ELECTROCHEMISTRY

Prepared by: Mr.P.L.Meena

WHAT IS ELECTROCHEMISTRY ?

Electrochemistry is the scientific study of the chemical species and reactions that take place at the interface between an electron conductor and an ion conductor in which an electron transfer occurs between the electrode and electroyte in a solution.

ELECTROCHEMICAL CELLS

An electrochemical cell is a device capable of either deriving electrical energy from chemical reactions, or facilitating chemical reactions through the introduction of electrical energy. A common example of an electrochemical cell is a standard 1.5-volt "battery".

Daniel cell

Representation of an Electrochemical Cell Anode; Anode electrolyte (C1) || Cathode electrolyte (C2);

Cathode Oxidation half cell Salt Bridge Reduction half cell

Zn; Zn2+ (1M) || Cu2+ (1M) ; Cu

In a electrochemical cell

The electrode where electrons are released or where oxidation occurs , is known as anode.

Anode is also called negative pole.The electrode where electrons are accepted

or where reduction occurs is known as cathode.

Cathode is also called positive pole.The electrons flow from anode to cathode

while flow of current is in the opposite directions.

GALVANIC CELLSA Galvanic cell, or Voltaic

cell, named after Luigi Galvani, or Alessandro Volta respectively, is an electrochemical cell that derives electrical energy from chemical reactions taking place within the cell. It generally consists of two different metals connected by a salt bridge, or individual half-cells separated by a porous membrane.

Galvanic Cells

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The difference in electrical potential between the anode and cathode is called:• cell voltage• electromotive force (emf)• cell potential

Cell Diagram

Zn (s) + Cu2+ (aq) Cu (s) + Zn2+ (aq)

[Cu2+] = 1 M & [Zn2+] = 1 M

Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s)

anode cathode

STANDARD ELECTRODE POTENTIALS

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Zn (s) | Zn2+ (1 M) || H+ (1 M) | H2 (1 atm) | Pt (s)

2e- + 2H+ (1 M) H2 (1 atm)

Zn (s) Zn2+ (1 M) + 2e-Anode (oxidation):

Cathode (reduction):

Zn (s) + 2H+ (1 M) Zn2+ + H2 (1 atm)

Electrode potentialThe potential difference between electrode and the

electrolyte

Cell potentialthe potential difference between the two

electrodes of a galvanic cells.

EMFThe potential difference between the two

electrodes when no current flowing through the cell,

NERNST EQUATIONIt shows the relationship between the

electrode potential and concentration of the solution.

Example.

EQUILIBRIUM CONSTANT FROM NERNST EQUATION

or

ELECTROCHEMICAL CELL AND GIBBS FREE ENERGY

Conductance of electrolytic solutionResistance-the obstruction to the flow of

current.R α l/AR=ρ L/ASI unit -Ω(ohm)Resistivity- the resistance of a conductor of

one meter length and one meter square area of cross section

Ρ=RA/lIts SI unit is Ω m(ohm meter)

Conductance –the reciprocal of resistance G=1/R=A/ρlIts SI unit siemens(S) or Ω-1 (ohm-1 )Conductivity (k)-the inverse of resistivity k=1/ρ=l/RAIts SI unit is siemens/meter(S/m)Molar conductivity-the conductivity of the

solution containing one mole of electrolyte and kept between the two electrode with unit area of cross section and distance of unit length.

Λm=k(S cm-1 ) x 1000(cm3 L-1 )/C(Mol/L) SI unit is S cm2 mol-1 limiting molar conductivity-the molar

conductivity at infinite dilution or zero concentration.

variation of conductivity and molar conductivity with concentration-

For Weak electrolyte-molar conductivity increase steeply on dilution due to number of ions as well as mobility of ions increase.

For strong electrolyte-molar conductivity slowly with dilution due to increase in movements of ions in dilution.

In case of weak electrolyte the value of limiting molar conductivity can not be obtained by

Extrapolation this problem was solved with help of kohlraushs law.

Kohlarahs law-it state that the limiting molar conductivity of an electrolyte is the sum of limiting molar conductivities of cation and anion.

Λ∞m = v+ λ∞

+ + v- λ∞

General Types of Batteries

Primary Cell Non-rechargeable

Provides electricity until it dies (i.e. achieves equlibrium)

Disposable as redox couple is non-reversible

e.g. Zinc/Manganese battery (Dry cell)

Secondary Cell Rechargeable

Provides electricity until it goes flat

Connect to external power source to reverse redox reactions

e.g. Pb/PbSO4 battery

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Principle Components of a Battery

Anode Current CollectorCathode Current Collector

Anode active mass

Cathode active mass

Separator

Container

Terminals

Electrolyte

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Lead / Acid Battery

Pb(s) + SO42-(aq) PbSO4(s) + 2e-

PbO2(s) + 4H+(aq) + SO42-(aq) + 2e- 2H2O(l) + PbSO4(s)

PbO2(s) + Pb(s) + 4H+(aq) + 2SO42-(aq) 2PbSO4(s) + 2H2O(l)

Electrolyte : H2SO4

Current collectors : Both Pb

Cathode

Anode

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