Electrochemical Cells

Definitions• Voltaic cell (battery): An electrochemical cell or group of

cells in which a product-favored redox reaction is used to product an electric current.

• Electrochemical cell: A combination of anode, cathode, and other materials arranged so that a product-favored redox reaction can cause a current to flow or an electric current can cause a reactant-favored redox reaction to occur

• Galvanic cell: A cell in which an irreversible chemical reaction produces electrical current

• Electrolytic cell: electrochemical reactions are produced by applying electrical energy

OutlineElectrochemistry

– Electrochemical cells– Modeling a electrochemical cells– Standard Hydrogen Electrodes– Using standard reduction potentials– Nernst– Quantifying current– Electrolysis

A Copper-Zinc battery – What Matters?

A Copper-Zinc battery – What Matters?What occurs atCopper electrode?

Copper plates out,mass increases

What occurs at Zinc electrode?Zinc is oxidized anddissolves, mass decreases

A Copper-Zinc battery – What Matters?

Consider reduction potentials:

Cu+2 + 2e- → Cu(s) 0.3419 VZn+2 + 2e- → Zn(s) -0.7618 V

Place Zn electrode in Copper Sulfate Solution – What happens?

Cu+2 + 2e- → Cu(s) 0.3419 VZn(s) → Zn+2 + 2e- 0.7618 V

Cu+2 + Zn(s) → Zn+2 + Cu(s) 1.1 V E > 0, spontaneous

Note, no need for electron to flow external to cell for reaction to occur!!

Copper is plated on Zn electrode

A Copper-Zinc battery – What Matters?

Consider reduction potentials:

Cu+2 + 2e- → Cu(s) 0.3419 VZn+2 + 2e- → Zn(s) -0.7618 V

Place Cu electrode in Zinc Sulfate Solution – What happens?

Cu(s) → Cu+2 + 2e- -0.3419 VZn+2 + 2e- → Zn(s) -0.7618 V

Zn+2 + Cu(s) → Cu+2 + Zn(s) -1.1 V E < 0, not spontaneous

No reaction occurs !!

Zn doesn’t plate on copper electrode?!

                                                                             

              

What are the ½ reactions?What is the overall reaction?

Identify the oxidation, reduction, anode, and cathode

SHE: Standard Hydrogen Electrode

2 H3O+(aq, 1.00 M) + 2e- <-> H2(g, 1 atm) + 2H2O(l)

Eo = 0V

Standard conditions:1M, 1atm, 25oC

Measuring Relative Potentials

                                                            

               

Standard Reduction Potentials

The half-reaction with the more positive standard reduction potential occurs at the cathode as reduction.

The half-reaction with the more negative standard reduction potential occurs at the anode as oxidation.   Eo

cell = Eocathode - E

oanode

Is potential always the same?

Standard conditions: 1 atm, 25oC, 1 M

What will influence the potential of a cell?

Mathematical Relationships: Nernst

The Nernst Equation:  Eo = standard potential of the cellR = Universal gas constant = 8.3145 J/mol*KT = temperature in Kelvinn = number of electrons transferredF = Faraday’s constant = 96,483.4 C/mol Q = reaction quotient (concentration of anode divided by the concentration of the cathode)

E = Eo - RT ln Q nF

Cu+2 + Zn(s) → Zn+2 + Cu(s) Q = [Zn+2]/[Cu+2]

Applying the Nernst Equation

This cell is operating at 25oC with 1.00x10-5M Zn2+ and 0.100M Cu2+?

Predict if the voltage will be higher or lower.

E = Eo - RT ln Q nF

Cu+2 + Zn(s) → Zn+2 + Cu(s)

Eo = standard potential of the cell

R = Universal gas constant = 8.3145 J/mol*K

T = temperature in Kelvin

n = number of electrons transferred

F = Faraday’s constant = 96,483.4 C/mol

Q = reaction quotient (concentration of anode divided by the concentration of the cathode)

E = Eo - RT ln Q nF Zn+2 + 2e- -> Zn -0.76 V

Cu+2 + 2e- -> Cu 0.34 V Eo

cell = Eocathode - E

oanode

25oC + 273 = K

n = 2

1.00x10-5M Zn2+ and 0.100M Cu2

Cu+2 + Zn(s) → Zn+2 + Cu(s)

Q = [Zn+2]/[Cu+2]

Were your predictions correct?