Electrochemistry.ppt

Electrochemistry

Electrochemistry



_{Electrochemistry}

 _{deals with interconversion between chemical and} electrical energy

 _{involves redox reactions}

Redox reactions (quick review)

Redox reactions (quick review)



_Oxidation

 _{loss of electrons}



_Reduction

 _{gain of electrons}



_{Reducing agent}

 _{donates the electrons and is oxidized}



_{Oxidizing agent}

Redox Reactions

Redox Reactions



_{Direct redox reaction}

 _{Oxidizing and reducing agents are mixed together}



_{Indirect redox reaction}

 _{Oxidizing and reducing agents are separated but} connected electrically

• Example

– Zn and Cu2+ can react indirectly

 _{Basis for electrochemistry}

Electrochemistry Terminology

Electrochemistry Terminology

Electrochemical Cell -- a device that converts electrical energy into chemical energy or vice versa

 _{Two Types:}

Electrolytic cell

Converts electrical energy into chemical energy Electricity is used to drive a non-spontaneous reaction

Galvanic (or voltaic) cell

Converts chemical energy into electricity (a battery!) A spontaneous reaction produces electricity

 _Conduction:

Metals: metallic (electronic) conduction -- free movement of electrons

Electrochemical Cells

Electrochemical Cells

Electrochemical Cells



_{Voltaic Cell}

 _{cell in which a spontaneous redox reaction generates} electricity

Electrochemical Cells

Voltaic Cell

Electrochemical Cells

Electrochemical Cells

Electrochemical Cells



_{Electrolytic Cell}

 _{electrochemical cell in which an electric current} drives a nonspontaneous redox reaction

Cell Potential

Cell Potential



Cell Potential (electromotive force), E

(V)

 _{electrical potential difference between the two} electrodes or half-cells

• Depends on specific half-reactions, concentrations, and temperature

• standard state conditions

high electrical

high electrical

potential

potential

low electrical

low electrical

potential

potential

Cell Potential

Cell Potential

Cell Potential

E

= E

- E

= E

- E

E°

= E°

- E°

= E°

- E°

Cell Potential

Cell Potential



_E°

_{= E°}

_{- E°}

_{= E°}

_{- E°}



_{Reference electrode}

has arbitrarily assigned E

_{used to measure relative Ecathode and Eanode for half-cell} reactions



_{Standard hydrogen electrode (S.H.E.)}

Standard Hydrogen Electrode

Standard Hydrogen Electrode



E = 0 V (by

definition; arbitrarily

selected)

Standard Electrode Potentials

Standard Electrode Potentials



_{Standard Reduction Potentials, E°}

 E°_cell measured relative to S.H.E. (0 V)

• electrode of interest = cathode

 _{If E° < 0 V:}

• Oxidizing agent is harder to reduce than H+

 _{If E° > 0 V:}

Standard Reduction Potentials Standard Reduction Potentials

Reduction Half-Reaction _E(V) F₂(g) + 2e- _2F-_{(aq) 2.87}

Au3+_{(aq) + 3e}- _{Au(s) 1.50}

Cl₂(g) + 2 e- _2Cl-_{(aq) 1.36}

Cr₂O₇2-_{(aq) + 14H}+_{(aq) + 6e}- _2Cr3+_{(aq) + 7H}

2O 1.33

O₂(g) + 4H+ _{+ 4e}- _2H

2O(l) 1.23

Ag+_{(aq) + e}- _{Ag(s) 0.80}

Fe3+_{(aq) + e}- _Fe2+_{(aq) 0.77}

Cu2+_{(aq) + 2e}- _{Cu(s) 0.34}

Sn4+_{(aq) + 2e}- _Sn2+_{(aq) 0.15}

2H+_{(aq) + 2e}- _H

2(g) 0.00

Sn2+_{(aq) + 2e}- _{Sn(s) -0.14}

Ni2+_{(aq) + 2e}- _{Ni(s) -0.23}

Fe2+_{(aq) + 2e}- _{Fe(s) -0.44}

Zn2+_{(aq) + 2e}- _{Zn(s) -0.76}

Al3+_{(aq) + 3e}- _{Al(s) -1.66}

Mg2+_{(aq) + 2e}- _{Mg(s) -2.37}

Uses of Standard Reduction

Uses of Standard Reduction

Potentials

Potentials



_{Compare strengths of reducing/oxidizing}

agents.

Standard Reduction Potentials Standard Reduction Potentials

Reduction Half-Reaction _E(V) F₂(g) + 2e- _2F-_{(aq) 2.87}

Au3+_{(aq) + 3e}- _{Au(s) 1.50}

Cl₂(g) + 2 e- _2Cl-_{(aq) 1.36}

Cr₂O₇2-_{(aq) + 14H}+_{(aq) + 6e}- _2Cr3+_{(aq) + 7H}

2O 1.33

O₂(g) + 4H+ _{+ 4e}- _2H

2O(l) 1.23

Ag+_{(aq) + e}- _{Ag(s) 0.80}

Fe3+_{(aq) + e}- _Fe2+_{(aq) 0.77}

Cu2+_{(aq) + 2e}- _{Cu(s) 0.34}

Sn4+_{(aq) + 2e}- _Sn2+_{(aq) 0.15}

2H+_{(aq) + 2e}- _H

2(g) 0.00

Sn2+_{(aq) + 2e}- _{Sn(s) -0.14}

Ni2+_{(aq) + 2e}- _{Ni(s) -0.23}

Fe2+_{(aq) + 2e}- _{Fe(s) -0.44}

Zn2+_{(aq) + 2e}- _{Zn(s) -0.76}

Al3+_{(aq) + 3e}- _{Al(s) -1.66}

Mg2+_{(aq) + 2e}- _{Mg(s) -2.37}

Li+_{(aq) + e}- _{Li(s) -3.04}

Uses of Standard Reduction

Uses of Standard Reduction

Potentials

Potentials



_{Determine if oxidizing and reducing agent react}

spontaneously

Cathode

(reduction) E°redn (cathode)

more +

Anode (oxidation)

E

° red

n

(

V

)

E°_redn (anode)

more

-Spontaneous rxn if

Spontaneous rxn if

E

Uses of Standard Reduction

Uses of Standard Reduction

Potentials

Potentials



Calculate E°

 E°_cell = E°_cathode - E°_anode

• Greater E°_cell, greater the driving force

Cell Potential

Cell Potential



Relationship between E

and G:

 G = -nFE_cell

• F = Faraday constant = 96500 C/mol e-’s, n = # e-’s

transferred redox rxn.

Equilibrium Constants from E

Equilibrium Constants from E



Relationship between E

and G:

 G = -nFE_cell

• F = Faraday constant = 96500 C/mol e-’s, n = # e-’s

transferred redox rxn

G < 0, E_cell > 0 = spontaneous



_{Under standard state conditions:}

Equilibrium Constants from E

Equilibrium Constants from E

 _{Relationship between Ecell and} _G:

 _{G = -nFEcell}

• F = Faraday constant = 96500 C/mol e-’s, n = # e-’s transferred redox rxn

• 1 J = CV

 G < 0, E_cell > 0 = spontaneous

 _{Under standard state conditions:}

 _{G° = -nFE°cell}

and

 G° = -RTlnK

so

The Nernst Equation

The Nernst Equation



G depends on concentrations

 G = G° + RTlnQ

and



G = -nFE

and G° = -nFE°

thus



-nFE

= -nFE°

+ RTlnQ

or

The Nernst Equation

The Nernst Equation



E

= E°

- (RT/nF)lnQ (Nernst eqn.)

 _{At 298 K (25°C), RT/F = 0.0257 V}

so



E

= E°

- (0.0257/n)lnQ

or



_{a A + b B = c C + d D}

and

[C]

[D]