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ELECTROCHEMISTRY

ELECTROCHEMISTRY. Prepared by: Mr.P.L.Meena. What is electrochemistry ?.

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ELECTROCHEMISTRY

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  1. ELECTROCHEMISTRY Prepared by: Mr.P.L.Meena

  2. What is electrochemistry ? • Electrochemistry is the scientificstudy 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.

  3. 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

  4. 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

  5. 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.

  6. Galvanic cells • A 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.

  7. Zn (s) + Cu2+(aq) Cu (s) + Zn2+(aq) Galvanic Cells • The difference in electrical potential between the anode and cathode is called: • cell voltage • electromotive force (emf) • cell potential Cell Diagram [Cu2+] = 1 M & [Zn2+] = 1 M Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s) anode cathode 19.2

  8. 2e- + 2H+ (1 M) H2(1 atm) Zn (s) Zn2+ (1 M) + 2e- Zn (s) + 2H+ (1 M) Zn2+ + H2 (1 atm) Standard Electrode Potentials Zn (s) | Zn2+ (1 M) || H+ (1 M) | H2 (1 atm) | Pt (s) Anode (oxidation): Cathode (reduction): 19.3

  9. Electrode potential • The potential difference between electrode and the electrolyte Cell potential • the potential difference between the two electrodes of a galvanic cells. EMF • The potential difference between the two electrodes when no current flowing through the cell,

  10. Nernst equation • It shows the relationship between the electrode potential and concentration of the solution. Example.

  11. Equilibrium constant from nernst equation or

  12. Electrochemical cell and gibbs free energy

  13. Conductance of electrolytic solution • Resistance-the obstruction to the flow of current. • R α l/A • R=ρ L/A • SI unit -Ω(ohm) • Resistivity- the resistance of a conductor of one meter length and one meter square area of cross section • Ρ=RA/l • Its SI unit is Ω m(ohm meter)

  14. Conductance –the reciprocal of resistance • G=1/R=A/ρl • Its SI unit siemens(S) or Ω-1 (ohm-1 ) • Conductivity (k)-the inverse of resistivity • k=1/ρ=l/RA • Its 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.

  15. 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-

  16. 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.

  17. 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∞

  18. 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

  19. Principle Components of a Battery Terminals Anode Current Collector Cathode Current Collector Anode active mass Cathode active mass Electrolyte Container Separator

  20. Lead / Acid Battery Anode Pb(s) + SO42-(aq) PbSO4(s)+ 2e- PbO2(s) + 4H+(aq) + SO42-(aq)+ 2e- 2H2O(l) + PbSO4(s) Cathode PbO2(s) + Pb(s) + 4H+(aq) + 2SO42-(aq) 2PbSO4(s)+ 2H2O(l) Electrolyte : H2SO4 Current collectors : Both Pb

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