Electrochemistry : Electrochemistry The study of the interchange
of chemical and electrical energy. Sample electrochemical processes:
1) Corrosion
4 Fe(s) + 3 O2(g) ? 2 Fe2O3(s)
2) Biological processes
C6H12O6 + 6 O2 ? 6 CO2 + 6 H2O
3) Batteries (Galvanic or Voltaic cells)
Electrochemical cells that produce a current (flow of electrons) as a result of a redox reaction
4) Electrolytic cells
Electrical energy is used to produce chemical change
Used to prepare or purify metals (such as sodium, aluminum, copper)
Chemical Change ? Electron Flow : Chemical Change ? Electron Flow Copper: Cu(s), Cu2+(aq)
Cu(s) ? Cu2+(aq) + 2e-
?G°rxn = ?G°f(Cu2+) = 65.6 kJ
Silver: Ag(s), Ag+(aq)
Ag(s) ? Ag+(aq) + e-
?G°rxn = ?G°f(Ag+) = 77.2 kJ Cu(s) ? Cu2+(aq) + 2e- ?G° = +65.6 kJ Ag+(aq) + e- ? Ag(s) ?G° = -77.2 kJ Cu(s) + 2 Ag+(aq) ? Cu2+(aq) + 2 Ag(s) ?G° = -88.8 kJ Spontaneous
wmax = -88.8 kJ Cu2+ in solution 2( ) 2( ) Ag(s)
Harnessing the Energy : Harnessing the Energy Separate the half-reactions
Creates a galvanic or voltaic cell Cu Ag 1 M CuSO4
Cu(s) ? Cu2+(aq) + 2e- 1 M AgNO3
Ag+(aq) + e- ? Ag(s) Luigi Galvani Alessandro
Volta Cu2+
SO42- Ag+
NO3- KNO3(aq) salt bridge Oxidation Reduction Anode Cathode -
(produces electrons) +
(attracts electrons) Red Cat
Line Notation for Galvanic Cells : Line Notation for Galvanic Cells Cu Ag 1 M CuSO4
Cu(s) ? Cu2+(aq) + 2e- 1 M AgNO3
Ag+(aq) + e- ? Ag(s) Cu2+
SO42- Ag+
NO3- Oxidation Reduction Anode
(-) Cathode
(+) Anode always on the left Cu(s)?Cu2+ (1 M)??Ag+ (1 M)?Ag(s) Cathode always on the right
Chemical Change ? Electrical Work : Chemical Change ? Electrical Work Chemical change produces electrical energy
Electrical energy can be used to do work! ?G = wmax Electrical work: w = -nFE
n = # of moles e- transferred
F = charge on a mole of e-
E = electrical potential
(electromotive force) Cell Potential (E) or Electromotive Force (emf): The driving force pushing the electrons from the anode to the cathode.
Units = Volts 1 Volt = 1 joule/coulomb
Standard Reduction Potentials : Standard Reduction Potentials The cell potential E°cell can be determined from the standard reduction potentials (E°red) for the half-reactions:
Reduction potential = tendency for reduction to happen
Positive E°red ? spontaneous reduction reaction
Negative E°red ? non-spontaneous reduction or spontaneous oxidation (reverse reaction)
Standard (o) = standard conditions (1 M solutions, 1 atm gases)
Standard Reduction Potentials : Standard Reduction Potentials Half-Reaction E° (V)
F2 + 2e- ? 2F- 2.87
Au3+ + 3 e- ? Au 1.50
Ag+ + e- ? Ag 0.80
Cu2+ + 2e- ? Cu 0.34
2H+ + 2e- ? H2 0.00
Ni2+ + 2e- ? Ni -0.23
Zn2+ + 2e- ? Zn -0.76
Al3+ + 3e- ? Al -1.66
Li+ + e- ? Li -3.05 ? E° = 0 (SHE)
Standard Hydrogen Electrode E° > 0
Spontaneous reduction E° < 0
Non-Spontaneous reduction Reduction potential = tendency for reduction to happen
Standard = standard conditions (1 M solutions, 1 atm gases) Spontaneous oxidation
(reverse rxn)
Standard Reduction Potentials : Standard Reduction Potentials Half-Reaction E° (V)
F2 + 2e- ? 2F- 2.87
Au3+ + 3 e- ? Au 1.50
Ag+ + e- ? Ag 0.80
Cu2+ + 2e- ? Cu 0.34
2H+ + 2e- ? H2 0.00
Ni ? Ni2+ + 2e- +0.23
Zn ? Zn2+ + 2e- +0.76
Al ? Al3+ + 3e- +1.66
Li ? Li+ + e- +3.05 E° > 0
Spontaneous reduction But remember, an oxidation CANNOT happen without a reduction Spontaneous oxidation
Standard Reduction Potentials : Standard Reduction Potentials Half-Reaction E° (V)
F2 + 2e- ? 2F- 2.87
Au3+ + 3 e- ? Au 1.50
Ag+ + e- ? Ag 0.80
Cu2+ + 2e- ? Cu 0.34
2H+ + 2e- ? H2 0.00
Ni2+ + 2e- ? Ni -0.23
Zn2+ + 2e- ? Zn -0.76
Al3+ + 3e- ? Al -1.66
Li+ + e- ? Li -3.05 Strongest Oxidizing Agent
(most easily reduced) Reduction potential = tendency for reduction to happen
Standard = standard conditions (1 M solutions, 1 atm gases) Strongest Reducing Agent
(most easily oxidized)
Cell Potential : Cell Potential E°cell = E°reduction + E°oxidation
Ag+(aq) + e- ? Ag(s) E° = 0.80 V
Cu2+(aq) + 2 e- ? Cu(s) E° = 0.34 V Reduction reaction: 2(Ag+(aq) + e- ? Ag(s)) E° = +0.80 V
Oxidation reaction: Cu(s) ? Cu2+(aq) + 2 e- E° = - 0.34 V Cu(s) + 2 Ag+(aq) ? Cu2+(aq) + 2 Ag(s) E°cell = +0.46 V E° is intensive, unlike ?Go Cu(s)?Cu2+ (1 M)??Ag+ (1 M)?Ag(s) The E°cell MUST be + and thus spontaneous for Galvanic cells
Free Energy and Cell Potential : Free Energy and Cell Potential ?G? = wmax = ?nFE?
n = number of moles of electrons transferred
F = Faraday’s constant
= 96,485 coulombs per mole of electrons (C/mol e-)
E° = standard cell potential (V or J/C) Michael Faraday Cu(s)?Cu2+ (1 M)??Ag+ (1 M)?Ag(s) E°cell = +0.46 V ?G° = -nFE°cell
?G° = -(2 mol e-)(96485 C/mol e-)(0.46 V)
?G° = -88,800 J or -88.8 kJ
Practice Time : Practice Time Given the following information, draw a galvanic cell.
Fe(s)?Fe2+(1 M)??Au3+(1 M)?Au(s)
Be sure to include the following:
Anode/Cathode reactions
Balanced overall reaction
Complete circuit (external wire with e- flow direction, salt bridge)
Label all parts of the cell (solution, electrode, etc.)
Fe(s)?Fe2+(1 M)??Au3+(1 M)?Au(s) : Fe(s)?Fe2+(1 M)??Au3+(1 M)?Au(s) Fe Au 1 M Fe2+
Fe(s) ? Fe2+(aq) + 2e- 1 M Au3+
Au3+(aq) + 3e- ? Au(s) Fe2+ Au3+ Oxidation Reduction Anode Cathode 3Fe(s) + 2Au3+(aq) ? 3Fe2+(aq) + 2Au(s) cations E°cell = +0.440V (Fe rxn) + 1.50 V (Au rxn) = 1.94 V
Reaction Quotient : Reaction Quotient The reaction quotient (Q) sets up a ratio of products and reactants
For a reaction, A + 2B ? 3C + 4D
[C]3[D]4
[A]1[B]2
Only concentrations (aq) or pressures (g) are used to solve for Q
Solids (s) and liquids (l) are not included in the expression Q =
Reaction Quotient practice : Reaction Quotient practice Write the Q expression for the following reaction
CH4(g) + O2(g) ? CO2(g) + H2O(g)
Reaction must be balanced first
CH4(g) + 2O2(g) ? CO2(g) + 2H2O(g)
(CO2)(H2O)2
(CH4)(O2)2 Q =
Reaction Quotient practice : Reaction Quotient practice Write the Q expression for the following reaction
Cu(s) + 2Ag+(aq) ? Cu2+(aq) + 2Ag(s)
(Cu2+)(Ag)2
(Cu)(Ag+)2
Is this correct?
NO: Solids aren’t included in the equation!
(Cu2+)
(Ag+)2 Q = Q =
Non-standard conditions:The Nernst Equation : Non-standard conditions:The Nernst Equation We can calculate the potential of a cell in which some or all of the components are not in their standard states (not 1 M concentration or 1 atm pressure). ?G = ?G° + RT lnQ ?G = -nFE ?G° = -nFE° -nFE = -nFE° + RT lnQ Walther Nernst R = 8.3415 J/mol K
T = temperature
n = moles of e-
F = Faraday’s constant
96,485 C/mol e-
Practice with the Nernst Equation : Practice with the Nernst Equation What will be the cell potential E of a Cu/Ag cell using 0.10 M Cu2+ and 1.0 M Ag+ solutions at 25°C? Cu(s) + 2 Ag+(aq) ? Cu2+(aq) + 2 Ag(s) E = 0.46 V – (-0.03 V) Cu(s)?Cu2+ (0.10 M)??Ag+ (1.0 M)?Ag(s) E = 0.49 V
Brain Warmup : Brain Warmup Half-Reaction E° (V)
Ag+ + e- ? Ag 0.80
Cu2+ + 2e- ? Cu 0.34
Zn2+ + 2e- ? Zn -0.76
Al3+ + 3e- ? Al -1.66 What is E° for each of the following reactions?
Which reaction(s) are spontaneous? 3 Ag+(aq) + Al(s) ? 3 Ag(s) + Al3+(aq)
Cu2+(aq) + Zn(s) ? Cu(s) + Zn2+(aq)
2 Al3+(aq) + 3 Zn(s) ? 2 Al(s) + 3 Zn2+(aq) E°
2.46 V
1.10 V
-0.90 V Spontaneous?
Y
Y
N Zn can reduce Cu2+, but not Al3+