= -0.77 v_____________________________________________Cell reaction Ag+ + Fe++ → Ag + Fe+++ E°cell=0.03 v

Dependence of Cell Potential Dependence of Cell Potential on Concentrationon Concentration

Nernest Equation

)Qlog(n

0591.0 -E=E

Cell potentialat conc. C

Standard cell potential No. of

transferred electrons

Ionic concentration(equilibrium constant)

ExampleExample

If E°cell is 0.48 v for the galvanic cell based on the reaction2Al(s) + 3 Mn++(aq) 2Al+++(aq) + 3 Mn(s)

What is the cell potential if [Mn++] = 0.5 M and [Al+++]=1.5 M

v47.0=]5.0[

]5.1[log

60591.0

-48.0=E

]Mn[

]AL[log

60591.0

-48.0=E

)Qlog(n

0591.0 -E=E

3

2

cell

3++

2+++

cell

cell

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