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19.2 Galvanic Cells 19.3 Standard Reduction Potentials 19.4 Spontaneity of Redox Reactions PowerPoint PPT Presentation


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Chapter 19 Electrochemistry Semester 2/2012. Ref: http://www.mhhe.com/chemistry/chang. 19.2 Galvanic Cells 19.3 Standard Reduction Potentials 19.4 Spontaneity of Redox Reactions 19.5 The Effect of Concentration on Emf 19.8 Electrolysis. 19.2 Galvanic Cells. anode oxidation. cathode

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19.2 Galvanic Cells 19.3 Standard Reduction Potentials 19.4 Spontaneity of Redox Reactions

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19 2 galvanic cells 19 3 standard reduction potentials 19 4 spontaneity of redox reactions

Chapter 19

Electrochemistry

Semester 2/2012

Ref: http://www.mhhe.com/chemistry/chang

19.2 Galvanic Cells

19.3 Standard Reduction Potentials

19.4 Spontaneity of Redox Reactions

19.5 The Effect of Concentration on Emf

19.8 Electrolysis


19 2 galvanic cells 19 3 standard reduction potentials 19 4 spontaneity of redox reactions

19.2 Galvanic Cells

anode

oxidation

cathode

reduction

Spontaneous(natural)

redox reaction


19 2 galvanic cells 19 3 standard reduction potentials 19 4 spontaneity of redox reactions

  • Cell = half-cell + half – cell

  • OxidationReduction

  • AnodeCathode

  • In Galvanic cell…

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

  • Zn is oxidized to Zn2+ ion

  • Zn electrode is Anode (Reducing Agent)

    Cu2+ is reduced to Cu

     Cu electrode is Cathode (Oxidizing Agent)


19 2 galvanic cells 19 3 standard reduction potentials 19 4 spontaneity of redox reactions

Cell Equation 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

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

anode

cathode


19 2 galvanic cells 19 3 standard reduction potentials 19 4 spontaneity of redox reactions

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 2 galvanic cells 19 3 standard reduction potentials 19 4 spontaneity of redox reactions

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

19.3 Standard Reduction Potentials

Standard reduction potential (E0) is the voltage associated with a reduction reaction at an electrode when all solutes are 1 M and all gases are at 1 atm.

Reduction Reaction

E0= 0 V

Standard hydrogen electrode (SHE)


19 2 galvanic cells 19 3 standard reduction potentials 19 4 spontaneity of redox reactions

E0 = 0.76 V

E0 = Ecathode - Eanode

E0 = EH /H - EZn /Zn

cell

cell

cell

Standard emf (E0 )

cell

0

0

0

0

2+

+

2

0.76 V = 0 - EZn /Zn

0

2+

EZn /Zn = -0.76 V

0

2+

Zn2+ (1 M) + 2e- ZnE0 = -0.76 V

Standard Electrode Potentials

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


19 2 galvanic cells 19 3 standard reduction potentials 19 4 spontaneity of redox reactions

0

0

0

Ecell = ECu /Cu–EH /H

2+

+

2

0.34 = ECu /Cu - 0

0

2+

0

ECu /Cu = 0.34 V

2+

E0 = Ecathode - Eanode

E0 = 0.34 V

cell

cell

0

0

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

2e- + Cu2+ (1 M) Cu (s)

H2 (1 atm) + Cu2+ (1 M) Cu (s) + 2H+ (1 M)

Standard Electrode Potentials

Pt (s) | H2 (1 atm) | H+ (1 M) || Cu2+ (1 M) | Cu (s)

Anode (oxidation):

Cathode (reduction):


19 2 galvanic cells 19 3 standard reduction potentials 19 4 spontaneity of redox reactions

  • Note:

  • The more positive E0 the greater the tendency for the substance to be reduced

  • The half-cell reactions are reversible

  • The sign of E0changes when the reaction is reversed

  • Changing the stoichiometric coefficients of a half-cell reaction does not change the value of E0


19 2 galvanic cells 19 3 standard reduction potentials 19 4 spontaneity of redox reactions

DG0 = -nFEcell

0

0

0

0

0

= -nFEcell

Ecell

Ecell

Ecell

F = 96,500

J

RT

V • mol

ln K

nF

(8.314 J/K•mol)(298 K)

ln K

=

0.0592 V

0.0257 V

log K

ln K

n (96,500 J/V•mol)

=

n

n

=

=

19.4 Spontaneity of Redox Reactions

DG = -nFEcell

n = number of moles of electrons in reaction

= 96,500 C/mol

DG0 = -RT ln K

E0cell > 0 spontaneous reaction


19 2 galvanic cells 19 3 standard reduction potentials 19 4 spontaneity of redox reactions

Spontaneity of Redox Reactions


19 2 galvanic cells 19 3 standard reduction potentials 19 4 spontaneity of redox reactions

DG0 = -nFE

0

RT

nF

E = E0 -

ln Q

0

0

E =

E =

E

E

0.0257 V

0.0592 V

log Q

ln Q

n

n

-

-

19.5 The Effect of Concentration on Cell Emf

DG = DG0 + RT ln Q

DG = -nFE

-nFE = -nFE0+ RT ln Q

Nernst equation

At 298 K

ln = 2.303log


19 2 galvanic cells 19 3 standard reduction potentials 19 4 spontaneity of redox reactions

19.8 Electrolysisis the process in which electrical energy is used to cause a non spontaneous chemical reaction to occur.


19 2 galvanic cells 19 3 standard reduction potentials 19 4 spontaneity of redox reactions

Electrolysis of Water

19.8


19 2 galvanic cells 19 3 standard reduction potentials 19 4 spontaneity of redox reactions

Electrolysis and Mass Changes

Quantitative Aspects

Case (i) Na + + 1eNa

1 mol. of electron produces 1 mol of Na Atom(22g)

1 F (96500 C)

Case (ii) Mg 2+ + 2eMg

2 mol. of electron produces 1 mol of Mg Atom(24g)

2 F (2x 96500C)

Case (iii) Al 3+ + 3eAl

3 mol. of electron produces 1 mol of Al Atom(26g)

3 F (3 x 96500 C)


19 2 galvanic cells 19 3 standard reduction potentials 19 4 spontaneity of redox reactions

charge ( C ) = current (A) x time (s)

1 mole of electron = 96500 coulomb

1 mol. of Na atom = 22 g

1 mol. of Mg atom = 24 g

1 mol. of Al atom = 26 g


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