Electrochemical Electrochemical SystemSystem
Week Topic Topic Outcomes
8-10 Electrochemical systems
•Electrochemical system and its thermodynamics
•Chemical reactions in electrochemical cells and the Nernstequation
•The relationship between the cell emf and the equilibrium constant, Gibbs energies and reaction entropies
•Thermodynamics of Galvanic cells and Fuel cell
It is expected that students are able to:
•Define the electrochemical system.
•Explain about electrochemical cell and differentiate galvanic cell and electrolyte cell
•Apply Nernst equation of thermodynamic principles to electrochemical cells.
•Determine ∆Go, ∆Ko, ∆So, ∆Ho, ∆Cpo
activity coefficient and pH of cell’s chemical reaction.
•Understand the cell emf in IUPAC convention.
•Estimate the liquid junction potential from emf measurement..
Topic Outcomes
Redox: Reduction-Oxidation, an electron transfer process
Oxidation: loss one or more e-
Reduction: gain one or more e-
Oxidation number: the charge of the atom or molecule would have
Pure element and neutral compound 0
Mono- and poly-atomic ions sum of oxidation number is ionic charge
Quick Review
CaBr2; Ca = +2, Br = -1,Oxidation number (pure element), =(+2)+2(-1)=0
Check your understanding
K2Cr2O7
O = -2 K = +1
Cr = ?
HCO3-
O = -2 H = +1
C = ?
Electrochemistry: study of the interchange between chemical change and electrical work
Electrochemical cells: systems utilizing a redoxreaction to produce or use electrical energy
Quick Review (Cont’d)
Becomes smaller
Becomes larger
Znanode
Cucathode
Cell and Cell Notation
Anode: Electrode site of oxidation
Cathode: Electrode site of reduction
Electrons flow from anode to cathode
Salt bridge: Maintains electrical neutrality, allows current to flow, prevent solutions mixing
Positive ion migrates to cathode
Negative ion migrates to anode
Cell notation: Anode | Salt Bridge | Cathode
The difference in electrical potential between the anode and cathode is called:
• cell voltage
• electromotive force (emf)
• cell potential
000reductionoxidationCell EEE += (Unit: Volts)
Note: Volt (V) = Joule (J) Coulomb (C)
Cell Potential
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.
E0 = 0 VReference electrode
Standard hydrogen electrode (SHE)
2e- + 2H+ (1 M) H2 (1 atm)
Reduction Reaction
Standard Electrode Potentials
(Anode)
(Cathode)
Note: Copper reduction potential, Eo = +0.34(If copper is reduced (at cathode), voltmeter can give a reading- positive)
Note: Zn reduction potential is negative.Change Zn as anode
E0 is for the reaction as written
The half-cell reactions are reversible; the sign of E0
changes when the reaction is reversed
Changing the stoichiometriccoefficients of a half-cell reaction does not change the value of E0
The more positive E0 the greater the tendency for the substance to be reduced
+ E°CELL : spontaneous reaction
E°CELL = 0 :equilibrium
- E°CELL: non-spontaneous reaction
More positive E°CELL means stronger oxidizing agent or more
likely to be reduced
Spontaneity of Redox Reaction000reductionoxidationCell EEE +=
Eoreduction=-Eooxidation
Write separate equations (half-reactions) for oxidation and reduction
For each half-reaction
Balance elements involved in e- transfer
Balance number e- lost and gained
To balance e- multiply each half-reaction by whole numbers
Add half-reactions/cancel like terms (e-)
Check that all atoms and charges balance
Balance the Half-Cell Reactions
e–
e – e–
e–
Zn 2+
SO4 2–
Zn(s)1.0 M Zn 2+
solution
Anode
1.0 M Cu 2+
solution
Cathode
Cu 2+
SO4 2–
Cu(s)
Check your understanding
??
?? ??
Check your understanding
A half-cell reduction reaction is
Fe(OH)2 (s) + 2e- Fe (s) + 2OH- (aq) Eo=-0.877 V
Is to combine with another half-cell, which the reduction reaction is given as:
a) Al3+ (aq) + 3e- Al (s) Eo= -1.66 V
b) AgBr (s) + e- Ag (s) + Br- (aq) Eo=+0.071 V
Determine Eocell, and whether Fe is oxidized or reduced.
∆G =∆G0 + RTlnQ ∆G0 = -nFE0cell
-nFEcell= -nFE0cell + RTln Q
Ecell= E0cell - RTln QnF
Ecell= E0cell - 0.0257ln Qn
Ecell= E0cell – 0.0592log Qn
Nernst Equation
Note: At equilibrium, ∆G=0, thus Ecell=0
the emf for an electrochemical cell to be calculated if the activity and E°are known.
Faraday, F: charge on 1 mole e-F = 96485 C/mole
∆Go = -nFEocell
ΔG0 = -RT ln K
ΔG0 = -nFE0cell-RT ln K = -nFE0cell
KnFRTECell ln
0 =( )
0257.096485
29831.8=
⎟⎠⎞
⎜⎝⎛
=
moleC
KmolK
J
FRT
Kn
Kn
ECell log0592.0ln0257.00 ==
n= moles of balanced e-
Entropy Change
The reaction entropy related to ΔG°R
pp
RR T
EnFTGS ⎟⎟
⎠
⎞⎜⎜⎝
⎛∂∂
=⎟⎟⎠
⎞⎜⎜⎝
⎛∂Δ∂
−=Δoo
o
Measurement of the T dependence of E° can be used to determine ΔS°R
To obtain the entropy change for the cell reaction
The reaction entropy related to ΔG°R
PTEnFTnFEH ⎟⎟
⎠
⎞⎜⎜⎝
⎛∂∂
+−=Δo
oo
Enthalpy Change
∆Ho = ∆Go + T∆So
Example (from text book)
The half-cell reduction reactions,
Fe3+ (aq) + e- Fe2+ (aq) Eo=+0.771 V
Fe2+ (aq) + 2e- Fe (s) Eo=-0.447 V
Calculate Eo for half-cell reduction reaction
Fe3+ (aq) + 3e- Fe (s)
Example
Calculate ΔGor and the equilibrium constant at 298.15K for two half-cell reduction reactions:
Cr2O72-(aq) + 14H+(aq) + 6e- → 2Cr3+(aq) + 7H2O(l)
Eº = +1.232 V
2H+(aq) + 2e- → H2(g)
Eº = 0.00 V
The standard potential E° for a given cell in which a two
electron transfer takes place is 1.10 V at 298.15 K and
(∂E°/∂T) = –6.50 × 10–5 V K–1. Calculate ΔG°rxn, ΔS°rxn,
ΔH°rxn. Assume that n = 2.
Example
Check your understanding
Calculate the equilibrium constant at
298.15 K for the following half-cell reactions
a) Cathode: NO3- + 4H+ + 3e- → NO + 2H2O Eº = +0.957 V
Anode: 2H2O → O2 + 4H+ + 4e- Eº = –1.229 V
b) Anode: Cd(OH)2 + 2e- → Cd + 2OH- Eº = –0.809V
Cathode: O2 + 2H2O + 4e- → 4OH- Eº = +0.401V
Calculate the equilibrium constant at 25°C for the
reaction occurring in the Daniell cell, if the standard
EMF is 1.10 V.
Calculate the equilibrium constant for the reaction
H2 + 2Fe3+ ⇔ 2H+ + 2Fe2+
Example
Galvanic cell vs. Galvanic cell vs. Electrolytic cellElectrolytic cell
Galvanic cell vs. Electrolytic cell
The external battery supplies the electrons. They enter through the cathode and come out through the anode.
The electrons are supplied by the species getting oxidized. They move from anode to the cathode in the external circuit.
Here, the anode is positive and cathode is the negative electrode. The reaction at the anode is oxidation and that at the cathode is reduction.
Anode is negative and cathode is the positive electrode. The reaction at the anode is oxidation and that at the cathode is reduction.
Both the electrodes are placed in a same container in the solution of molten electrolyte
The two half-cells are set up in different containers, being connected through the salt bridge or porous partition
Redox reaction is not spontaneous and electrical energy has to be supplied to initiate the reaction
Redox reaction is spontaneous and is responsible for the production of electrical energy
Converts electrical energy into chemical energy
Converts chemical energy into electrical energy
Electrolytic cellElectrochemical cell (Galvanic Cell)
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