Ch 12 electrochemistry

DISCLAIMER: these notes are provided to assist you with mastering the course material but they are not intended as a replacement of the lectures.
Neither do they contain the comments and ancillary material of the lectures; they are just a set of points that you might bring to the lectures to annotate
instead of having to write everything down and/or they may assist you in organizing the material after the lectures, in conjunction with your own notes.
Ch 12 Electrochemistry
• Refer back to chapter 3.7 for oxidation state/number;
• Recall: oxidation as loss of electrons; reduction as gain of electrons; oxidation - reduction
(or, redox) reactions; balancing redox equations; redox titrations (introduced in chapter 6.4,
pp. 163-170)
12.1 Electrochemical Cells
• energy of a spontaneous redox reaction can be used to do electrical work in a voltaic (=
galvanic) cell by forcing electron transfer through an external path
• eg., first with direct transfer, then indirect/external (alternative to Zn/Cu example in
Introduction):
Ag(s)2(aq.)Cu(aq.)Ag2Cu(s) 2
+→+ ++
• direct reaction: Cu strip in a beaker of Ag+
(aq.)
• blue colour of Cu2+
gradually builds up
• Cu begins to dissolve
• Ag begins to deposit (black) on Cu strip
• reaction via external path: Voltaic Cell
• 2 metal strips (= electrodes) in separate compartments or beakers
• Cu2+
and Ag+
solutions as nitrates, one in each
• 2 means of connection of solution
• salt bridge, Fig. 12.1 and 12.2
• porous glass disk
• two electrodes:
anode (oxidation): Cu(s) to Cu2+
(aq)
cathode (reduction): Ag+
(aq) to Ag(s)
• electron flow is from anode (removed from Cu) to cathode (supplied to Ag+
)
• this is general
• need for solution connection: maintain charge neutrality (eg. Na+
and NO3
-
move
through bridge)
• special aspects of this example, not necessarily in all voltaic cells:
• electrodes can be simple conductors, not chemical participants (eg. Pt foil)
• need not have deposition of solid at either electrode
Galvanic & Electrolytic Cells
• external path as above, Example: Cd(s) + Ni2+
(aq) → Ni(s) + Cd2+
(aq)
Chem 59-110 (’02)

2
• at anode Cd is oxidized, Cd2+
goes into solution, balanced by NO3
-
coming from salt
bridge
• at cathode, Ni2+
is reduced, comes out of solution as Ni, balanced by NO3
-
going into
bridge
• negative charge as NO3
-
flows through bridge from cathode (Ni) to anode (Cd)
• negative charge as e-
flows through wire from anode (Cd) to cathode (Ni)
• conversely, Na+
moves through bridge to Ni side (Ni2+
consumed) from Cd side (Cd2+
produced)
• Galvanic Cell: spontaneous reaction; current flows between electrodes; potential difference
(see below) measured as a positive voltage; electrical energy produced; examples above
• Electrolytic Cell: non-spontaneous reaction is forced by applying a potential between
electrodes; electrical energy consumed; example, reverse of Cu/Ag reaction above:
(aq.)Ag2Cu(s)Ag(s)2(aq.)Cu2 ++
+→+
12.2 Free Energy & Cell Voltage
• in Galvanic cells above, spontaneous direction of e-
flow from Zn to Cu, from Cu to Ag and
from Cd to Ni, ie.- to lower (potential) energy, from anode to cathode
• potential difference measured in volts, V:
1V 1
J
C
=
• in Zn/Cu eg., Zn(s) Cu (aq.) Zn (aq.) Cu(s)2 2
+ → ++ +
, potential is 1.10 V
= driving force or electron pressure
= electromotive force, emf
= cell potential, ∆Ecell
Standard Cell Voltage
• define: ∆Eo
cell (or, ∆Eo
), standard emf or standard cell potential with 1 M in all solution
components (1 bar for gases) at 25o
C; cell potentials positive for the spontaneous (product-
favoured) direction
• note: reverse the reaction, change the sign of ∆Eo
∆Eo
and ∆Go
(Emf & Free-Energy Change)
• both are measures of a reaction’s tendency to proceed to products
state)(std.EFnΔGE;FnΔG oo
∆−=∆−=
n = number of moles of electrons transferred
F = the Faraday, molar equivalent of e- charge
= 96,500 C/(mol e-)
59-110 (’02), ch 12, Electrochemistry

3
= 96,500 J/(V.mol e-)
• Example 12.4: for Zn/Cu2+
reaction, Eo
= +1.10 V, calculate ∆Go
= - 212 kJ
Calculating the Cell Potential, ∆Eo
Half-Cell Potentials (Voltages)
• cell potential comprised of two half-cell potentials, standard oxidation potential (for one
couple) and standard reduction potential (for other couple)
∆ E E Ecell
o
ox
o
red
o
= +
• scaled according to a reference half-cell, standard hydrogen electrode, at 0 V (Fig. 12.3):
2 H (aq., 1M) 2e H (g, 1atm)2
+ −
+ →
• eg. voltaic cell with Zn anode and standard hydrogen electrode as cathode:
at anode: Zn(s) Zn (aq.) 2e2
→ ++ −
0EEEV0.76E o
ox
o
red
o
ox
o
cell +=+==∆
• this determines the standard oxidation half-cell potential for the Zn/Zn2+
couple
• opposite reaction, reduction of Zn2+
to Zn has a standard reduction potential of same
magnitude, opposite sign
• similarly for Cu/Cu2+
, but here Cu is the cathode and the hydrogen electrode is the anode
• cell potential = + 0.34V = standard oxidation potential for Cu/Cu2+
• return to Zn + Cu2+
reaction:
∆Eo
cell = Eo
ox + Eo
red
= 0.76 + 0.34 = 1.10 V (Fig. 12.4)
• can also determine half-reaction (half-cell) potentials by measurement against a known couple
other than the standard hydrogen electrode
• Example, Ni/Ni2+
is done vs. Zn/Zn2+
• do Example 12.5
• for simplicity of tabulation, all standard half-cell potentials are listed as reduction potentials,
eg. Appendix E
• note: half-cell potential an intensive property, hence not multiplied by stoichiometry numbers
Using Standard Electrode Potentials
Oxidizing and Reducing Agents
• trends in Appendix E
• the more positive an E°, the greater the tendency for the half-reaction to occur as written

4
• F2 most easily reduced, strongest oxidizing agent
• Li+
poorest oxidizing agent
• Li strongest reducing agent
• reaction between any substance on the left column and any one lower on the right column is
spontaneous (see below)
Predicting ∆Eo
& Spontaneity of Redox Reactions
• positive emf indicates spontaneous process
• predict from half-cell potentials:
eg. Fe(s) 2Ag (aq.) Fe (aq.) 2Ag(s)2
+ → ++ +
• half-cells:
Fe(s) Fe (aq.) 2e E 0.44 V
2Ag (aq.) 2e 2Ag(s) E 0.80 V
2
ox
o
red
o
→ + =
+ → =
+ −
+ −
• overall: ∆E° = 1.24 V ∴ spontaneous
• note: complication if adding or subtracting half-cell reactions to get a new half-cell reaction;
then have to account for stoichiometry
(aside: Electrical Work)
• voltaic cell: wmax = -nF∆E (spontaneous)
• electrolytic cell: wmin = -nF∆E (non-spontaneous)
12.3 Concentration Effects & the Nernst Equation
KlnRTEFn
KlnRTΔG
o
o
−=∆−∴
−=
K)298(atVK,log
n
0.059
Kln
n
0.026
Klog
nF
RT
2.30Kln
nF
RT
Eo
==
==∆
• from: ∆ ∆G G RTlnQo
= + (NB: implicit in ch. 9.7)
QlogRT2.30EFnEFn o
+∆−=∆−∴
E E
RT
nF
ln Q = E
2.30RT
nF
log Q Nernst Equ'n
= E
0
n
ln Q = E
0.059
n
log Q (at 298 K)
o o
o o
= − −
− −
.026
• used to calculate emf under non-standard conditions, Example 12.7

5
• or, to calculate a concentration, if cell emf measured. Example 12.9
• or, to calculate an equilibrium constant, Example 12.8
• application: the pH meter (read for interest)
• read 12.4: Batteries & Fuel Cells for interest
12.5 Corrosion & Its Prevention
• nearly all metals undergo thermodynamically favoured oxidation in air
• "skin" of oxide frequently protects against further oxidation
• cathodic and anodic areas on the metal, Fig. 12.14
• for iron, if limited O2 (slow) :
anode: Fe(s) Fe 2e ; E 0.44V2
ox
o
→ + =+ −
cathode: 2 H O(l) 2e H (g) + 2 OH (aq); E - V2 2
-
red
o
+ → =−
083.
• for iron, if O2 and water available (fast) :
anode: Fe(s) Fe 2e ; E 0.44V2
ox
o
→ + =+ −
V1.23EO(l);H2e4(aq.)H4(g)O:cathode o
red22 =→++ −+
• and Fe2+
oxidized further, near cathode:
4 Fe (aq.) O (g) (4 2x) H O(l) 2 Fe O .x H O(s) 8 H (aq.)2
2 2 2 3 2
+ +
+ + + → +
• can protect iron from corrosion:
• anodic inhibition: coating with tin; good until surface broken; more recently, promote
Fe2O3 formation as “skin”
• cathodic inhibition: coating with Zn; Zn is sacrificial anode in preference to Fe; "cathodic
protection", i.e.- force the metal to become a cathode
12.7 Electrolysis
• reverse of foregoing: non-spontaneous redox reactions "driven" by electrical energy
• eg. molten NaCl
• note sign convention opposite to voltaic cell
• used in some metals production
• electrolysis of aqueous solutions
• eg. brine:
• at cathode, H2 production preferred:
V0.83E(aq.),OH2(g)He2O(l)H2 o
red22 −=+→+ −−
Na (aq.) e Na(s), E 2.71Vred
o+ −
+ → = −
• while at anode, little thermodynamic preference:

6
2Cl (aq.) Cl (g) 2e , E 1.36 V
2H O(l) 4H (aq.) O (g) 4e , E 1.23 V
2 ox
o
2 2 ox
o
− −
+ −
→ + = −
→ + + = −
(if use NaI, where I-
→ I2, Eo
= -0.535 V; more clearcut)
• and, due to "overvoltage", Cl
-
oxidation preferred
• net:
2Cl (aq.) 2H O(l) Cl (g) H (g) 2OH (aq.)2 2 2
− −
+ → + +
• all products are industrial commodities
• minimum emf required: 0.83 + 1.36 = 2.19 V
Electrolysis with Active Electrodes
• electroplating: make anode of metal to be deposited onto an object; make object the cathode
• eg. nickel
anode: Ni(s) Ni (aq.) 2e ; E 0.28V2
ox
o
→ + = ++ −
(in preference to water electrolysis)
cathode: Ni 2e Ni(s)2+ −
+ →
Suggested Problems
1; 11 – 21; 27 – 33; 51, 53; all odd

Ch 12 electrochemistry

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