Electrochemistry
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Electrochemistry. The electricity produced by chemical reactions or ….. The chemical changes brought about by electricity. Electrochemical reactions = Oxidation-Reduction Reactions. Cell = System where chemical reactions occur. Electrode = means of adding/removing electric current

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Electrochemistry

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Electrochemistry

Electrochemistry

The electricity produced by chemical reactions or …..

The chemical changes brought about by electricity.

Electrochemical reactions = Oxidation-Reduction Reactions

Cell = System where chemical reactions occur

Electrode = means of adding/removing electric current

to/from system.

Cathode = electrode at which reduction occurs. (usually -)

Anode = electrode at which oxidation occurs. (usually +)

Conductivity = flow of electrons

many solid metals conduct electricity

electrolyte solutions (acids, bases, or soluble ionic cpds)


Electrochemistry

Electrochemistry – Cell Types

Electrolytic Cells – External source of electricity drives

a nonspontaneous (DG > 0) reaction.

Electrolysis

Used to convert water into H2 and O2 gas.

Used for electroplating – put a thin layer of one metal unto

another metal at cathode.

Voltaic Cells – A spontaneous (DG < 0) reaction is used to

generate an electric current.

e.g. batteries


Electrochemistry

1.100 V

e-

e-

+

Zn

Cu

Salt bridge (KCl – 5% agar)

Cl-

K+

Zn2+

Cu2+

1M ZnSO4

Zn(s) Zn2+ + 2e-

Oxidation - anode

Cu2+ + 2e-  Cu(s)

reduction - cathode

1M CuSO4

Voltaic Cell

Zn Cu

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


Electrochemistry

0.462 V

e-

e-

+

Cu

Ag

Salt bridge (KCl – 5% agar)

Cl-

K+

Cu2+

Ag+

1M CuSO4

Cu(s) Cu2+ + 2e-

Oxidation - anode

Ag+ + e-  Ag(s)

reduction - cathode

1M AgNO3

Voltaic Cell

Cu|Cu2+(1M)||Ag+(1M)|Ag


Electrochemistry

0.763 V

e-

e-

+

Zn

H2

Salt bridge (KCl – 5% agar)

E (SHE) = 0

Cl-

K+

E = -0.763

Zn2+

H+

1M ZnCl2

Zn(s) Zn2+ + 2e-

Oxidation - anode

2H+ + 2e-  H2(g)

reduction - cathode

1M HCl

SHE

Standard Hydrogen

electrode

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

Zn


Electrochemistry

0.337 V

e-

e-

+

Cu

Salt bridge (KCl – 5% agar)

E (SHE) = 0

Cl-

K+

E = 0.337

H+

Cu2+

1M HCl(aq)

H2(g) 2H+ + 2e-

Oxidation - anode

Cu2+ + 2e-  Cu(s)

reduction - cathode

1M CuSO4

Voltaic Cell

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

H2 Cu


Electrochemistry

Reduction ½ rx Standard Reduction Potential

E volts

Zn2+ + 2e-  Zn(s) -0.763 V

Fe2+ + 2e-  Fe(s) -0.44 V

Ni2+ + 2e-  Ni(s) -0.25 V

Pb2+ + 2e-  Pb(s) -0.126 V

2H+ + 2e-  H2(g) 0.000 V

Cu2+ + 2e-  Cu(s) +0.337 V

Hg2+ + 2e-  Hg(s) +0.789 V

Ag2+ + e-  Ag(s) +0.799 V

Cl2 + 2e-  2Cl- +1.360 V

The higher up a ½ rx is on the table the more readily that

element/substance is oxidized. (Reverse ½ reaction & E˚)


Electrochemistry

Calculating The Standard Electrical Potential for any cell

from the tabulated Standard Reduction Potentials.

1. Choose the appropriate ½ rxs from table

2. Write the ½ rx for the more (+) or less (-) substance

  • Write the ½ rx for the less (+) or more (-) substance

  • as an oxidation reaction. (reverse the sign on E)

  • Write the net balanced reaction (the electrons must be balanced

  • but do not multiply E by the balancing coefficient!)

  • Sum the values of E for each ½ rx to get the Cell’s

  • standard electrical potential.


Electrochemistry

Reduction ½ rx Standard Reduction Potential

E volts

Zn2+ + 2e-  Zn(s) -0.763 V

Fe2+ + 2e-  Fe(s) -0.44 V

Ni2+ + 2e-  Ni(s) -0.25 V

Pb2+ + 2e-  Pb(s) -0.126 V

2H+ + 2e-  H2(g) 0.000 V

Cu2+ + 2e-  Cu(s) +0.337 V

Hg2+ + 2e-  Hg(s) +0.789 V

Ag2+ + e-  Ag(s) +0.799 V

Cl2 + 2e-  2Cl- +1.360 V

The higher up a ½ rx is on the table the more readily that

element/substance is oxidized. (Reverse ½ reaction & E˚)


Electrochemistry

Reaction Spontaneity and DG

DG < 0  reaction is spontaneous – proceeds as written

DG = 0  reaction is at equilibrium

DG > 0  reaction proceeds in the reverse direction unless

enough energy provided to drive reaction forward.

DG = DG + RT ln Q

DG = - RT ln K

  • Represents standard conditions where P = 1atm

    and the [ ]s of all reagents are 1M. Q = 1


Electrochemistry

Reaction Spontaneity and E

E > 0  reaction is spontaneous – proceeds as written

with voltage output

E = 0  reaction is at equilibrium – no current flow

E < 0  reaction proceeds in the reverse direction

unless enough current provided to drive reaction

forward. (e.g. electrolysis)

The Nernst Equation

E = E - 2.303RT log Q

nF

DG = DG + RT ln Q


Electrochemistry

??? V

e-

e-

+

Zn

Cu

Salt bridge (KCl – 5% agar)

Cl-

K+

Zn2+

Cu2+

1.2 M ZnSO4

Zn(s) Zn2+ + 2e-

Oxidation - anode

Cu2+ + 2e-  Cu(s)

reduction - cathode

0.80 M CuSO4

Voltaic Cell

Zn Cu

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

E = -0.763 - 0.0592 log(1/1.2)

n


Electrochemistry

The Nernst Equation

E = E - 2.303 RT log Q lump constants together….

nF

E = E - 0.0592 log Q at 25 C

n

E = 0.0592 log K = RT ln K at 25 C

n nF

Zn|Zn2+(1M)||Cu2+(1M)|Cu E = 1.10 V

Zn|Zn2+(1.2M)||Cu2+(0.8M)|Cu E = ??? V


Electrochemistry

96,485 C (J V-1 mol-1)

Reaction Spontaneity and the Nernst Equation

E = E - 0.0592 log Q at 25 C

n

E = 0.0592 log K = RT ln K at 25 C

n nF

DG˚ = -RT ln K

DG˚ = -nFE˚ or DG = -nFE

Calculate DG˚ or DG and K from cell potentials, E˚/E.


Electrochemistry

anode

cathode

Electrolysis of Water

2H2O2H2(g) + O2(g)

Reduction: (2H+ + 2e- H2(g)) 2

Oxidation: 2H2O 4H+ + O2 + 4e-


Electrochemistry

Na+ Cl-

Na+ + e- Na (floats)

Reduction - cathode

2Cl- Cl2(g) + 2e-

oxidation - anode

Graphite or Platinum

are common inert

electrodes

+

Molten NaCl (801˚ C)

Electrolysis of molten NaCl


Electrochemistry

Electron Stoichiometry

A Coulomb (C) is the SI unit of charge

1 e- = 1.602 x 10-19 C or 1.602 x 10-19 C per e-

1 mole of e- = 96,485 C = 1 Faraday (F)

Current = charge per time = C s-1 = Ampere (A)

How many grams of H2 gas can be produced from water

through which 1.35 x 106 C have been passed?

Reduction: 2H+ + 2e- H2(g)


Electrochemistry

3.4 amps for 15 seconds –

How much Cu is electroplated?

Cu(s)→ Cu2+ + 2e-

oxidation – anode

electrode ‘dissolves’

Cu2+ + 2e- → Cu(s)

reduction – cathode

Cu plated onto electrode

+

Cu2+ SO42-

Electroplating copper

3.4C x 15s x 1 mol e- x 1 mol Cu x 63.546g Cu = 16.8 mg

1s 96,485C 2 mol e- 1 mol Cu


Electrochemistry

Anode = Zn

Cathode = C

1.6 V

Dry Cell

ZnCl2; NH4Cl; MnO2 paste

Zn + 2NH4+→ Zn2+ + 2NH3 + H2

Anode: (oxidation) Zn → Zn2+ + 2e-

Cathode: (reduction) 2NH4+ + 2e-→ 2NH3 + H2

2MnO2 + H2→ 2MnO(OH) (removes H2 gas)

Zn2+ + 4NH3 → [Zn (NH3)4]2+ (removes NH3)


Electrochemistry

Anode = Zn

Cathode = C

1.5 V

Alkaline

Cell

ZnCl2; KOH; MnO2 paste

Zn + 2MnO2 + 2H2O → Zn(OH)2(s) + 2MnO(OH)

Anode: (oxidation) Zn + 2OH-→ Zn(OH)2 + 2e-

Cathode: (reduction) 2MnO2 + 2H2O+ 2e-→ 2MnO(OH) + 2OH-

Alkaline cells have longer shelf lives


Electrochemistry

Anode = Cd

Cathode = NiO2

1.4 V

Ni - Cd

rechargeable

Ni(OH)2(s); KOH; Cd(OH)2(s)

Cd + NiO2(s) + 2H2O → Cd(OH)2(s) + Ni(OH)2(s)

Rechargeable Batteries – Nickel-Cadmium

Anode: (oxidation) Cd + 2OH-→ Cd(OH)2(s) + 2e-

Cathode: (reduction) NiO2 + 2H2O+ 2e-→ Ni(OH)2(s) + 2OH-


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