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Electrochemical Cells
Galvanic Cells
Voltaic / Galvanic Cell Apparatus which produce electricity
Electrolytic Cell Apparatus which consumes electricity
 Consider:
Zn
Initially there is a
flow of eAfter some time the
process stops
Electron transport
stops because of
charge build up
Cu
Chapter 9.5
Build up of
positive charge
Build up of
negative charge
The charge separation will lead to
process where it cost too much
energy to transfer electron.
Sign Convention of Voltaic Cell
Completing the Circuit
Electron transfer can occur if the circuit is closed
Parts:
@ Anode: Negative Terminal (anions).
Source of electron then repels electrons. Oxidation occurs.
Zn(s) g Zn+2(aq) + 2e- : Electron source
@ Cathode: Positive Terminal (cation)
Attracts electron and then consumes electron. Reduction occurs.
Electron target: 2e- + Cu +2(aq) g Cu(s)
3 process must happen if e- is to flow.
Two conductors
A. e- transport through external circuit
B. In the cell, ions must migrate
Electrolyte solution
Salt Bridge / Porous membraneC. Circuit must be closed (no charge build up)
Anode (-)
Cathode (+)
A
Black
Overall:
Red
Negative
electrode
generates
electron
Zn(s) + Cu+2(aq) g Zn+2(aq) + Cu(s)
Positive
electrode
accepts
electron
C
B
Oxidation
Occur
Sign of E is also reversed
E = -1.10 V
Zn(s) g Zn +2(aq)
Oxidation:
Reduction: Cu +2(aq) g
Reduction
Occur
Cathode/Cation(+)
CELL
= E
red
(Cathode) - E
Line notation: Convenient convention for electrochemical cell
Consider :
Schematic Representation

Anode: Zn g Zn+2 + 2e-

Cathode:
2.
Anode
[oxidation (-)]

E = 0.34 V
ox
CELL
(anode)
Line Notation Examples
Line Notation Convention
1.
E = 0.76 V
Cu(s)
1.10 V = E
or E
Anode/Anion (-)
E = 1.10 V
Note when the reaction is reverse: Cu(s) + Zn+2(aq) g Cu+2(aq) + Zn(s)
Cathode
[reduction (+)]
“ “ phase boundary
Zn(s) + Cu +2(aq) g Zn+2(aq) + Cu (s
Cu+2 + 2e- g Cu
Shorthand “Line” notation

Zn (s) Zn+2
(aq)(1.0M)
Cu+2(aq) (1.0M) Cu(s)
(where potential may develop)
2nd Example : Zn(s) + 2H+ (aq) g Zn+2(aq) + H2(g)
3.
“ “ Liquid junction
4.
Concentration of component
Anode: Zn g Zn+2 + 2e-

Cathode:
2H+ + 2e- g H2 (g)
Shorthand “Line” notation
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1
Zn(s) ZnSO4 (aq,1.0M)

3
CuSO4 (aq,1.0M) Cu(s)

Zn (s) Zn+2 (aq)(1.0M)
H+(aq) (1.0M), H2(g, 1atm) Pt(s)
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Other Voltaic Cell & Their Line Notation
Zn(s) | Zn+2 (aq)||H+(aq) , H2 (g,1atm)|Pt
C(s)| I-(aq) , I2 (g,1atm) || MnO4-(aq) , Mn+2
Cr(s) | Cr+3 (aq)||Ag+(aq) | Ag(s)
(aq) |
C(s)
Standard Hydrogen Electrode
 Under standard-state condition, the potential
for the reduction of H+ at 25 C is taken to be
EXACTLY ZERO.
2 H+ (1 M) + 2 e-  H2 (1 atm) E = 0 V
 Think about
Zn(s) | Zn2+ (1M) || Cu2+ (1M) | Cu (s)
the measured emf =1.10 V at 25 C
 The redox reaction is spontaneous in the
forward direction if the standard emf of the
cell is positive.
Example
 A galvanic cell consists of a Mg electrode in a
1.0 M Mg(NO3)2 solution and a Ag electrode in
a 1.0 M AgNO3 solution. Calculate the standard
emf of this electrochemical cell at 25 C.
Standard Reduction Potential
 The voltage associated with a reduction
reaction at an electrode when all solutes are
1 M and all gases are at 1atm.
 The hydrogen electrode serves as the
reference for calculating the potential of a
single electrode.
-hydrogen gas is bubbled into a hydrochloric
acid solution at 25 C;
-the platinum electrode provides the surface
and serves as an electrical conductor.
Rules for using E
1. Applies to reactions as written.
2. The more positive E , the great tendency to
be reduced (the stronger oxidizing agent).
3. The half-cell reactions are reversible and
the sign of E will be changed.
4. Diagonal rule.
5. Multiplying reactions by some coefficient
does not change E (intensive property –
not affected by size of electrode or volume
of electrolyte, but concentration must be
1M).
Standard Reduction Potentials
E
cell =
E
reduction
–E
oxidation
Modern voltmeters can read “+” or “–” volts. By
always connecting the negative lead to SHE, the
voltmeter reads E of
the half-cell directly.
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Predicting Spontaneity
Guess which half-cell is the anode.
Write half-reactions accordingly.
Calculate E cell.
If E cell > 0, reaction is spontaneous
(proceeds in the forward direction as
written).
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